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• Unit 1 : INTRODUCTION TO ORGANIC CHEMISTRY

  UNIT 1 INTRODUCTION TO ORGANIC CHEMISTRY

  Key unit competency

  Apply IUPAC rules to name organic compounds and explain types of isomers for

  organic compounds

  Learning objectives

  • Classify organic compounds as aliphatic, alicyclic and aromatic

  • Determine different formulae for given organic compounds

  • Describe the common functional groups and relate them to the homologous series

  • Use IUPAC rules to name different organic compounds

  • Describe the isomers of organic compounds

  
  

  Introductory activity

  Consider the following substances: Sodium chloride, starch, table sugar,

  magnesium carbonate, glucose, sodium hydrogen carbonate, water.

  1. Heat a small sample of each ( 5g for solids, 5ml for liquids) in a crucible

  2. Record your observations.

  3. From the observations, classify the substances listed above.

  4. What criterion do you use for that classification?

  5. Interpret your observations

  Organic chemistry is defined as the study of the compounds mainly composed by

  carbon and hydrogen atoms, and sometimes oxygen, nitrogen, phosphorus, sulphur

  and halogens atoms. The study of the rest of the elements and their compounds falls

  under the group of inorganic chemistry. However, there are some exceptions such

  as carbonates, cyanides, carbides, carbon oxides, carbonic acid, carbon disulphide

  which are considered as inorganic compounds. Since various organic compounds

  contained carbon associated with hydrogen, they are considered as derived from

  hydrocarbons. Thus, a more precise definition of organic chemistry is: “the study of

  hydrocarbons and the compounds which could be thought of as their derivatives’’.

  The organic and inorganic compounds can be differentiated based on some of their

  properties as summarised in the following table.

  

  Why to study organic chemistry as a separate branch?

  The organic chemistry involves the study of all chemical reactions that are commonly

  used in industries and many other organic reactions that take place in living systems.

  Materials used in everyday life, food processing and other manufacturing objects are

  obtained based on organic chemistry. Some other reasons are highlighted below.

  • Large number of compounds: up to now, no one knows exactly the number

  of organic compounds that are present in nature.

  • Built of relatively few elements: The elements frequently encountered

  in organic compounds are carbon, hydrogen, oxygen, nitrogen, sulphur,

  phosphorous, and halogens;

  • Unique characteristic of carbon to undergo catenation: carbon atom is

  unique among other elements whose atoms possess the capacity to unite

  with each other by the covalent bonds resulting in a long chain of carbons ( i.e:

  polysaccharides, proteins, polyesters, polyamides…).

  Isomerism is the existence of compounds that have the same molecular formula

  but different arrangements of atoms; these compounds are called “isomers”.

  • Functional groups as basis of classification: Organic molecules contain

  active atoms or groups of atoms which determine their chemical behaviour.

  These are called functional groups joined in a specific manner. Therefore,

  organic compounds with similar functional groups display similar properties and form a class

  • Combustibility: organic compounds are combustible.

  • Nature of chemical reactions: organic compounds being formed by covalent

  bonds, they are slow and often have a low yield.

  Importance of organic chemistry

  The organic chemistry is a subject that plays an important role in modern life. In

  general, there is no art, science or industry where knowledge of organic chemistry

  is not applied.

  Examples where organic chemistry is applied:

  1)Application in daily life.

  In our day-to-day life, we find many substances or materials that are commonly used

  and the later are made of organic compounds.

  • Food: starch, fats, proteins, vegetables,...

  • Clothes: cotton, wool, nylon, dacron, ....

  • Fuels: petrol, diesel oil, and kerosene

  • Dyes of all kinds

  • Cosmetics (body lotion,…)

  • Soaps and detergents

  • Medicine: cortisone, sulphonamide, penicillin,…

  • Drugs: morphine, cocaine,...

  • Stationery: pencils, paper, writing ink,…

  • Insecticides,rodenticides,ovicides …

  2)Applications in industry

  The knowledge of organic chemistry is required in many industries such as

  manufacture of food, pharmacy, manufacture of dyes and explosives, alcohol

  industry, soil fertilisers, petroleum industry, etc.

  
  

  3)Study of life processes

  Organic chemistry in other words is the chemistry of life. For example the vitamins, enzymes, proteins and hormones are important organic compounds

  produced in our body to ensure its proper development.

  

  

  (https://chemistry.tutorvista.com/organic-chemistry/hydrocarbons.html)

  1.1.1. Aliphatic compounds

  Aliphatic compounds are organic compounds in which the carbon atoms are

  arranged in a straight or branched chain.

  

  1.1.2. Alicyclic compounds

  Alicyclic compounds are organic compounds that contain one or more carbon rings

  that may be saturated or unsaturated

  

  1.1.3. Aromatic compounds

  Aromatic compounds are compounds that contain a closed ring that consists of

  alternating single and double bonds with delocalised pi electrons.

  

  Aromatic compounds are designated as monocyclic, bicyclic and tricyclic if they

  contain one, two or three rings, respectively.

  Examples:

  

  Note: Heterocyclic compounds: Are also classified as cyclic compounds which

  include one or two atoms other than carbon (O, N, S) in the ring.Thus furan, thiophene

  and pyridine are heterocyclic compounds.

  

  

  1.2. Types of formulas for organic compounds

  
  

  
  

  Atoms bond together to form molecules and each molecule has a chemical formula.

  In organic chemistry, we can distinguish empirical, molecular and structural formulas.

  1.2.1. Empirical formula

  The empirical formula is the simplest formula which expresses the ratio of the number

  of atoms of each element present in a particular compound. The empirical formula

  is determined using the percentage composition according to the following steps.

   i. The percentage of each element, considered as grams of that element in 100g

  of the compound, is divided by its atomic mass. This gives the number of moles

  of the element in 100g of the compound.

   ii. The result in i. is then divided by the lowest ratio (number of moles in 100g of

  the compound), seeking the smallest whole number ratio.

   iii. If the atomic ratios obtained in ii. are not the whole number, they should be

  multiplied by a suitable common factor to convert each of them to the whole

  numbers (or approximatively equal to the whole numbers). Minor fractions are

  ignored by rounding up or down (ex: 7.95 = 8).

  
  

  
  

  
  

  1.1.2. Molecular formula

  The molecular formula is a formula expressing the exact number of atoms of each

  element present in a molecule.

  Molecular formula = Empirical formula x n 

  

  Example 1:

  An organic compound contains 31.9% by mass of carbon, 6.8% hydrogen and 18.51%

  nitrogen and the remaining percentage accounts for oxygen. The compound has a

  vapour density of 37.5. Calculate the molecular formula of that compound.

  

  Vapour density = a half molecular mass

  Molecular mass = 2 x vapour density = 2 x 37.5 = 75g/mol

  

  Note: From the above calculations, we can extend our generalized expression: : From the above calculations, we can extend our generalized expression:

  % of Oxygen = 100 – (% hydrohen + % carbon)

  1.2.3. Structural formulas

  Structural formula shows how the different atoms in a molecule are bonded (i.e.

  linked or connected)

   There are three types of structural formulas: displayed, condensed and skeletal

  (stick) formulas.

  

  1.3. Functional groups and homologous series

  1.3.1 Functional groups

  
  

  
  

  A functional group is an atom or group of atoms in a molecule which determines

  the characteristic properties of that molecule. Examples of some fuctionnal groups

  are indicated in the Table 1.2.

  
  

  1.3.2. Homologous series
  

  
  

  
  

  When members of a class of compounds having similar structures are arranged in

  order of increasing molecular mass, they are said to constitute a homologous series.

  Each member of such a series is referred to as a “homologous” of its immediate

  neighbours. For example, the following sequence of straight chain of alcohols forms

  a homologous series.

  
  

  
  

  
  

  Characteristics of a homologous series

  1. Any member of the series differs from the next by the unit –CH₂-(methylene group)

  2. The series may be represented by a general formula of alcohols which is

  C_nH_2n+1 OH where n =1,2,3, etc.

  3. The chemical properties of the members of a homologous series are

  similar, though in some series the first members show different behaviour.

  4. The physical properties such as density, melting point and boiling point

  generally increase within the molecular mass.

  
  

  
  

  
  

  1.4. General rules of nomenclature of organic compounds according to IUPAC

  
  

  
  

  The organic compounds are named by applying the rules set by the International

  Union of Pure and Applied Chemistry (IUPAC). The purpose of the IUPAC system of

  nomenclature is to establish an international standard of naming compounds to

  facilitate the common understanding.

  In general, an IUPAC name has three essential parts:

  • A prefix that indicates the type and the position of the substituents on the

  main chain.

  • The base or root that indicates a major chain or ring of carbon atoms found

  in the molecule’s structure. e.g. Meth- for one carbon atom, eth- for 2 carbon

  atoms, prop- for 3 carbon atoms, hex- for five carbon atoms, etc.

  • The suffix designates the functional group.

  Example -ane for alkanes, -ene for alkenes, -ol for alcohols, -oic acid for carboxylic

  acids and so on.

  Steps followed for naming organic compounds:

  1. Identify the parent hydrocarbon:

  It should have the maximum length, or the longest chain

  

  3. Identification of the side chains.

  Side chains are usually alkyl groups. An alkyl group is a group obtained by a

  removal of one hydrogen atom from an alkane. The name of alkyl group is obtained

  by replacing -ane of the corresponding alkane by –yl (Table 1.3).

  

  

  4.If the same substituent occurs two or more times, the prefix di, tri,tetra, ...is

  attached to substituent’s name. Its locants separate the prefix from the name of the

  substituent.

  

  5.Identify the remaining functional groups, if any, and name them. Different side

  chains and functional groups will be listed in alphabetical order.

  

  The prefixes di, tri, tetra,...are not taken into consideration when grouping

  alphabetically. But prefixes such iso-, neo- are taken into account.

  Example:

  

  Identify the position of the double/triple bond.

  Example:

  

  The sum of the numbers which show the location of the substituents is the possible smallest

  

  The correct name will be the one which shows the substituents attached to the third

  and fifth carbon, respectively and not to the fourth and the fiveth carbon atom.

  Numbers are separated by commas Hyphens are added between numbers and

  words. Successive words are merged in one word.

  

  

  1.5. Isomerism in organic compounds

  
  

  
  

  
  

  Isomerism is the existence of compounds that have the same molecular formula

  but different arrangements of atoms; these compounds are called “isomers”.

  Isomers have different physical or/and chemical properties and the difference may

  be great or small depending on the type of isomerism.

  There are two main classes of isomerism: Structural isomerism and stereoisomerism.

  
  

  1.5.1. Structural isomerism
  

  

  1. Position isomerism

  Position isomers are compounds with the same molecular formula but different

  positions of the functional group or substituent(s).

  

  2. Chain isomerism

  Chain isomers are compounds with the same molecular formula, belonging to the

  same homologous series, with chain of carbon atoms of different length.

  

  3. Functional isomerism

  Functional (group) isomers are compounds which have the same molecular formula

  but different functional groups.

  Examples:

  

  1.5.2. Stereoisomerism

  
  

  1. Geometrical isomerism

  Geometrical isomers or cis-trans isomers are compounds with the same molecular

  formula, same arrangement of atoms but differ by spatial arrangements.

  This type of isomers is mainly found in alkenes due to the restricted rotation around

  the carbon-carbon double bond.

  Note: For more information, visit the website below. (https://www.youtube.com/

  watch?v=7tH8Xe5u8A0).

  The necessary condition for an alkene to exhibit geometrical isomerism is that each

  carbon doubly bonded has two different groups attached to it.

  
  

  
  

  2. Optical isomerism
  

  

  Optical isomers are compounds with the same molecular formula and arrangements

  of atoms but have different effect on the plane polarised light.

  • A compound that rotates the plane polarised light is said to have an optical activity.

  • This type of isomerism occurs in compounds containing an asymmetric

  (asymmetrical) carbon atom or chiral centre^1.

  .

  • When a molecule has chiral centre, there are two non superimposable isomers

  that are mirror images of each other.

  • Such compounds are called enantiomers

  
  

  
  

  
  

  In a mirror, the left hand is the image of the right hand and they are non

  superimposable, i.e. they are enantiomers. An achiral object is the same as its mirror

  image, they are nonsuperimposable.

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  

  URLs: 2Files: 3
• Unit 2 : ALKANES

  UNIT 2: ALKANES

  Key unit competency

  Relate the physical and chemical properties of the alkanes to the preparation

  methods, uses and isomerism.

  Learning objectives

  • Name straight chain alkanes up to carbon-20

  • Define homologous series

  • Use IUPAC system to name straight and branched alkanes

  • Describe the preparation methods of the alkanes

  • Prepare and collect methane gas

  • Respect of procedure in experiment to carry out preparation of methane or propane

  • Describe and explain the trend in physical properties of homologous series ofalkanes

  • Be aware of the dangers associated with combustion reactions of the alkanes

  • Write reaction for free radical mechanism for a photochemical reaction

  • State the chemical properties of the alkanes

  • Develop practical skills,interpret results make appropriate deductions.

  • Appreciate the importance of the alkanes in daily life

  • Appreciate the dangers caused by the alkanes to the environment as major

  sources of air contaminants

  • State the uses of the alkanes

  

  Alkanes are the simplest class of organic compounds. They are made of carbon and

  hydrogen atoms only and contain two types of bonds, carbon-hydrogen (C-H) and

  carbon-carbon (C-C) single covalent bonds. They do not have functional groups.

  Alkanes form a homologous series with the general formula C_nH_2n+2 where n is the

  number of carbon atoms in the molecule. The first member of the family has the

  molecular formula CH₄ (n=1) and is commonly known as methane and the second

  member with molecular formula is C₂H₆ (n=2) is called ethane.

  These compounds are also known as saturated hydrocarbons. This name is more

  descriptive than the term “alkane’’ because both their composition (carbon and

  hydrogen) and the fact that the four single covalent bonds of each carbon in their

  molecules are fully satisfied or ‘’saturated’’.

  The name alkane is the generic name for this class of compounds in the IUPAC

  system of nomenclature. These hydrocarbons are relatively unreactive under

  ordinary laboratory conditions, but they can be forced to undergo reactions by

  drastic treatment. It is for this reason that they were named paraffins (Latin parum

  affinis = little activity).

  2.1. Nomenclature of alkanes

  
  

  
  

  
  

  IUPAC Rules for the nomenclature of alkanes

  a. Find and name the longest continuous carbon chain.

  b. Identify and name groups attached to this chain.

  c. Number the chain consecutively, starting at the end nearest a substituent

  group.

  d. Designate the location of each substituent group by an appropriate

  number and name.

  e. Assemble the name, listing groups in alphabetical order. The saturated

  hydrocarbon form homologous series (series in which members have similar

  chemical properties and each differs from the preceding by a methylene

  group –CH₂-).

  The first four members are known by their common names, from C₅ and

  above the Roman prefixes indicating the number of carbon atoms is written

  followed by the ending “ane” of the alkanes.

  Note: Alkyl groups are obtained when one hydrogen atom is removed from

  alkanes; therefore their names are deduced from the corresponding alkanes

  by replacing “ane” ending with “yl” desinence (Table 2.1)

  
  

  
  

  
  

  Prefixes di, tri, tetra, sec, tert, are not considered when alphabetizing.
  

  f. In case of chains of the same length, the priority is given for part where

  many branched of alkyl groups appear.

  

  g. For cyclanes or cycloalkanes, the prefix “cyclo” is recommended, followed
  by the name of the alkanes of the same carbon number.

     But in case of ramified cyclanes, the priority is for the ring.

  

  

  Note: If there are more than one substituent, the numbering is done so that the

  sum of the numbers used to locate the locants is minimum. This is the lowest sum

  rule.

  

  

  2.2. Isomerism

  

  Alkanes show structural isomerism. The easiest way to find isomers is to draw the

  longest chain of carbon atoms first and then reduce it by one carbon first until

  repetition begins to occur.

  

  d. Substitute one hydrogen on carbon (3) by the mathyl group

  

  e.Longest chain reduced further to 4 carbon atoms by cutting 2 methyl group

  

  -Putting the methyl group on position 1 or 5 gives you the same straight chain

  isomer.

  

  2.3 Occurrence of Alkanes

  
  

  
  

  1. The alkanes exist in nature in form of natural gases and petroleum. Natural

  gas and petroleum existence are the results of decomposition of died

  bodies after many years ago.

  2. The most natural gas is found in lake Kivu as methane gas but in form of

  traces like ethane, propane and butane

  3. Petroleum is the most world energy, it is formed by decomposition by

  bacteria for millions of years died marine living things and as the last

  product is petroleum and natural gases which are separated in fractional

  distillation of their crude oil and the results are obtained according to their

  boiling point. 

  
  

  
  

  
  
  
  

  2.4. Preparation of alkanes

  
  

  
  

  
  

  
  

  

  
  Note: The reaction is practically used to reduce by one carbon the length of carbon chain. It is referred as decarboxylation of sodium carboxylates.
  
  

  Other reactions used for the preparation of alkanes are the following:

  1. Addition reaction of hydrogen to alkenes and alkynes in the presence of

  catalyst like Nickel, Palladium or platinum produces alkanes: this reaction

  is called hydrogenation reaction of alkenes and alkynes; it is also called a

  reduction reaction of alkenes and alkynes.

  

  2. From halogenoalkanes or Alkyl halides

   On reduction of alkyl halides with Zn and concentrated hydrochloric acid, alkyl

  halides are converted to alkanes. 

  

  

  b) Alkyl halides when heated with sodium metal in ether solution give higher alkanes (alkanes with more carbon atoms) (Wurtz reaction).

  

  c) When Alkyl halides are treated with Zn-Cu couple, in the presence of ethanol,

  alkanes are formed.

  Note: Zn-Cu couple is obtained by adding Zinc granules in aqueous copper (II)

  sulphate solution where copper is deposited on the Zn pieces. 

  
  
  

  3. From carbonyl compounds

  Reduction of carbonyl compounds, with amalgamated Zinc (alloy made of zinc

  and mercury) and HCl. This is the Clemmensen reduction). 

  
  
  

  

  

  2.5. Physical properties of alkanes

  

  

  The above Table shows that the boiling and melting points of homologue alkanes

  increase with the number of carbon i.e. molecular mass.

  Explanation:

  The boiling and melting points depend on the magnitude of the Van Der Waal’s

  forces that exist between the molecules. These forces increase in magnitude with

  molecular mass. 

  Note: Branched chain isomers have lower boiling and melting points than their

  straight chain isomers, because straight chain isomers are closer packed than the

  branched chain isomers.

  

  

  2.6. Chemical properties of alkanes

  

  Generally, alkanes are quite inert towards common reagents because:

  • The C-C bond and C-H bonds are strong and do not break easily.

  • Carbon and hydrogen have nearly the same electronegativity valuehence

  • C-H bond only slightly polarized; generally C-H bond is considered as covalent.

  • They have unshared electrons to offer.

  They, however, undergo the following reactions.

  1. Reaction with oxygen

  Alkanes react with oxygen to produce carbon dioxide (if oxygen is enough to

  burn all quantity of hydrocarbons), or carbon monoxide or carbon if oxygen is in

  insufficient quantity, and water. This reaction is called “combustion”

  

  Carbon dioxide (CO₂) produced from the burning of alkanes or fossil fuels for heating, transport and electricity generation is the major atmospheric pollutant that increases the green house potential of the atmosphere .Carbon dioxide is the major Green House Effect (GHE) gas.

  Burning wood and forests produce also carbon dioxide and lead to the increase of that gas in the atmosphere. Methane as another GHE gas is produced by human activities, agriculture (Rice), and cattle-rearing.
  

  
  

  There are many natural ways of reducing atmosphere carbon dioxide:

  i. Water in seas dissolves millions of tonnes of gas (but less now than it did in
  

  the past, since the average ocean temperature has increased by 0.5 ^oC in the last 100 years, and gases are less soluble in hot than in cold water).

  ii. Plankton can fix the dissolved carbon dioxide into their body mass by photosynthesis

  iii. Trees fix more atmospheric carbon dioxide than do grass and other vegetation through photosynthesis according to the equation below.

  
  

  There are other ways than natural ways of reducing GHE gases and among them

  there are the use of technologies that reduce the green house gas emissions, the

  recycling of the GHE.

  Notice: (i) Br2

   reacts as Cl2

   but slowly while iodine reacts hardly or does not.

  Notice: (i) Br2 reacts as Cl2 but slowly while iodine reacts hardly or does not

  ,Fluorine, the most electronegative element of the periodic table reacts with

  alkanes to give coke,

  i.e. a decomposition reaction:

  
  
  

  A mechanism of a reaction is a description of the course of the reaction which

  shows steps of the reaction and the chemical species involved in each step.

  The mechanism for the reaction between methane and bromine is the following:

  

  

  (ii) Due to radical formation involved, the main product of reaction is the one from

  the most stable radical, starting with tertiary, secondary, primary and methyl in

  decreasing order of stability.

  A tertiary free radical is better stabilised by the electron donating methyl groups

  than the secondary, primary and methyl ones where the carbon atom is attached

  to more hydrogen atoms

  

  3. Dehydrogenation of alkanes gives alkenes under heat and a catalyst like V₂O₅

  
  

  4. Cracking

  On heating or in the presence of a catalyst, large molecules of alkanes are

  decomposed into smaller alkanes and alkenes. If the cracking is performed on

  heating, it is referred as themocracking.

  If the cracking is performed using a catalyst; it is referred as catalytic cracking and

  many products result from one reactant as shown below.

  

  

  2.7. Uses of alkanes
  

  

  1. Methane (CH₄)

  Methane finds many uses:

  • It is used as a fuel at homes, ovens, water heaters, kilns and automobiles as it

  combusts with oxygen to produce heat.

  • Highly refined liquid methane is used as rocket fuel.

  • Methane is used as fuel for electricity generation.

  • It is used as a vehicle fuel in the form of liquefied natural gas (LNG).

  • Methane can be used as raw material in the production of urea, a fertilizer.

  In general, methane is more environmental friendly than gasoline/petrol and

  diesel.

  2. Butane (C₄H₁₀)

  • Butane is a key ingredient of synthetic rubber.

  • It is used as fuel in cigarette lighters.

  • When blended with propane and other hydrocarbons, it may be referred to

  commercially as LPG, for liquefied petroleum gas.

  • Butane gas cylinders are used in cooking.

  • Also used in aerosol spray cans.

  3. Propane (C₃H₈)

  • Propane is used as a propellant for aerosol sprays such as shaving creams

  and air fresheners.Used as fuel for home heat and back up electrical

  generation in sparsely populated areas that do not have natural gas

  pipelines.

  • Propane is commonly used in movies for explosions

  4. Ethane (C₂H₆)

  • Ethane is used in the preparation of ethene and certain heavier

  hydrocarbons.

  • Ethane can be used as a refrigerant in cryogenic refrigeration systems.

  5. Pentane (C₅H₁₂)

  • Pentane is used in the production of polystyrene foams and other foams.

  • Used in laboratories as solvents.

  • It is also an active ingredients of pesticides.

  • Used as solvent in liquid chromatography

  6. Hexane (C₆H₁₄)

  • It is used in the formulation of glues for shoes, leather products, and roofing.

  • It is also used to extract cooking oils such as canola oil or soy oil from seeds.

  • Hexane is used in extraction of pyrethrine from pyrethrum; e.g. Horizon

  SOPYRWA (a pyrethrum factory in Musanze District).

  • Also for cleansing and degreasing a variety of items, and in textile

  manufacturing.

  7. Heptane (C₇H₁₆)

  • Heptane is used as solvent in paints and coatings.

  • Pure n-heptane is used for research, development and pharmaceutical

  manufacturing

  • Also as a minor component of gasoline.

  • It is used in laboratories as a non-polar solvent.

  

  

  

• Unit 3 : ALKENES AND ALKYNES

  UNIT 3: ALKENES AND ALKYNES

  Key unit competency

  Relate the physical and chemical properties of alkenes and alkynes to their

  reactivity and uses

  Learning objectives

  • Explain the reactivity of alkenes in comparison to alkanes

  • Explain the existence of geometrical isomerism in alkenes

  • Describe the industrial process of preparing alkenes and alkynes

  • Apply IUPAC rules to name alkenes and alkynes

  • Carry out an experiment to prepare and test ethene gas

  • Outline the mechanisms for electrophilic addition reactions for alkenes and

  alkynes

  • Write the structural formulae of straight chain alkenes and alkynes

  • Apply Markovnikov’s rule to predict the product of hydrohalogenation of

  alkenes

  • Classify alkynes as terminal and non-terminal alkynes using their different

  structures

  • Appreciate the combustion reaction as source of fuels.

  • Appreciate the uses and dangers of addition polymers (polythene used for

  polythene bags, polypropene for plastic bottles etc.)

  Introductory activity

  Observe the following picture and answer the questions that follow.

  

  1. What is the collective name of the substances used to manufacture the

  items showed in the above picture?

  2. a. What are the raw materials used in the manufacture of the substances

  identified in 1)?

  b. These raw materials may be obtained from different sources. Discuss

  this statement.

  c. Do you expect these raw materials be soluble or not in water? Justify

  your answer.

  3. Even though the items which appear in the picture above are interesting,

  they also present some disadvantages. Discuss this statement.

  3.1. Definition, structure and nomenclature of alkenes

  
  

  Alkenes are a homologous series of hydrocarbons which contain a carbon-carbon

  double bond. Since their skeleton can add more hydrogen atoms, they are referred

  as unsaturated hydrocarbons.

  The general formula of alkenes is C_nH_2n.

  
  

  Alkenes are abundant in the nature and play important roles in biology. Ethene,

  for example, is a plant hormone, a compound that controls the plant’s growth and

  other changes in its tissues.

  Ethene affects seed germination, flower maturation, and fruit ripening.

  They are described as unsaturated hydrocarbons because they can undergo

  addition reactions.

  The double bond in alkenes is made of one sigma bond and one pi bond. This gives rise to the impossibility of rotation around the double bond. The hybridization state in alkenes is sp2 and the structure around each carbon doubly bonded is trigonal planar with a bond angle value of 120^o

  
  

  .IUPAC names of alkenes are based on the longest continuous chain of carbon

  atoms that contains the double bond.

  The name given to the chain is obtained from the name of the corresponding

  alkane by changing the suffix from –ane to –ene.

  If the double bond is equidistant from each end, number the first substituent that

  has the lowest number. If there is more than one double bond in an alkene, all of

  the bonds should be numbered in the name of the molecule, even terminal double

  bonds. The numbers should go from lowest to highest, and be separated from one

  another by a comma.

  The chain is always numbered from the end that gives the smallest number for the

  location of the double bond.

  In naming cycloalkenes, the carbon atoms of the double bond are numbered

  1 and 2 in the direction that gives the smallest numbers for the location of the

  substituents.

  If a compound contains two or more double bonds, its location is identified by a

  prefix number. The ending is modified to show the number of double bonds:

  • a diene for two double bonds,

  • a triene for two three bonds

  • a tetraene for four double bonds

  

  

  3.2. Isomerism in alkenes

  
  

  Alkenes exhibit two types of isomerism: structural isomerisms and

  stereoisomerism.

  1. Structural isomerism

  Alkenes show as well position isomerism, chain isomerism and functional isomerism.

  In position isomerism, the position of the double bond changes but the length

  of the chain remains the same.

  
  

  
  

  Alkenes and cycloalkanes have two fewer hydrogen atoms than alkanes. That is

  why, they have the same molecular formula. However, they belong to different

  homologous series. Therefore, they are functional group isomers. This isomerism

  that relates open chain compounds to ring chain compounds is referred to as ring

  isomerism.

  
  

  
  

  2. Stereoisomerism

  Due to the impossibility of rotation around the double bond, alkenes give rise to

  cis-trans or geometrical isomerism.

  

  3.3. Preparation of alkenes

  
  

  
  

  Different methods are used for the preparation of alkenes. Most of them are

  elimination reactions.

  1. Dehydration of alcohols

  An alkene may be obtained by dehydration of an alcohol. The reaction involves

  the loss of H and OH (water) from adjacent carbons of an alcohol to form an  alkene. The dehydration is carried out by heating an alcohol with concentrated sulphuric acid or 85% phosphoric acid.

  
  

  If two or more alkenes may be obtained, the one having more substituents on

  the double bond generally predominates. This is the Zaitsev’s rule.

  

  This is due to the stability of the intermediate carbocation. The carbocation

  produced in step 2 may undergo a transposition (rearrangement) of a hydride ion

  or a methyl group giving a more stable carbocation and therefore a more stable

  alkene.

  

  Mechanism

  
  

  
  

  
  

  
  

  2. Dehydrohalogenation of halogenoalkanes

  

  alkenes. The reaction follows the Zaitsev’s rule.

  
  

  
  

  
  

  3. Dehalogenation of dihalogenoalkanes

  When a compound containing two halogen atoms on the adjacent carbon

  atoms is treated with magnesium or zinc it transforms to an alkene.

  

  

  3.4. Laboratory preparation and chemical test for ethene

  
  

  Activity 3.4
  

  Preparation of ethene Set up the apparatus as shown in the Figure below (Figure 3.1) and follow the instructions to perform the experiment on the preparation ofethane

  
  

  Requirements:

  Chemicals:

  • Ethanol, aluminium oxide, lime water, mineral wool, bromine water,

  acidified potassium permanganate solution (very dilute), water.

  Additional apparatus:

  • Boiling tube

  • Rubber stopper with hole

  • Delivery tube

  • Trough

  • Test- tube rack

  • 5 test tubes

  • 5 rubber stoppers for test tubes

  • Spatula Procedure and setting

  • Bunsen burner

  • Glass rod

  • Splint

  • Matches 

  1. Preparation of ethene:

  - Pour some ethanol into the boiling tube to a 3 cm depth

  - Add some glass wool to soak up the ethanol, using a glass rod to push the wool

  down the tube.

   - Clamp the boiling tube in a horizontal position using a retort stand.

   - Put a small amount of aluminium oxide about half way along the boiling tube.

  - Complete the set up of the apparatus as shown in the diagram above.

  - Light the Bunsen burner, adjust it to a blue flame and heat the aluminium

  oxide. (Make sure the test tube is filled with water when you start to collect the

  gas produced.)

   - As the aluminium oxide gets hot the heat reaches the ethanol at the end of the

  tube. The ethanol then changes to vapour, passes over the hot aluminium oxide

  and is dehydrated to produce ethene gas.

  - Collect 5 test tubes of the gas and put a stopper on each tube when it is filled.

  - When the test tubes have all been filled, loosen the retort stand and raise the

  apparatus so that the delivery tube no longer dips into the water. This avoids

  suck back of water as the tube begins to cool which could cause the boiling

  tube to crack. Turn off the Bunsen burner. 

  2. Testing the properties of ethene

  Addition of bromine:

  - Taking great care, add about 1ml of the test tube of bromine water to one of

  the test tubes of ethene.

   - Replace the stopper and shake the tube a few times.

   - Record your observations.

  - Write down your conclusions 

  - Addition of acidified potassium permanganate:

  - Add about 1ml of very dilute potassium permanganate solution to one of the

  test tubes of ethene and shake the tube a few times.

   - Record your observations.

  - Write down your conclusions

   Combustion:

  - Remove the stopper of one of the tubes filled with ethene and apply a light to

  the mouth of the test tube using a lighted splint.

  - Allow the gas to burn and when it has stopped burning add a small amount of

  lime water to the test tube, stopper it and shake the tube a few times.

  - Write down your observations.

  Interpretation

  When ethanol is heated in the presence of aluminium oxide, a gas is produced. This

  gas does not react with lime water. This means that the produced gas is not carbon

  dioxide. The equation of the reaction is: 

  

  The gas decolourises bromine water. Bromine water is a test used to identify the

  presence of a carbon-carbon double bond or triple bond. The bromine adds across

  the double bond and a dibromoalkane is formed. The reaction between alkene and

  bromine water is shown below:

  

  If you shake an alkene with bromine water (or bubble a gaseous alkene through

  bromine water), the solution becomes colourless. Alkenes decolourise bromine

  water.

  The Figure 3.2 shows Bromine water added to ethene: before the reaction (left)

  the color of bromine appears, and after the reaction (right) the colour of bromine

  disappears.

  

  When ethene reacts with acidified potassium manganate (VII), the purple colour of

  the permanganate solution turned to colourless or light pink indicating the presence

  of the carbon – carbon double bond.The reaction is the following:

  

  The gas burns with a smoky flame producing carbon dioxide and heat energy. The

  carbon dioxide produced turns into milky lime water.

  

  

  3.5. Physical properties of alkenes

  

  • Alkenes which have less than 5 carbon atoms are gaseous at ordinary

  temperature, the other are liquid up to 18 while others are solids as the

  number of carbon atoms increases.

  • Boiling points and melting points of alkenes are less than those of alkanes

  but also increase as the molecular weight increase.

  • Alkenes are insoluble in water but soluble in most organic solvents.

  • Cis-alkenes have a slightly higher boiling point than the trans-isomers

  because the dipole moments in trans structures cancel each others----.

  

  3.6. Chemical properties

  3.6.1. Addition reactions
  3.6.1.1. Electrophilic additions

  
  

  
  

  
  

  Alkenes are far more reactive than alkanes due to the carbon-carbon double bond.

  These compounds are unsaturated and they can easily undergo addition reactions

  to yield saturated products.

  
  

  The double bond in alkenes is a region of high density of electrons. Therefore, this

  region is readily attacked by electrophiles. An electrophile is an atom, a molecule

  or an ion which is electron-deficient; i.e. it is a Lewis acid or an electron pair acceptor.

  
  

  
  

  
  

  1. Addition of hydrogen halides

  Hydrogen halides (HCl, HBr, HI) react with alkenes to yield halogenoalkanes. The

  reaction is carried out either with reagents in the gaseous state or in inert solvent

  such as tetrachloromathane. 

  

  When hydrogen halides add to unsymmetrical alkenes, the reaction leads to the

  formation of two products in two steps. The first step leads to the formation of

  two different carbocations with the major product formed from the more stable carbocation. This is the Markownikov’s rule. That is “The electrophilic addition of an unsymmetric reagent to an unsymmetric double bond proceeds by involving the most stable carbocation.

  
  

  
  

  
  

  2. Addition of water

  The hydration of alkenes catalysed by an acid is an electrophilic addition. Ethene

  can be transformed into ethanol. The first step consists of adding concentrated

  sulphuric acid. The second step consists of the hydrolysis of the product of the

  first step.

  In industry the reaction is carried out at approximately 300 °C in the prence of

  phosphoric acid as a catalyst.

  

  3. Addition of cold concentrated sulphuric acid

  When cold concentrated sulphuric acid reacts with alkene, an alkyl hydrogen

  sulphate is obtained. If the starting alkene is unsymmetrical, two different alkyl

  hydrogen sulphates are obtained. If the alkyl hydrogen sulphate is warmed in

  the presence of water, an alcohol is obtained.

  

  4. Addition of halogens

  The addition of halogens (halogenation) on alkenes yields vicinal

  dihalogenoalkanes. The reaction takes place with pure reagents or by mixing

  reagents in an inert organic solvent.

  When a chlorine or bromine molecule approaches an alkene, the pi electrons

  cloud interact with the halogen molecule causing its polarisation.

  

  

  3.6.1.2. Hydrogenation

  In the presence of a catalyst (Pt, Ni, Pd), alkenes react with hydrogen to give

  alkanes.

  

  Alkenes are readily oxidised due to the presence of the double bond.

  1. Reaction with oxygen

  i. Transformation to epoxides

  Ethene react with oxygen in the presence of silver as a catalyst to yield epoxyethane.

  
  

  
  

  

  3. Reaction with potassium permanganate

  Alkenes react with dilute potassium permanganate solution to give diols. The

  reaction takes place in the cold.

  The colour change depends on the medium of the reaction.

  

  

  4. Hydroformylation

  The hydroformylation is a process by which alkenes react with carbon monoxide

  and hydrogen in the presence of rhodium catalyst to give aldehydes.

  
  

  3.6.3. Addition polymerisation 

  
  

  
  

  Alkenes undergo addition polymerisation reaction to form long chain polymers.i.e

  a polymer is a large molecule containing a repeating unit derived from small unit

  called monomers. A polymerisation reaction involves joining together a large number of small molecules to form a large molecule.

  
  

  Many different addition polymers can be made from substituted ethene

  compounds.

  Each polymer has its physical properties and therefore many polymers have wide

  range of uses.

  Mechanism for the polymerisation of ethene.

  1. Initiation

  It is a free radical initiation.

  

  2. Propagation

  

  3. Termination 

  

  where the part between brackets indicates a unit of the formula of the polymer

  that repeats itself in the formula; n indicates the number of the units in a formula of

  a polymer and is a very large number.

  Summary of most alkene polymers obtained from alkenes as monomers and their

  uses (Table 3.1)

  

  

  3.7. Structure, classification and nomenclature of alkynes

  
  

  
  

  A triple bond consists of one sigma bond and two pi bonds. Each carbon of the triple

  bond uses two sp orbital to form sigma bonds with other atoms. The unhybridised 2p

  orbitals which are perpendicular to the axes of the two sp orbitals overlap sideways

  to form pi bonds.

  According to the VSEPR model, the molecular geometry in alkynes include bond

  angle of 180^o around each carbon triply bonded.Thus, the shape around the triple

  bond is linear.

  
  

  There are two types of alkynes: terminal alkynes and non-terminal (internal)

  alkynes

  A terminal alkyne has a triple bond at the end of the chain e.g.: : R-C ≡ C − H

  A non-terminal alkyne has a triple bond in the middle of the chain: R − C ≡ C − R'

  Examples

  

  Alkynes are named by identifying the longest continuous chain containing the

  triple bond and changing the ending –ane from the corresponding alkane to –yne

  
  

  
  

  
  

  
  

  3.8. Laboratory and industrial preparation of alkynes

  1. Preparation of ethyne

  Activity: 3.8

  Set up the apparatus as shown in the diagram below. 

  
  

  
  

  
  

  Procedure:

  • Place 2g of calcium carbide in a conical flask

  • Using the dropping funnel, add water drop by drop.

  • Collect the gas produced in the test tube.

  • Remove the first tube and connect a second test tube.

  • To the first test tube add two drops of bromine water. Record your observations

  • To the second tube add two drops of potassium manganate (VII). Record your observations.

  
  

  Ethyne (acetylene) can be prepared from calcium carbide which is obtained by

  reduction of calcium oxide by coke at high temperature.

  

  

  When bromine water is added to acetylene, the red colour of bromine is discharged.

  The solution becomes colourless.

  The decolourisation of bromine water is a test for unsaturation in a compound.

  When potassium manganate (VII) is added to acetylene, its purple colour is

  discharged.

  2. Alkylation of acetylene

  The hydrogen atom of ethyne as that of other terminal alkynes is slightly acidic and therefore it can be removed by a strong base like NaNH₂ or KNH₂.The products of the reaction are acetylides. Acetylides react with halogenoalkanes to yield higher alkynes.

  

  3. Dehydrohalogenation

  The dehydrohalogenation of vicinal or geminal dihalogenoalkanes yields alkynes

  

  4. Dehalogenation

  The dehalogenation of a tetrahalogenoalkane yield an alkyne.

  

  3.9. Physical properties of alkynes

  
  

  Alkynes are non-polar compounds with physical properties similar to those of

  alkenes with the same number of carbon atoms. Their linear structure gives them

  greater intermolecular forces than alkenes

  
  

  
  

  
  

  3.10. Chemical reactions of alkynes

  
  

  Addition reactions

  As unsaturated hydrocarbons, alkynes are very reactive. Because they are unsaturated

  hydrocarbons, alkynes undergo addition reactions. Alkynes can add two moles of reagents.

  Even though they have a higher electron density than alkenes, they are in general less reactive because the triple bond is shorter and therefore the electron cloud is less accessible.

  1. Addition of hydrogen halides

  Alkynes react with hydrogen halides to yield vicinal dihalogenoalkanes, the reaction follows the Markownikov’s rule. The reaction takes place in four steps.

  
  

  2. Addition of water

  Alkynes react with water in the presence of sulphuric acid and mercury sulphate at 60^o C to give carbonyl compounds.

  
  

  3. Hydrogenation

  The hydrogenation of alkynes in the presence of palladium catalyst gives alkanes

  The reaction requires two moles of hydrogen for a complete saturation.

  

  In the presence of Lindlar catalyst, the alkynes are partially hydrogenated giving alkenes4

  
  

  A Lindlar catalyst is a heterogeneous catalyst that consists of palladium deposited on

  calcium carbonate and poisoned with different lead derivatives such as lead oxide

  or lead acetate. A heterogeneous catalyst is the one which is in the phase different

  from that of the reactants.

  4. Reaction with metals

  Terminal alkynes react with active metals to yield alkynides and hydrogen gas.

  Internal alkynes do not react as they do not have an acidic hydrogen atom.

  

  5. Reaction with metal salts

  When a terminal alkyne is passed through a solution of ammoniacal silver nitrate, a white precipitate of silver carbide is formed.

  

  When a terminal alkyne is passed through a solution of ammoniacal copper(I)

  chloride, a red precipitate of copper(I)carbide is formed. 

  

  The reactions above are used to:

  • Differentiate between terminal and non-terminal alkynes.

  • Differentiate ethene and ethyne

  The reaction shows that hydrogen atoms of ethyne are slightly acidic, unlike those

  of ethene.

  

  3.11. Uses of alkenes and alkynes

  Activity 3.11

  Look at the picture below and appreciate the importance of alkenes and alkynes.

  

  • Alkenes are extremely important in the manufacture of plastics which have

  many applications such as: packaging, wrapping, clothing, making clothes,

  artificial flowers, pipes, cups, windows, ...

  • Ethene is a plant hormone involved in the ripening of fruits, seed germination,

  bud opening;

  

  • Ethene derivatives are also used in the making of polymers such as polyvinylchloride (PVC), Teflon,...

  • Alkenes are used as raw materials in industry for the manufacture of alcohols, aldehydes, ...

  • Alkynes are used in the preparation of many other compounds. For exampleethyne is used in the making of ethanal, ethanoic acid, vinyl chloride, trichloroethane, ...

  • Ethyne (acetylene) is used as a fuel in welding and cutting metals.

  • Propyne is used as substitute for acetylene as fuel for welding.

  

  

  

  

  

  

  

  

• Unit 4 : HALOGENOALKANES (ALKYL HALIDES)

  UNIT 4: HALOGENOALKANES (ALKYL HALIDES)

  Key unit competency

  The learner should be able to relate the physical and chemical properties of halogenoalkanes to their reactivity and their uses

  Learning objectives

  • Define halogenoalkanes and homologous series.

  • Explain the reactivity of halogenoalkanes.

  • Explain the physical properties of halogenoalkanes.

  • Describe preparation methods for halogenoalkanes.

  • Explain different mechanisms in halogenoalkanes.

  • Explain the uses and dangers associated with halogenoalkanes.

  • Draw displayed structural formulae of halogenoalkanes and give names

  using IUPAC system.

  • Classify halogenoalkanes according to developed formula as primary,

  secondary and tertiary.

  • Write reaction mechanisms of halogenoalkanes as SN1, SN2, E1 and E2.

  • Test for the presence of halogenoalkanes in a given sample organic

  compound.

  • Appreciate the uses and dangers of halogenoalkanes in everyday life.

  • Develop the awareness in protecting the environment.

  • Develop team work approach and confidence in group activities and

  presentation sessions.

  

  4.1. Definition and nomenclature of halogenoalkanes

  
  

  
  

  1. Definition

  Halogenoalkanes compounds are compounds in which the halogen atoms like

  chlorine, bromine, iodine or fluorine are attached to a hydrocarbon chain. When

  the halogen atom is attached to a hydrocarbon chain the compound is called a halogenoalkane or haloalkane or an alkyl halide. 

  Halogenoalkanes contain halogen atom(s) attached to the sp³ hybridised carbon atom of an alkyl group. 

  
  

  2. Nomenclature of halogenoalkanes

  Halogenoalkanes are organic compounds that contain a halogen atom: F, Cl, Br, I.

  They are named using the prefixes fluoro-, chloro-, bromo- and iodo-.

  Numbers are used if necessary to indicate the position of the halogen atom in the

  molecule.

  

  4.2. Classification and isomerism

  
  

  
  

  
  

  4.2.1. Classification of halogenoalkanes

  There are three types of halogenoalkanes:
  

  

  A primary halogenoalkane has a halogen atom attached to the ended carbon atom

  of the chain. A secondary halogenoalkane has a halogen atom attached to a carbon

  bonded to two other carbon atoms while a tertiary halogenoalkane has a halogen

  atom attached to a carbon bonded to three other carbon atoms.

  

  4.2.2. Isomerism

  Halogenoalkanes exhibit both chain and position isomerism.

  Example: Molecular formula C₄H₉Br

  a. Chain isomerism: This arises due to arrangement of carbon atoms in chains of

  different size.

  

  b. Position isomerism: This arises due to the different positions taken by the

  halogen atom on the same carbon chain.

  The following compounds are position isomers: CH₃ CH₂ CH₂ CH₂-Br and CH₃ CH₂ CH Br CH₃; because the atoms of bromine are on different positions of the chain.

  Hence, all isomers of the compound with molecular formula C₄H₉ Br are the following.

  

  

  4.3. Physical properties of halogenoalkanes

  
  

  
  

  
  

  1. Volatility

  Volatility is a property that shows if a substance transforms easily or not into vapour

  or gaseous form. This property depends on the nature of the bonds that make up

  the molecule of the substance. Generally non polar covalent compounds are more

  volatile than polar covalent compounds. We know that halogens when bonded to

  other atoms form polar bonds because they possess high electronegativities: F =

  4.0, Cl = 3.0, Br = 2.8, I = 2.5, and C = 2.5.

  
  

  The more the difference of electronegativities of the atoms that form the bond,

  the more polar is the bond. This explains the high polarity of C-F bond with an

  electronegativity difference of 1.5, and the low polarity of C-Cl and C-Br bonds where

  the electronegativity differences are 0.5 and 0.3 respectively.

  The presence of polarity or charge distribution results into more attraction between

  polar molecules called dipole-dipole attraction forces, one type of Van der Waals

  forces, as shown below:

  
  

  The dashed line represents the attraction forces between the polar molecules or

  dipoles.

  Therefore, more energy must be supplied to separate polar molecules and this

  explains why melting and boiling temperatures of fluoroalkanes and chloroalkanes

  are higher than those of alkanes of similar molecular mass.

  As we have already learnt, molecules of organic halogen compounds are generally

  polar. Due to the greater polarity as well as higher molecular mass as compared

  to the parent hydrocarbons, the intermolecular forces of attraction (dipole-dipole

  and Van der Waals) are stronger in the halogen derivatives. That is why the boiling

  points of chlorides, bromides and iodides are considerably higher than those of the

  hydrocarbons of comparable molecular mass (Table 4.1).

  

  Chloromethane, bromomethane, chloroethane and some chlorofluoromethanes

  are gases at room temperature. Higher members are liquids or solids.

  The attractions get stronger as the molecules get bigger in size. The pattern of

  variation of boiling points of different halides is depicted in Figure 4.1. For the same

  alkyl group, the boiling points of alkyl halides increase in the order: RF <RCl < RBr, <

  RI This is because with the increase in size and mass of halogen atom, the magnitude of Van der waal forces increases.

  2. Solubility

  The solubility is the capacity of a substance to dissolve in a given solvent; in chemistry the most common solvent we refer to is water. It is a result of the interaction between the molecules of the substance, a solute, and the molecules of the solvent.

  Polar molecules can interact with water molecules, but the attractive forces set

  up between water molecules and molecules concerned are not as strong as the

  hydrogen bonds present in water. Halogenoalkanes therefore, although they

  dissolve more than alkanes, are only slightly soluble in water.

  3. State

  The state of matter is the physical appearance of that matter: solid, liquid and

  gaseous.

  Chloromethane, bromomethane, chloroethane and chloroethene are colourless

  gases at room temperature and pressure. The higher members are colourless

  liquids with a sweet pleasant smell

  4. Density

  The density is a measure of the quantity of matter by volume unit. Cotton wool is

  less dense than sand because if you compare the quantity of matter cotton wool

  and sand contained in for instance 1m³, you find that there more matter in sand than in cotton wool.

  The density of halogenoalkanes increases in the order RCl < RBr < RI, since the

  atomic weight of halogens increases in order Cl < Br < I. Iodo, bromo and polychloro

  derivatives are denser than water but chloro derivatives are less dense than water.

  

  

  4.4. Preparation methods of halogenoalkanes

  1. From alkenes and alkynes

  

  
  

  

  

  Direct halogenation of alkanes in the presence of ultraviolet light gives alkyl halides and a hydrogen halide.

  

  

  
  

  
  

  
  

  
  

  

  

  

  
  

  

  

  
  

  
  

  

  

  

  
  

  
  

  

  

  
  

  

  
  

  

  
  

  

  

  
  

  

  

  
  

  

  
  

  

  
  

  

  

  
  

  
  
• UNIT 5: ALCOHOLS AND ETHERS

  UNIT 5: ALCOHOLS AND ETHERS

  
  

  Key unit competency:
  

  To be able to compare the physical and chemical properties of alcohols and ethers to their preparation methods, reactivity and uses.

  
  

  Learning objectives:

  • Distinguish between alcohols from other organic compounds by representing the functional group of alcohols

  • Classify primary, secondary and tertiary alcohols by carrying out the method of identification

  • Write the name of alcohols by using IUPAC system

  • Describe the physical properties of alcohols to other series of organic compounds

  • Carry out the method of preparation of alcohols

  • Describe the local process of making alcohol by fermentation.

  • Explain the effect of oxidation on urwagwa when it overstays

  • Compare the physical, chemical and the method of preparation of alcohols to ethers

  • State the use of ethers

  

  
  

  5.1. Definition and nomenclature

  
  

  
  

  
  

  5.1.1. Definition
  

  Alcohols are organic compounds that are derivatives of hydrocarbons where one or more hydrogen atoms of hydrocarbon is or are replaced by hydroxyl (-OH) group. They are represented by the general formula: C_nH_2n+1OH or ROH where R is a radical: alkyl group made by a chain of carbon atoms
  

  
  

  Alcohols are called monohydric if only one hydroxyl group is present (eg: CH3CH2-OH). Dihydric alcohols are those with two hydroxyl group (diol: vicinal and gem), trihydric (triols) and polyhydric are those with many – C-OH groups. The functional group attached is –OH group to any atom of carbon.
  

  
  

  5.1.2. Nomenclature

  According to IUPAC system, alcohols are named by replacing the final ‘‘e’’ of the parent hydrocarbon with ‘‘ol’’, then specify the position of -OH group before ending by ol
  

  

  
  

  
  

  5.2. Classification and isomerism

  
  

  
  

  Alcohols are classified as:
  

  Primary alcohols: These have only one alkyl group attached to the carbon carrying the –OH.

  
  

  
  

  
  

  • Functional isomers:

  Except methanol which has one carbon, other alcohols are isomers with ethers

  another chemical function of general formula R-O-R’ where R and R’are alkyl groups or aryl groups but not hydrogen.

  
  

  
  

  
  

  5.3. Physical properties

  
  

  
  

  a.Boiling points

  
  

  The chart shows the boiling points of some simple primary alcohols and alkanes with up to 4 carbon atoms.

  
  

  • The boiling point of an alcohol is always much higher than that of the alkane with the same number of carbon atoms.

  • The boiling points of the alcohols increase as the number of carbon atoms increases.

  • The boiling point of alcohols with branches is lower than that of unbranched alcohols with the same number of carbon atoms. This is because increased branching gives molecules a nearly spherical shape and the surface area of contact between molecules in the liquid. This results in weakened intermolecular forces and therefore in lower boiling points.

  • Tertiary alcohols exhibit the lowest boiling point than secondary and primary alcohols:

  • Primary alcohol                      > Secondary alcohol              > Tertiary alcohol
  

                                                  Highest boiling point                    lowest boiling point

  
  

  
  

  The patterns in boiling point reflect the patterns in intermolecular attractions: In the case of alcohols, there are hydrogen bonds set up between the slightly positive hydrogen atoms and lone pairs on oxygen in other molecules.

  
  

  
  

  
  

  b.Solubility of alcohols in water

  
  

  The lower members of alcohols are completely soluble in water because mixed hydrogen bonds between water and alcohol molecules are formed. As the length of hydrocarbon group of the alcohol increases, the solubility decreases.

  
  

  c. Volatility

  
  

  Alcohols are volatile and the volatility decreases as the molecular mass increases. Compared to alkyl halides, alcohols are less volatile. Polyalcohol are viscous or solids. Example: propane-1, 2, 3-triol (glycerine). This is due to stronger intermolecular forces than those of monoalcohols.

  
  

  
  

  5.4. Alcohol preparations

  
  

  b. From alkenes

  Alkenes react with water in the presence concentrated sulphuric acid to yields

  alcohols

  
  

  c. From carbonyl compounds

  When aldehydes and ketones are reduced by hydrogen in the presence of a

  suitable catalyst like Pt, Ni or Pd, they form primary and secondary alcohols

  respectively.

  

  

  d. From esters

  Esters on hydrolysis in the presence of mineral acid or alkalis produce alcohols and

  carboxylic acids.

  

  e. From Grignard reagents

  The reaction between carbonyl compound and Grignard reagent (alkyl magnesium

  halides) produces an alcohol with more carbon atoms. The reaction is a nucleophilic addition on a carbonyl compound.

  

  f. From primary amine to give primary alcohol

  Primary amines react with nitrous acid to produce primary alcohols.

  

  5.5. Preparation of ethanol by fermentation

  
  

  
  

  This method is mainly used to prepare ethanol industrially. Ethanol is prepared from starch
  

  (e.g. maize, cassava, millet, sorghum) and sugar(e.g. banana juice, molasses) by fermentation process.

  
  

  Fermentation can be defined as any of many anaerobic biochemical reactions in which enzymes produced by microorganisms catalyse the conversion of on substance into another.

  
  

  
  

  The ethanol obtained by fermentation process is only about 11%. This is made concentrated by distillation which converts it to about 95% ethanol. This on further distillation yields a constant boiling mixture whose composition does not change (an azeotropic mixture). Therefore, 100% ethanol is obtained by either:

  
  

  i. Adding quick lime which removes water

  ii. Distilling with of benzene as a third component

  Note: Methanol can be prepared industrially by the reaction of carbon monoxide

  and hydrogen at 300 °C and a pressure of 200 atmospheres.

  
  

  
  

  
  

  
  

  
  

  5.6.1. Oxidation 

  
  

  
  

  Aldehydes formed by oxidation of primary alcohols tend to undergo further oxidation to carboxylic acid.

  
  

  Ketones formed by oxidation of secondary alcohols are not further oxidised, unless if the oxidising agent is hot and concentrated in which case bonds around the –CO_ group are broken and two smaller carboxylic acids are formed.

  
  

  
  

  
  

  
  

  Tertiary alcohols resist oxidation because they have no hydrogen atom attached on the functional carbon atom. Oxidation also occurs when the alcohol is in gaseous phase by used of silver or copper catalyst under 500 °C and 300 °C respectively; and the vapour of the alcohol is passed with air (oxygen) over heated silver.

  
  

  
  

  An acidified potassium dichromate solution is turned from orange to green when it reacts with primary and secondary alcohols.

  
  

  Secondary alcohols having the following structure R-CHOH-CH₃ only undergo oxidation, on treatment with iodine solution in the presence of sodium hydroxide to give yellow precipitate of tri-iodomethane.
  

  
  

  Note: This is a reaction which is characteristic of methyl ketones, CH₃-CO-R’; but iodine here acting as an oxidizing agent first oxidizes the CH₃-CHOH-R’toCH₃-CO-R’ ; then the methyl ketone formed then gives the yellow precipitate of CHI₃ (Iodoform). From the reaction involved we have the Iodoform test.

  
  

  
  

  5.6.2. Reaction with sulphuric acid

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  5.6.4. Reaction with strong electropositive metals and metal hydroxides

  
  

  
  

  
  

  5.6.5 Action of hydrohalic acids (HX)
  

  
  

  
  

  
  

  
  

  Notice:

  i. Reaction with concentrated hydrochloric acid is catalyzed by anhydrous zinc chloride.

  
  

  ii. This reaction is called LUCAS test and is used to distinguish between simple primary, secondary or tertiary alcohols. In this reaction, the alcohol is shaken with a solution of zinc chloride in concentrated hydrochloric acid.
  

  
  

  Observations: Immediate cloudiness indicates presence of a tertiary alcohol. If the solution becomes cloudy within 5 minutes then the alcohol is a secondary one. Primary alcohol would show no cloudiness at room temperature since the reaction is very slow.

  
  

  For example all alcohols which are isomers of C₄H₁₀O can be distinguished by the LUCAS test.
  

  
  

  Alcohols are also transformed into halogenoalkanes using phosphorus halides and thionyl chloride.

  
  

  
  

  
  

  
  

  
  

  5.7. Uses of alcohols

  
  

  
  

  Ethanol is the alcohol found in alcoholic drinks. Alcoholic fermentation converts starch sugar into ethanol. For example grapes are used to produce wine, ripe banana to produce urwagwa, honey for spirits are obtained by distilling the ethanol –water product obtained when sugar is fermented.

  
  

  Drinking alcohol, i.e. the ethylic alcohol also called ethanol, is a normal social activity; but excess of it is dangerous for our health. Hence excess of alcoholic consumption must be avoided.For non-adult youth, consumption of alcohol in any form is illegal in Rwanda and many other countries.

  
  

  There are some alcoholic drinks produced in Rwanda and in the Region that are prohibited to be sold in Rwanda. However, alcohols have many other applications in daily life as indicated in the Table 5.1.
  

  
  

  

  
  

  

  
  

  Ethanol produced by sugar cane fermentation has been used as alternative fuel to gasoline (petrol). It has been mixed with gasoline to produce gasohol.

  
  

  

  
  

  

  
  

  

  
  

  1. Ethers are sparingly soluble in water but are soluble in organic solvents.

  2. The polar nature of the C-O bond (due to the electronegativity difference of the atoms) results in intermolecular dipole-dipole interactions.

  3. An ether cannot form hydrogen bonds with other ether molecules since there is no H to be donated (no -OH group).

  4. Their melting and boiling points increase with the increase in molecular mass because of increasing the magnitude of Van der Waal’s forces with size.

  5. The boiling points of ethers are much lower than those of alcohols of similar molecular mass. This is because of the intermolecular hydrogen bonding which are present in alcohols but are not possible in ethers.
  

  
  

  

  
  

  

  
  

  

  
  

  Since they are saturated compounds and non-polar, they are relatively chemically inert reason why their chemical reactions are very few.

  
  

  

  
  

  

  
  

  a. Ethers can act as the Lewis base due to the two non-bonded electron pair on oxygen to form coordinative bonds with Grignard reagent. This explains clearly why organ magnesium compounds are manipulated in ether solvent but not in water since in water, there is a reaction which generate alkanes.
  

  
  

  

  
  

  
  

  

  
  

  Lower ethers are used as anesthesia since they produce inert local cooling when sprayed on a skin, ether are also used as local anesthesia for minor surgery operation.Lower ethers are volatile liquid which on evaporation produce low temperature they are therefore used as refrigerants.

  
  

  Ether itself is one of the most important organic solvents for fats, oils, resins, and alkaloids.

  
  

  

  
  

  
• Unit 6 : CARBONYL COMPOUNDS: ALDEHYDES AND KETONES

  UNIT 6: CARBONYL COMPOUNDS: ALDEHYDES AND KETONES

  Key unit competency

  To be able to compare the chemical nature of carbonyl compounds to their

  reactivity and uses.

  Learning objectives

  • Describe the reactivity of carbonyl compounds

  • State the physical properties of aldehydes and ketones

  • Describe the preparation reactions of ketones and aldehydes

  • Explain the mechanism of nucleophilic addition reactions of carbonyl

  compounds

  • Prepare ketones from secondary alcohols by oxidation reactions

  • Compare aldehydes and ketones by using Fehling’s solution and Tollens’

  reagent

  • Write and name carbonyl compounds and isomers of ketones and aldehydes

  • Write equations for the reactions of carbonyl compounds with other

  substances

  • Compare the physical properties of carbonyl compounds to those of alcohols

  and alkenes

  • Differentiate the methyl ketones from other ketones by using the iodoform

  test

  • Carry out an experiment to distinguish between carbonyl compounds and

  other organic compounds

  • Carry out an experiment to distinguish between ketones and aldehydes

  • Carry out an experiment to prepare ethanol and propan-2-one.

  6.1. Definition and nomenclature of carbonyl compounds

  
  

  Introductory activity

  Many fruits such as mangoes and honey contained sugar.The following images

  represent mangoes, honey and some sugars such as fructose and glucose. 

  

  6.1.1 Definition

  
  

  
  

  Carbonyl compounds are compounds that contain carbon-oxygen double bond

  (C=O). Carbonyl compounds are classified into two general categories based on the

  kinds of chemistry they undergo. In one category there are aldehydes and ketones;

  in the other category there are carboxylic acids and their derivatives. This unit looks

  on category of aldehydes and ketones.

  
  

  Aldehyde molecules

  For aldehydes, the carbonyl group is attached to hydrogen atom and alkyl group as

  shown in the molecule of propanal below. Methanal is the smallest aldehyde, it has

  two hydrogen atoms attached to carbonyl group.

  

  If you are going to write this in a condensed form, you write aldehyde as –CHO,

  don’t write it as -COH, because that looks like an alcohol functional group.

  Ketone molecules

  
  

  
  

  6.1.2 Nomenclature

  Aldehydes

  
  

  The systematic name of an aldehyde is obtained by replacing the terminal “e” from

  the name of the parent hydrocarbon with “al.”In numbering the carbon chain of an

  aldehyde, the carbonyl carbon is numbered one.

  
  

  
  

  Ketones

  The systematic name of a ketone is obtained by removing the terminal“e” from the

  name of the parent hydrocarbon and adding “one.”The chain is numbered in the

  direction that gives the carbonyl carbon the smallest number.Ketone contains

  a carbon-oxygen double bond just like aldehyde, but for ketone carbonyl groupis

  bonded to two alkyl groups. 

  

  

  6.2. Isomerism

  

  6.2.1 Functional isomerism in aldehydes and ketones

  
  

  Isomers are molecules that have the same molecular formula, but have a different

  arrangement of the atoms in space. Functional group isomers are molecules that

  have same molecular formula but contain different functional groups, and they

  belong to different homologous series of compounds.

  

  6.2.2. Position isomerism in ketones 

  Position isomerism is isomerism where carbon skeleton remains constant, but the

  functional group takes different positions on carbon skeleton.

  
  

  
  

  6.2.3. Chain isomerism in aldehydes and ketones

  In chain isomerism the same number of carbons forms different skeletons.

  Aldehydes with 4 or more carbon atoms and ketones with five or more carbon

  atoms show chain isomerism. 

  
  

  
  

  6.3. Physical properties of aldehydes and ketones

  
  

  6.3.1. Solubility in water aldehydes and ketones

  
  

  The small molecules of aldehydes and ketones are soluble in water but solubility

  decreases with increase of carbon chain. Methanal, ethanal and propanone - the

  common small aldehydes and ketones are soluble in water at all proportions.

  Even though aldehydes and ketones don’t form hydrogen bond with themselves,

  they can form hydrogen bond with water molecules.

  
  

  The slightly positive hydrogen atoms in a water molecule can be sufficiently

  attracted to the lone pair on the oxygen atom of an aldehyde or ketone to form a

  hydrogen bond.

  Other intermolecular forces present between the molecules of aldehyde or ketone

  and the water are dispersion forces and dipole-dipole attractions

  Forming these attractions releases energy which helps to supply the energy

  needed to separate the water molecules and aldehyde or ketone molecules from

  each other before they can mix together.

  Apart from the carbonyl group, hydrocarbon chains are non polar, they don’t

  dissolve in water. By forcing hydrocarbon chain to mix with water molecules, they

  break the relatively strong hydrogen bonds between water molecules without

  replacing them by other attractions good like hydrogen bonds. This makes the

  process energetically less profitable, and so solubility decrease.

  6.3.2. Boiling points of aldehydes and the ketones

  Methanal is a gas and has a boiling point of -21°C, and ethanal has a boiling point

  of +21°C. The other aldehydes and ketones are liquids or solids, with boiling points

  rising with rising of molecular mass hence rising of strength of Van der Waals force. 

  Comparing the physical properties of carbonyl compounds to those of alcohols

  and alkanes

  
  

  Physical properties of covalent compounds depend on intermolecular forces.

  Compounds that have similar molecular mass but different intermolecular forces

  have different physical properties.

  

  Alcohols have higher boiling point than aldehydes and ketones of similar lengths. In

  the alcohol, there is hydrogen bonding, but the molecules of aldehydes and ketones

  don’t form hydrogen bonds.Aldehydes and ketones are polar molecules but alkanes

  are non polar molecules. 

  
  

  
  

  6.4. Chemical properties of carbonyl compounds

  6.4.1. Nucleophilic addition reactions

  
  

  
  

  a. Polarity of carbonyl group

  By comparing carbon-carbon double bond and carbon- oxygen double bond

  the only difference between bonds C=C and C=O is distribution of electrons. The

  distribution of electrons in the pi bond is heavily attracted towards the oxygen atom,

  because oxygen atom is much more electronegative than carbon.

  
  

  During chemical reactions nucleophiles will attack carbon of the carbonyl functional

  group which bears apartial positive charge. While electrophile will attack oxygen of

  the carbonyl functional group which bears a partial negative charge.

  b. Reaction of HCN with aldehydes and ketones

  Hydrogen cyanide adds to aldehydes or ketones to form cyanohydrins or

  hydroxynitriles.The product has one more carbon atom than the reactant. For

  example, ethanal reacts with HCN to form 2-hydroxypropanenitrile:

  

  Because hydrogen cyanide is a toxic gas, the best way to carry out this reaction is to

  generate hydrogen cyanide during the reaction by adding HCl to a mixture of the

  aldehyde or ketone and excess sodium cyanide. Excess sodium cyanide is used in

  order to make sure that some cyanide ion is available to act as a nucleophile. The

  solution will contain hydrogen cyanide (from the reaction between the sodium or

  potassium cyanide and the HCl)

  The pH of the solution is maintained in range 4 - 5, because this gives the fastest

  reaction. The reaction takes place at room temperature.

  c. The mechanism of reaction between HCN and propanone

  
  

  
  

  d. Application of the reaction

  The product of the reaction above has two functional groups:

  • The -OH group which behaves like ordinary alcohol and can be replaced by

  other substituent like chlorine, which can in turn be replaced to give other

  functional group, for example, an -NH₂

   group;

  • The -CN group which can be hydrolysed into a carboxylic acid functional

  group -COOH.

  
  

  
  

  e.Reaction of NaHSO3 with aldehydes or ketones

  The aldehyde or ketone is shaken with a saturated solution of sodium hydrogen

  sulphite in water. Hydrogen sulphite with negative charge act as nucleophile,

  where the product formed is separated as white crystals. Propanone react hydrogen

  sulphite, as below:

  

  Impure aldehyde and ketone can be purified by using this reaction. Impure

  aldehyde or ketone is shaken with a saturated solution of sodium hydrogensulphite

  to produce the crystals. Impurities don’t form crystals; these crystals formed are

  filtered and washed to remove any impurities. Addition of dilute acid to filtered

  crystals regenerates the original aldehyde. Dilute alkali also can be added instead

  dilute acid. 

  

  

  6.4.2. Condensation reactions

  
  

  
  

  a. Experimental reaction

  The procedure of the preparation of Brady’s reagent and carbonyl compounds

  changesslightly depending on the nature of the aldehyde or ketone, and the solvent

  in which 2,4-dinitrophenylhydrazine is dissolved in. The Brady’s reagent for activities

  (6.4.1) is a solution of the 2,4-dinitrophenylhydrazine in methanol and sulphuric acid.

  Add a few drops of Brady’s reagent to either aldehyde or ketone. A bright orange or

  yellow precipitate indicates the presence of the carbonyl group in an aldehyde or

  ketone. 

  b. Structural formula of 2,4-dinitrophenylhydrazine.

  The carbon of benzene attached to hydrazine is counted as number one.In

  2,4-dinitrophenylhydrazine, there are two nitro groups, NO₂,attached to the phenyl

  group in the 2- and 4- positions.

  
  

  c. The reaction of carbonyl compounds with 2,4-dinitrophenylhydrazine

  Brady’s reagent is a solution of the 2,4-dinitrophenylhydrazine in methanol

  and sulphuric acid. The overall reaction of carbonyl compounds with

  2,4-dinitrophenylhydrazine is:

  

  Where R and R’ represent alkyl groups or hydrogen(s); if both or only one is hydrogens

  the starting carbonyl compound is an aldehyde. If both R and R’ are alkyl groups

  the carbonyl compound is a ketone. The following molecule shows clearly how the

  product is formed.

  

  The product formed is named”2,4-dinitrophenylhydrazone”. The simple difference

  consists in replacing suffix “-ine” by “-one”.

  The reaction of 2,4-dinitrophenylhydrazine with ethanal produces ethanal

  2,4-dinitrophenylhydrazone; The reaction of 2,4-dinitrophenylhydrazine with

  butanal produces butanal 2,4-dinitrophenylhydrazone. This is an example of

  condensation reaction.

  During the chemical reaction, the change takes place only on nitrogen (-NH₂) of

  hydrazine in 2,4-dinitrophenylhydrazine. If the -NH₂ group is attached to other

  groups a similar reaction as that of 2,4-dinitrophenylhydrazine will take place: 

  
  

  
  

  
  

  6.4.3. Oxidation reactions using KMnO₄/H+ and K₂Cr₂O₇/H+

  
  

  
  

  a. Difference in reactivity of ketones and aldehydes with K₂Cr₂O₇

  
  

  By considering the structural formulae of aldehydes and ketones, the difference is

  only the presence of a hydrogen atom attached to the carbonyl functional group in

  the aldehyde whereas ketones have a alkyl group instead.

  
  

  
  

  During chemical reaction aldehydes react with oxidizing agent; hydrogen on

  carbonyl functional group is replaced by oxygen, look on figure below. The presence

  of hydrogen atom makes aldehydes very easy to oxidize, in other words aldehydes

  are strong reducing agents.

  For ketone, absence of hydrogen on carbonyl functional group makes ketones to

  resist oxidation. But very strong oxidising agents like potassium permanganate

  solution oxidize ketones - and they do it in a destructive way, by breaking carboncarbon bonds.

  Aldehyde oxidation can take place in acidic or alkaline solutions. Under acidic

  solutions, the aldehyde is oxidized to a carboxylic acid. Under alkaline solutions, acid

  formed react with base to form a salt of carboxylic acid. 

  

  b. Oxidation of aldehyde by K₂Cr₂O₇/H+ solution

  
  

  Add few drops of the aldehyde or ketone to a solution of potassium dichromate

  (VI) acidified with dilute sulphuric acid. If the color doesn’t change in the cold, the

  mixture is warmed gently in a beaker containing hot water.

  
  

  6.4.4. Oxidation reactions using Tollens’ reagent

  
  

  a. Difference in reactivity of Ketones and Aldehydes with Tollens’ reagent

  Aldehydes can also be oxidized into carboxylic ions in basic medium.Tollens’ reagent

  is a solution of diamminesilver (I) ion, [Ag(NH₃)₂]⁺ and OH-.In order to identify if a

  substance is aldehyde or ketone, add few drops of Tollens reagent to test tubes

  containing aldehyde or ketone and warm gently in a hot water bath for a few minutes. The formations of sliver mirror or grey precipitate is an indication of the presence of aldehyde.

  
  

  
  

  6.4.5. Oxidation reactions using Fehling ;or Benedict; solution

  

  a. Difference in reactivity of Ketones and Aldehydes with Fehling or Benedict

  solution.

  Fehling’s solution and Benedict’s solution react with aldehyde in the same way; both

  solutions contain Cu²⁺ and OH^- . In Fehling’s solution Cu²⁺ is complexed with tartrate

  ligand butin Benedict’s solution Cu²⁺ is complexed with citrate ligand.

  
  

  Don’t worry about ligands, important reagents are Cu²⁺ and OH^- , ligands tartrate and

  citrate are used to prevent formation of precipitate copper (II) hydroxide or copper

  (II) carbonate.

  
  

  A few drops of Fehling’s solution or Benedict’s solution is added to the aldehyde or

  ketone and the mixture is warmed gently in a hot water bath for a few minutes.

  
  

  
  

  
  

  6.4.6. Iodoform reaction with aldehydes and ketones

  Activity 6.4.6

  Materials:

  
  

  
  

  a. Reagents for iodoform reaction

  There are two different mixtures that can be used to do iodoform test, these

  mixture are:

  • Iodine and sodium hydroxide solution

  • Potassium iodide and sodium chlorate (I) solutions

  Don’t worry about Potassium iodide and sodium chlorate(I) solutions, Potassium

  iodide and sodium chlorate(I) react to form final solution containI₂ and OH-. Both mixtures contain the same reagents.

  Each of these mixtures contains important reagent I₂ and OHwhich react with

  aldehyde or ketone. When I₂ and OH^-  is added to a carbonyl compound containing the group CH₃ CO (blue in the cycle) as shown below, pale yellow precipitate (triiodomethane) is formed.

  
  

  
  

  a. Description of iodoform test

  For iodine and sodium hydroxide solution

  Iodine solution, I₃^- , is added to aldehyde or ketone, followed by just enough sodium

  hydroxide solution to remove the colour of the iodine. If pale yellow precipitate

  doesn’t form in the cold, it may be necessary to warm the mixture very gently. The positive result is pale yellow precipitate of CHI₃ 

  For potassium iodide and sodium chlorate (I) solutions

  Potassium iodide solution is added to a small amount of aldehyde or ketone, followed

  by sodium chlorate (I) solution. If pale yellow precipitate doesn’t form in the cold,

  warm the mixture very gently. The positive result is pale yellow precipitate of CHI₃.

  Reaction of iodoform test

  The reagents of iodoform test are I₂ and OHsolution. The reaction takes place into

  two main steps:

  • Three hydroxides, OH^- , remove three hydrogens from methyl group and the

  place of hydrogen is taken by iodide.

  
  

  
  

  
  

  6.5. Preparation methods of aldehydes and ketones

  6.5.1. Oxidation of alcohols 

  
  

  
  

  a. alcohol by K₂\Cr₂O₇/H+

  Potassium dichromate (VI) acidified with dilute sulphuric acid is used as oxidizing agent during the preparation of aldehyde or ketone. Primary alcohol is oxidized to  aldehyde, oxygen atom from the oxidising agent removes two hydrogens; one from the -OH group of the alcohol and the other hydrogen comes from the carbon that is attached to hydroxide functional group .

  

  

  b.Technique of stopping oxidation of aldehyde

  The aldehyde produced by oxidation of alcohol could make further oxidation to a

  carboxylic acid if the acidified potassium dichromate (VI) is still present in solution

  where reaction takes place. In order to prevent this further oxidation of aldehyde to

  carboxylic the following technique are used.

  • Use an excess of the alcohol than potassium dichromate (VI). Potassium

  dichromate (VI) is limiting reactant hence there isn’t enough oxidising agent

  present to carry out the second stage of oxidizing the aldehyde formed to a

  carboxylic acid.

  • Distil off the aldehyde as soon as it forms. Removing the aldehyde as soon as it

  is formed this means that aldehyde is removed from solution where oxidizing

  agent is, to prevent further oxidation. Ethanol produces ethanal as shown by

  the following reaction.

  

  

  c. Oxidation of alkene by KMnO₄/H⁺

  

  

  6.5.2. Preparation of ketone by distillation of calcium acetate

  
  

  Procedure: Transfer 15g of calcium acetate in 50ml round bottom flask fixed on

  a stand, and place it on a heating mantle fitted with a condenser and a receiver

  flask. Adjust the temperature until the condensation starts. Use the aluminium

  foil to insulate the flask. Heat the flask and collect the acetone in receiver flask. The

  obtained product is a crude acetone and needs to be purified.Set up a distillation

  apparatus and distil the crude product to obtain pure acetone (56^Oc). Do not forget

  to use stirrer bar which must be placed in the round bottom flask containing the

  acetone.

  
  

  6.6. Uses of aldehydes and ketones

  Aldehydes and ketones have many uses for example in industries such as

  pharmaceutical industry and in medicine.

  
  

  a. Formaldehyde:

  Formaldehyde is a gas at room temperature but is sold as a 37 percent solution in water.

  Formaldehyde is used as preservative and germicide, fungicide, and insecticide

  for plants and vegetables. Formaldehyde is mainly used in production of certain

  polymers like Bakelite (Figure 6.1). Bakelite and formaldehyde is used as

  monomers in production of Bakelite

  

  b. Acetone as solvent: 

  Acetone  is soluble in water at all proportions and also dissolves in many organic

  compounds. Boiling point of acetone is low, 56 °C, which makes it easier to be removed by

  evaporation. Acetone is an industrial solvent that is used in products such as paints,

  varnishes, resins, coatings, and nail polish removers.

  c. Aldehydes and ketones

  Organic molecules that contain ketones or aldehydes functional group are found

  in different foods such as irish potatoes, yellow bananas.

  Aldehydes and ketones perform essential functions in humans and other living organisms. For examples sugars, starch, and cellulose, which are formed from simple molecules that have aldehyde or ketone functional group

  d. Aldehydes and ketones in human’s body

  
  

  
  

  Aldehydes and ketones functional group are found in humans hormones like progesterone,

  testosterone. 

  

  

  

• Unit 7 : CARBOXYLIC ACIDS AND ACYL CHLORIDES

  UNIT 7: CARBOXYLIC ACIDS AND ACYL CHLORIDES

  Key unit competency:

  The learner should be able to compare the chemical nature of the carboxylic acids

  and acid halides to their reactivity.

  Learning objectives

  • Explain the physical properties and uses of carboxylic acids and acyl chlorides

  • Describe the inductive effect on the acidity of carboxylic acid

  • Explain the reactions of carboxylic acids and acyl chlorides

  • Apply the IUPAC rules to name different carboxylic acids acyl chlorides

  • Write the structural formula and isomers of carboxylic acids

  • Distinguish between carboxylic acids from other organic compounds using

  appropriate chemical test

  • Prepare carboxylic acids from oxidation of aldehydes or primary alcohols

  • Compare the physical properties of carboxylic acids to those of alcohols

  • Outline the mechanisms of esterification and those of reaction of acyl

  chlorides with ammonia , amines and alcohols

  • Develop a culture of working as a team group activities and self-confidence

  in presentation

  • Appreciate the uses of carboxylic acids as the intermediate compounds in

  industrial processes such as aspirin, vinegar and perfumes

  

  Carboxylic acid is classified in the family of organic compounds due to the presence

  of carboxyl group (-COOH) in their chemical formula. The general formula for

  carboxylic acids is R-COOH where R- refers to the alkyl group of the molecule.

  

  

  7.1.1. Nomenclature

  
  

  Carboxylic acids are named by following the general rules of naming organic

  compounds, where the suffix ‘oic acid’ is added to the stem name of the longest

  carbon chain that contains the acid functional group. The side branches are also

  positioned by starting from the carbon with carboxylic functional group.

  The carboxylic group takes priority to other functional group when numbering

  carbons in thecase of substituted chain.

  

  

  

  • Optical isomers

  Optical isomers have the same molecular formula and the same structural formula,

  but they are different in the spatial arrangement of atoms and their optical properties.

  An organic compound shows optical isomerism, when there is chiral carbon (a

  carbon atom attached to four diverse groups) in its structure. A chiral carbon is also

  known as asymmetric carbon. 

  

  

  

  Just as the right hand and left hand are mirror images of another but not superimposable, optical isomers, also known as enantiomers, are different from each other and can have different properties. For example, muscles produce D-lactic acid when they contract, and a high amount of this compound in muscles causes muscular pain and cramps.

  These molecules are optical isomers, because they have opposite optical activities.They can be distinguished by a plane-polarized light where one enantiomer rotates the light to the right while the other rotates it to the left.

  

  Enantiomers are often identified as D- or L- prefixes because of the direction in which

  they rotate the plane polarized light as shown in figures 7.2 and 7.3. Enantiomers

  that rotate plane polarized light in clockwise direction are known as dextrorotatory

  (right-handed) molecules and enantiomers that rotate plane polarized light in

  anticlockwise direction are known as levorotatory (left-handed) molecules.

  A solution containing equal amounts of enantiomers, 50% levorotatory and 50%

  dextrorotatory is known as a racemic mixture that will not rotate polarized light,

  because the rotations of the two enantiomers cancel each other out. 

  
  

  

  7.2. Physical properties of carboxylic acids

  
  

  
  

  a. Physical state

  Many carboxylic acids are colorless liquids with disagreeable odors. Aliphatic

  carboxylic acids with 5 to 10 carbon atoms are all liquids with a “goaty” odors (odor

  of cheese). These acids are also produced by the action of skin bacteria on human

  sebum (skin oils), which accounts for the odor of poorly ventilated storerooms.

  The acids with more than 10 carbon atoms are wax-like solids, and their odor

  diminishes with increasing molar mass and resultant decreasing volatility.

  Anhydrous acetic acid freezes at (17^oC) slightly below ordinary room temperature,

  reason why it is called glacial acetic acid (Figure 7.4). But a mixture of acetic acid

  with water solidifies at much lower temperature.

  

  b. Melting and boiling point

  Carboxylic acids show a high degree of association through hydrogen bonding.

  Because of this, they have high melting and boiling points compared to other

  organic compounds of the same mass or number of carbon atoms.

  
  

  
  

  Carboxylic acids have high melting and boiling points because their hydrogen

  bonds enhance the possibility of bringing two acid molecules together by forming

  a kind of dimer.

  

  

  

  7.3. Acidity of carboxylic acids

  
  

  
  

  
  

  Solutions of carboxylic acid turn blue litmus paper red; they do not change the

  color of red litmus paper; therefore, they are acids as other mineral acids such as

  HCl (aq).

  Organic or carboxylic acids are weak acids in opposition to some mineral acids such

  as hydrochloric acids which are strong acids.According to Arrhenius’ theory of acids

  and bases, strong acids dissociate completely in water to give hydrogen ion, H⁺(aq) or

  H₃O⁺, whereas weak acids dissociate partially. The hydrogen ion released combines with a water molecule to form H₃O⁺ a hydrate positive ion called hydronium H₃O⁺:

  

  The carboxylate ion formed by ionization of the acid is more stable than the acid

  because it has many resonance structures.

  Ethanoic acid is a weaker acid than methanoic because its methyl group has a

  positive inductive effect; that is to mean that it pushes electrons towards the O-H

  bond hence make hydrogen ion stable and not easily leaving.

  

  The greater the number of such groups, the greater the effected and therefore

  the weaker will be the acid.For example, 2,2-dimethylpropanoic is weaker than

  2-methylpropanoic acid which is in turn weaker than propanoic.

  The same rule applies to the increase in the length of the alkyl group chain.Butanoic

  acid is a weaker acid than propanoic acid which shows that the acidity strength

  decreases as the alkyl chain increases.

  On the other hand, when an electron withdrawing group (a group with a negative

  inductive effect) is present, the opposite effect is observed. For example,

  chloroethanoic acid is a stronger acid than ethanoic acid. This is because chlorine

  being electronegative, will withdraw electron towards itself thus reducing the

  electron density around the O-H bond thus weakening it. It causes O-H bond to

  easily break, and the concentration of hydrogen ions will be high in the solution.

  

  The more the number of groups with negative inductive effect, the greater is the

  effect and hence the more acidic will be the solution. Trifluoroacetic acid is more

  acidic than trichloroacetic, dichloroacetic, chloroacetic and acetic acid because

  fluorine is more electronegative than chlorine and hydrogen. It will strongly

  withdraw electron towards itself, hence makes easier for the proton to leave.It must

  also be noted that the further away the electronegative element, the less the effect.

  For example, 3-chlorobutanoic acid is therefore a weaker acid than 2-chlorobutanoic

  acid.

  

  7.4. Preparation of carboxylic acids

  
  

  
  

  
  

  Carboxylic acids are common and vital functional group; found in amino acids,

  fatty acids etc. and provide the starting raw material for acid derivatives such as

  acyl chlorides, amides, esters and acid anhydrides. There are several methods

  of preparation of carboxylic acids where the most common are discussed in this

  section.

  
  

  7.4.1. From primary alcohols and aldehydes

  
  

  Different carboxylic acids can be prepared by oxidation of either primary alcohols

  or aldehydes. In the process, the mixture of alcohol is heated under reflux with an

  oxidizing agent such acidified potassium permanganate or potassium dichromate.

  Primary alcohols are first oxidized to aldehydes then further oxidation of aldehydes

  produces carboxylic acid.

  
  

  
  

  7.4.2. Hydrolysis of acid nitriles and amides with acid or alkali

  When nitriles are hydrolyzed by water in acidic medium and the mixture is submitted

  to heat, the reaction yields carboxylic acids.

  
  

  
  

  
  

  7.4.3. From dicarboxylic acid

  Monocarboxylic acids can be prepared by heatingcarboxylic acids which have two

  carboxylic functional groups attached to the same carbon atom.

  
  

  Note that the reaction is used to reduce length of the carbon chain. The mono

  carboxylic acid prepared has one carbon atom less than the starting dicarboxylic

  acid.

  7.4.4. From organomagnesium compounds (Carboxylation reaction)

  Grignard reagents react with carbon dioxide gas, and when the intermediate compound formed is hydrolyzed it finally forms carboxylic acid.

  

  Mechanism:

  
  

  
  

  7.4.5. From alkenes (Oxidation of alkenes)

  Carboxylic acids are also obtained by heating alkenes with concentrated acidified

  potassium permanganate. The reaction unfortunately forms a mixture of compounds

  that mustbe later separated.

  
  

  Note that the hydrolysis of carboxylic acid derivatives such as amide, esters, acyl

  chloride and acid anhydrides also produce the corresponding acids.

  7.4.6. Laboratory preparation of acetic acid

  
  

  
  

  
  

  7.5. Reactions of carboxylic acids

  
  

  
  

  
  

  
  

  
  

  
  

  The reaction of acids with carbonates is the basis for the chemical test of carboxylic

  acid functional group and it can be used to distinguish carboxylic acids from other

  functional groups in qualitative analysis.

  Carbon dioxide produced is also tested by lime water and it turns lime water milky

  (Figure 7.8)

  
  

  
  

  It is noted that the oxygen atom in the ester formed comes from the alcohol and the

  one in water is from the acid. In the mechanism of esterification, the acid loses -OH

  group while the alcohol loses H-atom.

  

  7.5.3. Reduction of carboxylic acids

  Carboxylic acids are reduced to primary alcohols on treatment with reducing agent

  such as LiAlH₄ in dry ether or by use of hydrogen in the presence of Ni catalyst. The

  reduction does not form aldehyde as an intermediate product, like in oxidation of

  primary alcohols.

  
  

  
  

  
  

  
  

  
  

  7.6. Uses of carboxylic acids

  
  

  
  

  
  

  Food industry and nutrition

  • Food additives: Sorbic acid, benzoic acid, etc.

  • Main ingredient of common vinegar (acetic acid).

  • Elaboration of cheese and other milk products (lactic acid).

  Pharmaceutical industry

  • Antipyretic and analgesic (acetylsalicylic acid or aspirin).

  • Active in the process of synthesis of aromas, in some drugs (butyric or butanoic acid).

  • Antimycotic and fungicide (Caprylic acid and benzoic acid combined with salicylic acid).

  • Active for the manufacture of medicines based on vitamin C (ascorbic acid).

  • Manufacture of some laxatives (Hydroxybutanedioic acid).

   Other industries

  • Manufacture of varnishes, resins and transparent adhesives (acrylic acid).

  • Manufacture of paints and varnishes (Linoleic acid).

  • Manufacture of soaps, detergents, shampoos, cosmetics and metal cleaning products (Oleic acid).

  • Manufacture of toothpaste (Salicylic acid).

  • Production of dyes and tanned leather (Methanoic acid).

  • Manufacture of rubber (Acetic acid).

  • Preparation of paraffin candles (Stearic acid)

  
  

  7.7. Acyl chlorides and nomenclature 

  

  Acyl halides are compounds with the general formula where the–OH group of

  carboxylic acid has been substituted by a halogen atom. The acyl remaining

  structure is represented as:

  

  Their isomers can be chain isomerism, positional isomerism and functional isomerism

  with chloro aldehydes and ketones, alcohols with double bond C=C and chlorine as

  a substituent, cyclic ethers with chlorine.

  Acid chlorides have not many applications in our everyday life, but industrially they

  are used in synthesis of perfumes and nylons, which are polymers of high importance

  in textile industry. They can also be used in pharmaceutical industries to synthesize

  drugs with aromatic ester or amide functional groups like aspirin or paracetamol.

  

  7.7.1. Physical properties

  
  

  
  

  Appearance

  Acyl chlorides are colourless fuming liquids. Their characteristic strong smell is

  caused by hydrogen chloride gas that is produced when they get in contact with

  moisture (see figure 7.7). For example, the strong smell of ethanoyl chloride is a

  mixture of vinegar odour and the acrid smell of hydrogen chloride gas.

  
  

  b. Solubility

  Acyl chlorides are slightly soluble in water due to their small dipole that can interact

  with the polarity of water molecule. They cannot be said to be soluble in water

  because they readily react with water. It is impossible to have a simple aqueous

  solution of acyl chlorides, rather we have the products of their reaction with water.

  
  

  c. Boiling and melting points

  Acyl chloride molecules interact by Van der Waals forces whose strength increases

  with the increase in molecular masses of the compounds. 

  
  

  The boiling and melting points of acyl chlorides increases as their molecular masses

  rise. They have lower boiling and melting points than alcohols and carboxylic acids

  of the same number of carbon atoms, because they lack hydrogen bonds.

  

  7.7.2. Reactions of acyl chlorides

  
  

  The chemistry of acyl chlorides is dominated by nucleophilic substitution, where

  a stronger nucleophile replaces chlorine atom of acyl chloride. They undergo

  nucleophilic substitution reactions more easily than alkyl halides and carboxylic

  acids because the nucleophile targets the carbon which is deficient in electrons and

  -Cl is better leaving group than -OH group.

  
  

  The common reactions of acyl chlorides include reactions with water, alcohols and

  ammonia and amines. These reactants have a very electronegative element that

  has a free lone pair of electrons to act as a nucleophile.

  
  

  
  

  
  

  Reaction with alcohols

  They react with alcohol to produce esters with high yields than esterification of an

  alcohol and carboxylic acid, since Cl-atom in acyl chloride is a better leaving group

  than O-H for the case of carboxylic acid. The difference in electronegativity is the

  main reason for this observation.

  
  

  
  

  Reaction with ammonia and amine

  
  

  Acyl chlorides react with ammonia and amines to yield amides. Ammonia, primary

  amines and secondary amines form primary amides, secondary amides and tertiary

  amides respectively.

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

• UNIT 8: ESTERS, ACID ANHYDRIDES, AMIDES AND NITRILES.

  UNIT 8: ESTERS, ACID ANHYDRIDES, AMIDES AND ITRILES.

  Key unit competency:

  To be able to relate the functional groups of esters, acid anhydrides, amides and

  nitriles to their reactivity, preparation methods and uses. 

  
  
  

  8.1. Structure and nomenclature of esters 

  8.1.1. Structure of esters

  In unit 7, the reactions of carboxylic acids were discussed. The reactions of carboxylic acids produce the derivatives of acids such as esters, acid halides, acid anhydrides and amides.

  The general molecular formula of esters is Cn H₂nO₂  and their general structural

  formula is: RCOOR’ or

  
  

  
  

  

  Esters are known for their distinctive odor and they are commonly responsible for

  the characteristic of food (fruits) aroma, flowers and fragrances. Esters are found in

  nature but they can be also synthesized. Both natural and synthetic esters are used

  in perfumes and as flavoring agents.

  8.1.2. Nomenclature of esters

  The nomenclature of esters follows some steps. When naming esters the alkyl

  group R’ is named followed by the name of RCOO- group.

  The group name of the alkyl or aryl portion is written first and is followed by the

  name of the acid portion. In both common and International Union of Pure and

  Applied Chemistry (IUPAC) nomenclature, the -ic ending of the corresponding acid

  is replaced by the suffix –ate. Some examples of names of esters are given in Table

  8.1.

  
  

  8.1.3. Physical properties and uses of Esters
  

  

  
  

  B. Comparing boiling points of alcohols, carboxylic acids and esters

  Materials and Chemicals

  Propan-1-ol, propanoic acid and methyl ethanoate, test tubes, test tube holders

  (lacks), heaters, and thermometers.

  Procedure

  1. Put 10 mL of each substance in a labeled test tube.

  2. Boil carefully substances are volatile and flammable

  3. Use a thermometer to measure the boiling point of each substance.

  4. Record the results and compare them. Suggest a reason for the difference

  in boiling points of the three substances.

  Conclusion: Esters have lower boiling points than alcohols and carboxylic

  acids because they lack hydrogen bonds. A compound having hydrogen bonds

  has a high boiling point because, to break that bond requires higher energy. 

  Other physical properties of esters

  i. Lower esters have sweet fruity smells

  ii. Melting and boiling points of esters increase as the molecular mass

  increases.

  iii. Small esters are fairly soluble in water but the solubility decreases as the

  length of the chain increases

  8.1.4. Uses of Esters

  Esters find various uses:

  i. They are used as organic solvent

  ii. Due to their aroma, they are used as constituent of fragrance, essential oils,

  food flavoring and cosmetics.

  iii. They are used to manufacture soaps, detergents and glycerol.

  iv. They are used to provide energy in the body

  v. Polyesters are used to produce plastics etc.

  
  

  8.2. Preparation and chemical properties of esters
  

  
  

  8.2.1. Preparation of Esters

  The preparation of esters involves different types of reaction such as esterification,

  reaction of an acid chloride with an alcohol and the reaction of acid anhydrides

  with alcohols.

  1.Esterification reaction

  In units five and seven, it is mentioned that esters can be produced by a reaction

  between alcohols and carboxylic acids in strong acidic medium acting as a catalyst.

  The acid is commonly a concentrated sulphuric acid, under reflux (Figure 8.3). The

  reaction is generally called “Esterification” (a condensation reaction which involves

  the addition of the alcohol and acid molecules followed by an elimination of a

  water molecule).

  

  

  
  

  

  
  

  

  

  
  

  

  

  

  

  

  
  

  

  

  

  

  

  

  

  

  

  
  
• Unit 9 : AMINES AND AMINO ACIDS

  UNIT 9: AMINES AND AMINO ACIDS

  Key unit competency:

  The learner should be able to relate the chemical nature of the amines and

  aminoacids to their properties, uses and reactivity.

  Learning objectives

  At the end of this unit, the students will be able to:

  • Explain the zwitterion forms in the solution of different pH.

  • Explain the isoelectric point in amino acids.

  • Describe the physical properties and uses of amines.

  • Describe the preparation methods of the amines.

  • Describe the reactions of amino acids and amines with other substances.

  • Classify amines as primary, secondary and tertiary amines.

  • Compare and contrast the physical properties of the amino acids to those of

  carboxylic acids and amines.

  • Test the presence of amines and amino acids in the solution. 

  
  

  
  

  9.1. Nomenclature and classification of amines

  

  Amines are one of organic compounds containing nitrogen. They are one of the

  most important classes of organic compounds which are obtained by replacing

  one or more hydrogen atoms by an alkyl or aryl group in a molecule of ammonia

  (NH₃). They are present in vitamins, proteins, hormones, etc. They are extensively

  used in the manufacturing of many drugs and detergents.

  
  

  9.1.1 Classification of amines

  Nitrogen has 5 valence electrons and so is trivalent with a lone pair. As per VSEPR

  theory, nitrogen present in amines is sp³ hybridized and due to the presence of lone pair, it is pyramidal in shape instead of tetrahedral shape which is a general structure for most sp3

   hybridized molecules. Each of the three sp³  hybridized orbitals of nitrogen overlap with orbitals of hydrogen or carbon depending upon the configuration of amines. Due to the presence of lone pair, the C-N-H angle in amines is less than 109 degrees which is characteristic angle of tetrahedral geometry. The angle in amines is near about 108 degrees.

  

  

  9.1.2 Nomenclature of amines

  In organic chemistry, the names of the compounds are given according to the

  guidelines provided by IUPAC. In this regards, amines are named by ending with –

  amine. The IUPAC system names amine functions as substituents on the largest alkyl

  group. 

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  9.2 Physical properties, natural occurrences and uses of amines.

  
  

  
  

  9.2.1 Physical properties of amines

  
  

  Tertiary amines do not bond to each other by hydrogen bond and they have boiling

  points similar to those of hydrocarbons of the same molecular weight. However,

  primary, secondary and tertiary amines form hydrogen bond with water and amines

  with low-molecular weight are generally soluble in water.

  
  

  Generally the boiling point of amines increases as the molecular weight increase

  and they boil at higher temperatures than alkanes but at lower temperatures than

  alcohols of comparable molar mass.

  The amines are soluble in organic solvent and the solubility decreases as the

  molecular weight increases. The Table 9.3 summarizes some physical properties of

  some amines.

  

  9.2.2 Natural occurrence of amines and their usage

  Natural amines occur in proteins, vitamins, hormones, etc. and they are also prepared

  synthetically to make polymers, drugs and dyes.

  Amines can be used as dyes (colorants) or as drugs: Primary aromatic amines are

  used as a starting material for the manufacture of azo dyes. They react with nitrous

  (II) acid to form diazonium salt which can undergo a coupling reaction in order to

  form an azo compound. As azo compounds are highly coloured, they are widely

  used in dyeing industries. Examples include Methyl orange and Direct brown 138.

  In medicine, amines can be used as drugs. 

  • Chlorpheniramine is an antihistamine that helps to relief allergic disorders

  due to cold, hay fever, itchy skin, insect bites and stings.

  • Diphenhydramine is the common antihistamine.

  • Chlorpromazine is a tranquillizer that anaesthetizes without inducing sleep.

  It is used to relieve anxiety, excitement, restlessness or even mental disorder.

  • Acetaminophen is also known as paracetamol or p-acetaminophenol, it is

  an analgesic that relieves pains such as headaches. It is believed to be less

  corrosive to the stomach and is an alternative to aspirin.

  Amines are widely encountered in biological and pharmacological studies. Some

  important examples are the 2-phenylethylamines, some vitamins, antihistamines,

  tranquilizers, and neurotransmitters (noradrenaline, dopamine and serotonin)

  which act at neuromuscular synapses.

  

  9.3 Preparation of amines.

  

  The amines can be prepared based on the following reactions:
  

  9.3.1 Alkylation of ammonia

  
  

  9.3.2. Gabriel phthalimide synthesis

  This procedure is used for the preparation of primary amines. Phthalimide on

  treatment with ethanolic potassium hydroxide forms potassium salt of phthalimide

  which on heating with alkyl halide followed by alkaline hydrolysis produces the

  corresponding primary amine. However, primary aromatic amines cannot be

  prepared by Gabriel phthalimide synthesis because aryl halides do not undergo

  nucleophilic substitution with the anion formed by phthalimide.

  
  

  
  

  
  

  9.3.3. Hoffmann bromamide degradation reaction

  Hoffmann developed a method for the preparation of primary amines by treating an

  amide with bromine in an aqueous or ethanolic solution of sodium hydroxide. This

  is a degradation reaction with migration of an alkyl or aryl group taking place from

  carbonyl carbon of the amide to the nitrogen atom.

  The reaction is valid for the preparation of primary amines only, and it yields

  uncontaminated compound with other amines. 

  

  9.3.4 Reduction of amides

  Similarly to reduction of amides, lithium aluminium hydride (LiAlH₄) reduces amides to amines. 

  
  

  9.3.5 Reduction of nitriles

  Nitriles are reduced to amines using hydrogen in the presence of a nickel catalyst, although acidic or alkaline conditions should not be used to avoid the possible hydrolysis of the -CN group. LiAlH₄ is more commonly employed for the reduction of nitriles on the laboratory scale.

  
  

  9.3.6 Reduction of nitro compounds

  Nitro compounds are reduced to amines by passing hydrogen gas in the presence

  of finely divided nickel, palladium or platinum and also by reduction with metals

  in acidic medium. Nitroalkanes can also be similarly reduced to the corresponding

  alkanamines.

  
  

  
  

  Reduction with iron scrap and hydrochloric acid is preferred because FeCl₂ formed gets hydrolysed to release hydrochloric acid during the reaction. Thus, only a small amount of hydrochloric acid is required to initiate the reaction.

  
  

  
  

  9.4. Chemical properties of amines

  
  

  
  

  Difference in electronegativity between nitrogen and hydrogen atoms and the

  presence of unshared pair of electrons over the nitrogen atom makes amines

  reactive. The number of hydrogen atoms attached to nitrogen atom also is involved

  in the reaction of amines; that is why the reactivity of amines differ in many reactions.

  Amines behave as nucleophiles due to the presence of unshared electron pair.

  The chemical properties of amines are summarized in the reactions below.

  
  

  9.4.1 Reactions of amines diluted with acids

  Amines, like ammonia, are bases. Being basic in nature, they react with acids to form salts.

  

  9.4.2 Reactions of amines (alkylation, acylation, and sulfonation)

  Acyl chlorides and acid anhydrides react with primary and secondary amines to

  form amides. Tertiary amines cannot be acylated due to the absence of a replaceable

  hydrogen atom. 

  

  9.4.3. Reaction with carboxylic acid

  Because amines are basic, they neutralize carboxylic acids to form the corresponding

  ammonium carboxylate salts. Upon heating at 200°C, the primary and secondary

  amine salts dehydrate to form the corresponding amides.

  

  9.4.4. Reaction with nitrous acid

  Nitrous acid, HNO₂ is unstable. It is produced indirectly using a mixture of NaNO₂ and a strong acid such as HCl or H₂SO₄ in diluted solution. Primary aliphatic amines react with nitrous acid to produce a very unstable diazonium salts which spontaneously decomposes by losing N₂ to form a carbenium ion. Further, the carbonium ion is used to produce a mixture of alkenes, alkanols or alkyl halides, with alkanols as major product said above

  

  Primary aromatic amines, such as aniline (phenylamine) forms a more stable

  diazonium ion at 0^oC –5°C. Above 5°C, it will decompose to give phenol and N₂. Diazonium salts can be isolated in the crystalline form but are usually used in solution and immediately after preparation, due to its rapid decomposition.

  
  

  
  

  9.4.5. Reactions with ketones and aldehydes

  Primary amines react with carbonyl compounds to form imines. Specifically,

  aldehydes become aldimines, and ketones become ketimines. In the case of

  formaldehyde (R’ = H), the imine products are typically cyclic trimers.

  
  

  Secondary amines react with ketones and aldehydes to form enamines. An

  enamine contains a C=C double bond, where the second C is singly bonded to N as

  part of an amine ligand.

  

  9.4.6. Neutralization reactions

  Tertiary amines (R₃N) react with strong acids such as hydroiodic acid (HI), hydrobromic acid (HBr) and hydrochloric acid (HCl) to give ammonium salts R₃NH⁺X^-.

  
  

  
  

  
  

  
  

  9.5. General structure of amino acids and some common examples

  
  

  9.5.1. General structure of amino acids

  Amino acids are organic compounds containing amine (-NH₂) and carboxyl (-COOH)

  functional groups, along with a side chain (R group) specific to each amino acid.

  The key elements of an amino acid are carbon (C), hydrogen (H), oxygen (O), and

  nitrogen (N). About 500 naturally occurring amino acids are known.

  The general structure of amino acid is shown by the functional group (-NH₂) and

  a carboxylic acid group (-COOH) attached to the same carbon and they are called

  α-amino acids.

  
  

  The R group is the part of the amino acid that can vary in different amino acids. It

  can be a hydrogen (in that case, the amino acid is called Glycine) or a –CH₃ group (Alanine) or other radicals.

  

  9.5.2. Common Amino Acids

  Among the 500 known amino acids, there are 20 important α-amino acids, as shown

  in the Table 9.4 below. Each amino acid has a common name. You will notice that

  the names in common used for amino acids are not descriptive of their structural

  formulas; but at least they have the advantage of being shorter than the systematic

  names. The abbreviations (Gly, Glu, …) that are listed in table below, are particularly

  useful in designating the sequences of amino acids in proteins and peptides. 

  

  

  

  

  

  The first amino acid to be isolated was asparagine in 1806. It was obtained from

  protein found in asparagus juice (hence the name). Glycine, the major amino acid

  found in gelatin, was named for its sweet taste (Greek glykys, meaning “sweet”). In

  some cases an amino acid found in a protein is actually a derivative of one of the

  common 20 amino acids.

  

  9.6. Comparison of physical properties of amino acids to those of carboxylic acids and amines

  
  

  
  

  The amino acids, carboxylic acids and amines have different functional groups; this

  is the base of their different physical properties as shown in the Table 9.5.

  
  

  
  

  The first amino acid to be isolated was asparagine in 1806. It was obtained from protein found in asparagus juice (hence the name). Glycine, the major amino acid found in gelatin, was named for its sweet taste (Greek glykys, meaning “sweet”). In some cases an amino acid found in a protein is actually a derivative of one of the common 20 amino acids.

  
  

  9.7. Chemical properties of amino acids

  
  

  The reactivity of amino acids involves the reactions of both amines and carboxylic

  acids. Some of these reactions are given below.

  
  

  
  

  
  

  
  

  
  

  9.7.1. Acid–base properties of amino acids

  As the name suggests, amino acids are organic compounds that contain both a

  carboxylic acid group and an amine group. Amino acids are crystalline, high melting

  point (>200°C) solids. Such high melting points are unusual for a substance with

  molecules of this size — they are a result of internal ionisation. Even in the solid

  state, amino acids exist as zwitterions in which a proton has been lost from the

  carboxyl group and accepted by the nitrogen of the amine group:

  
  

  So instead of hydrogen bonds between the amino acid molecules there are stronger

  ionic (electrovalent) bonds. This is reflected in the relative lack of solubility of amino

  acids in non- aqueous solvents compared with their solubility in water.

  
  

  Zwitterions exhibit acid–base behaviour because they can accept and donate

  protons. In acids a proton is accepted by the carboxylic acid anion, forming a unit

  with an overall positive charge:

  
  

  In alkalis the reverse occurs with the loss of a proton from the nitrogen atom: 
  

  
  

  Carboxylic acids have acidic properties and react with bases. Amines have basic

  properties and react with acids. It therefore follows that amino acids have both

  acidic and basic properties.

  9.7.2. Isoelectric point in aminoacids (pI)

  The isoelectric point (pI), is the pH at which a particular molecule carries no net

  electrical charge in the statistical mean. This means it is the pH at which the amino

  acid is neutral, i.e. the zwitterion form is dominant. The pI is given by the average of

  the pKa that involve the zwitterion, i.e. that give the boundaries to its existence.

  The table below shows the pKa values and the isoelectronic point, pI, are given

  below for the 20 α-amino acids (Table 9.6).

  pKa₁= α-carboxyl group, pKa₂= α-ammonium ion, and pKa₃  = side chain group

  
  

  
  

  There are 3 cases to consider: 

  1. Neutral side chains

  These amino acids are characterised by two pKa values: pKa₁ and pKa₂ for the

  carboxylic acid and the amine respectively. The isoelectronic point will be halfway

  between, or the average of, these two pKa values . This is most readily appreciated

  when you realise that at very acidic pH (below pKa₁) the amino acid will have an

  overall positive charge and at very basic pH (above pKa₂) the amino acid will have

  an overall negative charge. 

  

  The other two cases introduce other ionisable groups in the side chain “R” described

  by a third acid dissociation constant, pKa₃

  _{2. Acidic side chains}

  The pI will be at a lower pH because the acidic side chain introduces an “extra” negative charge. So the neutral form exists under more acidic conditions when the extra -ve has been neutralised. For example, for aspartic acid shown below, the neutral form is dominant between pH 1.88 and 3.65, pI is halfway between these two values, i.e. pI = 1/2 (pKa₁ + pKa₃ ), so pI = 2.77

  

  3. Basic side chains

  The pI will be at a higher pH because the basic side chain introduces an “extra”

  positive charge. So the neutral form exists under more basic conditions when the

  extra positive has been neutralised. For example, for histidine, which has three acidic

  groups of pKa’s 1.82 (carboxylic acid), 6.04 (pyrrole NH) and 9.17 (ammonium NH),

  the neutral form is dominant between pH 6.04 and 9.17; pI is halfway between these

  two values, i.e. , so pI = 7.60. 

  
  

  _{9.7.3. Reaction with strong acids} 

  _{In the following reaction, amino acids react with strong acids such as hydrochloric acid:}

  9.7.4. Reaction with nitrous acid (deamination) 

  The amine function of α-amino acids and esters reacts with nitrous acid in a similar manner to that described for primary amines. The diazonium ion intermediate loses molecular nitrogen in the case of the acid, but the diazonium ester loses a proton and forms a relatively stable diazo compound known as ethyl diazoethanoate:

  

  The diazo ester is formed because of the loss of N₂ from the diazonium ion which

  results in the formation of a quite unfavourable carbocation.

  9.7.5. Reaction with sodium hydroxide

  Amino acids react with strong bases such as sodium hydroxide:
  

  
  

  9.7.6. Reaction of amino acids with sodium carbonate

  Amino acids are instantly dissolved by strong hydrochloric acid but are in part

  recovered unchanged on dilution and evaporation. They are not decomposed by

  sodium carbonate but are easily decomposed by sodium hydroxide. (Dakin & West,

  1928).

  
  

  9.8. Optical isomers of amino acids

  
  

  In chemistry, the term “isomer” means molecules that have the same molecular

  formula, but have a different arrangement of the atoms in space.

  Simple substances which show optical isomerism exist as two isomers known as

  enantiomers. Where the atoms making up the various isomers are joined up in a

  different order, this is known as structural isomerism. Structural isomerism is not

  a form of stereoisomerism, which involve the atoms of the complex bonded in the

  same order, but in different spatial arrangements. Optical isomerism is one form of

  stereoisomerism; geometric isomers are a second type.

  

  The general formula for an amino acid (apart from glycine, 2-aminoethanoic acid)

  is shown below.The carbon at the centre of the structure has four different groups

  attached. In glycine, the “R” group is another hydrogen atom.

  

  The lack of a plane of symmetry means that there will be two stereoisomers of an

  amino acid (apart from glycine) - one the non-superimposable mirror image of the

  other.

  For a general 2-amino acid, the isomers are:

  The R group, usually referred to as a side chain, determines the properties of each

  amino acid. Scientists classify amino acids into different categories based on the

  nature of the side chain. A tetrahedral carbon atom with four distinct groups is

  called chiral. The ability of a molecule to rotate plane polarized light to the left,

  L (levorotary) or right, D (dextrorotary) gives it its optical and stereo chemical

  fingerprint. 

  All the naturally occurring amino acids have the right-hand structure in the

  diagram above. This is known as the “L-” configuration. The other one is known as

  the “D-” configuration.

  When asymmetric carbon atoms are present in a molecular compound, there are

  two ways in which the groups attached to that carbon can be arranged in the three

  dimensions, as we have just shown with the two models above. Chemically, optical

  isomers behave in the same way. Biologically, they do not. One will react properly,

  but the other will not.

  

  9.9. Peptides and polypeptides

  
  

  
  

  9.9.1. Formation of peptide bonds

  
  

  Amino acid molecules can also react with each other; the acidic –COOH group in one

  molecule reacts with the basic –NH₂ group in another molecule. When two amino

  acids react together, the resulting molecule is called a dipeptide, forming an amide

  linkage (peptide bond), with the elimination of a water molecule.

  
  

  Each amino acid possesses a carboxylic acid group and an amine group. The

  possibilities for constructing polypeptides and proteins are enormous. Let us

  consider two simple amino acids, glycine (2-aminoethanoic acid) and alanine

  (2-aminopropanoic acid). The figures below show that these can be joined in two

  ways:

  
  

  Note the amide link between the two amino acids. An amide link between two

  amino acid molecules is also called a peptide link. The reaction is a condensation

  reaction as a small molecule. The dipeptide product still has an –NH₂ group at one

  end and a –COOH group at the other end. Therefore the reaction can continue, to

  form a tripeptide initially, and then ever-longer chains of amino acids. The longer

  molecules become known as polypeptides, and then proteins as they get even

  longer sequences of amino acids. A typical protein is formed from between 50 and

  200 amino acids joined in a variety of sequences.

  9.9.2. Structure of peptides and polypetides

  
  

  A series of amino acids joined by peptide bonds form a polypeptide chain, and each

  amino acid unit in a polypeptide is called a residue. A polypeptide chain has polarity

  because its ends are different, with α-amino group at one end and α-carboxyl

  group at the other. By convention, the amino end is taken to be the beginning of

  a polypeptide chain, and so the sequence of amino acids in a polypeptide chain

  is written starting with the aminoterminal residue. Thus, in the pentapeptide TyrGly-Gly-Phe-Leu (YGGFL), tyrosine is the amino-terminal (N-terminal) residue and

  leucine is the carboxyl-terminal (C-terminal) residue Leu-Phe-Gly-Gly-Tyr (LFGGY) is

  a different pentapeptide, with different chemical properties.

  
  

  This above illustration of the pentapeptide Tyr-Gly-Gly-Phe-Leu (YGGFL) shows

  the sequence from the amino terminus to the carboxyl terminus. This pentapeptide,

  Leu-enkephalin, is an opioid peptide that modulates the perception of pain. The

  reverse pentapeptide, Leu-Phe-Gly-Gly-Tyr (LFGGY), is a different molecule and

  shows no such effects.

  A polypeptide chain consists of a regularly repeating part, called the main chain

  or backbone, and a variable part, comprising the distinctive side chains. The

  polypeptide backbone is rich in hydrogen-bonding potential. Each residue contains

  a carbonyl group, which is a good hydrogen-bond acceptor and, with the exception

  of proline, an NH group, which is a good hydrogen-bond donor. These groups

  interact with each other and with functional groups from side chains to stabilize

  particular structures, as will be discussed later.

  

  A polypeptide chain consists of a constant backbone (shown in blue) and variable

  side chains (shown in green).

  9.9.3. Uses of amino acids as building blocks of proteins

  
  

  Like carbohydrates and lipids, proteins contain the elements carbon (C), hydrogen

  (H) and oxygen (O), but in addition they also always contain nitrogen (N).sulphur(S)

  is often present as well as iron (Fe) and phosphorus (P). Before understanding how

  proteins are constructed, the structure of amino acids should be noted. 

  
  

  The process of construction of proteins begins by amino acids bonding together, as

  seen earlier, through peptide bonds. When many amino acids join together a long-chain polypeptide is produced. The linking of amino acids in this way takes place

  during protein synthesis.

  The simplest level of protein structure, primary structure, is simply the sequence of

  amino acids in a polypeptide chain. The primary structure (Figure 9.2) of a protein

  refers to its linear sequence of amino de (–S–S–) bridges. One of those sequences is:

  –Gly–Ile–Val–Cyst–Glu–Gln–Ala–Ser–Leu–Asp–Arg–Asp–Arg–Cys–Val–Pro–

  The primary structure is held together by peptide bonds that are made during the

  process of protein biosynthesis. The two ends of the polypeptide chain are referred

  to as the carboxyl terminus (C-terminus) and the amino terminus (N-terminus) based

  on the nature of the free group on each extremity.

  

  For example, the hormone insulin (Figure 9.3) has two polypeptide chains, A and B,

  shown in diagram below. Each chain has its own set of amino acids, assembled in a

  particular order. For instance, the sequence of the A chain starts with glycine at the

  N-terminus and ends with asparagine at the C-terminus, and is different from the

  sequence of the B chain. You may notice that the insulin chains are linked together

  by sulfur-containing bonds between cysteines.

  
  

  
  

  
  

  
  

  
  

  
  

  
  

• Unit 10 : PHASE DIAGRAMS

  UNIT 10. PHASE DIAGRAMS

  Key unit competency:

  To be able to interpret the phase diagrams for different compounds.

  • Define a phase

  • Explain the term phase equilibrium

  • Explain the effect of change of state on changing pressure and temperature

  • Define heterogeneous and homogeneous equilibria

  • Define triple point, critical point, normal boiling and melting points of

  substances

  • Relate the physical properties of compounds to their phase diagrams.

  • Locate triple point, critical point, normal boiling and melting points on the

  phase

  • diagrams

  • Compare the phase diagrams for water with that carbon dioxide

  • Develop analysis skills, team work, and attentiveness in interpreting the

  phase diagrams and in practical activities

  

  10.1. Phase equilibrium

  

  10.1.1. Definition of key terms

  A phase is a homogeneous portion of a system which has uniform physical

  characteristics. It can be separated from other parts of the system by a clear boundary

  (limit). A phase can be a solid, liquid, vapor (gas) or aqueous solution which is uniform

  in both chemical composition and physical state.

  

  A component: it is a chemical species which may be used to specify the

  composition of a system. For example;

  • A three-phase system of water (i.e. water, ice, and vapor) is a one component

  system. The constituent substance of the three phases is water only.

  • A mixture of water and ethanol is a one phase, two components system

  because there are two different chemical compositions.

  • An equilibrium: it is the state of a reaction or physical change in which the

  rates of the forward and reverse processes are the same and there is no net

  change on the amount of the equilibrium components

  • A phase equilibrium: it is a balance between phases, that is, the coexistence

  of two or more phases in a state of dynamic equilibrium. The forward process

  is taking place at the same rate as the backward process and therefore the

  relative quantity of each phase remains unchanged unless the external

  condition is altered.

  

  In homogeneous equilibrium, all substances are in the same phase while in

  heterogeneous equilibrium, substances are in distinct phases.

  

  10.2. Homogeneous and heterogeneous equilibria

  
  

  1. Homogeneous equilibrium

  A system with one phase only is described as a homogeneous system and when

  this system is at equilibrium, it is said to be a homogenous equilibrium.

  In general, a homogeneous equilibrium is one in which all components are present

  in a single phase. In a case of a chemical reaction, both reactants and products exist

  in one phase (gaseous phase, liquid phase or aqueous solution and solid phase).

  For example, in the esterification of acetic acid and ethanol the equilibrium is

  homogeneous because all involved substances are in the same liquid phase.

  All the reactants and products are liquids

  2. Heterogeneous equilibrium

  A system consisting of more than one distinct phases is described as heterogeneous

  system. A heterogeneous equilibrium is a system in which the constituents are

  found in two or more phases. The phases may be any combination of solid, liquid,

  gas, or solutions.

  For example, in the manufacture of quick lime from lime stone the following equilibrium is involved:

  

  It is a heterogeneous equilibrium because some of the components are solids (lime

  stone and quick lime) and another is a gas (carbon dioxide).

   

  10.3. Phase diagrams

  

  Aphase diagram is a graph illustrating the conditions of temperature and pressure

  under which equilibrium exists between the distinct phases (states of matter) of a

  substance. Phase diagrams are divided into three single phase regions that cover

  the pressure-temperature space over which the matter being evaluated exists:

  liquid, gaseous, and solid states. The lines that separate these single-phase regions

  are known as phase boundaries. Along the phase boundaries, the matter being

  evaluated exists simultaneously in equilibrium between the two states that border

  the phase boundary.

  The general form of a phase diagram for a substance that exhibits three phases is

  shown below in the Figure 10.1

  

  Under appropriate conditions of temperature and pressure of a solid can be in

  equilibrium with its liquid state or even with its gaseous state.The phase diagram

  allows to predict the phase of substance that is stable at any given temperature and

  pressure. It contains three important curves, each of which represents the conditions

  of temperature and pressure at which the various phases can coexist at equilibrium.

  i. Boiling point

  The line TC is the vapor pressure curve of the liquid. It represents the equilibrium

  between the liquid and the gas phases. The temperature on this curve where the

  vapor pressure is equal to 1atm and it is the normal boiling point of the substance.

  The vapor pressure curve ends at the critical point (C) which is the critical

  temperature corresponding to the critical pressure of the substance which is the

  pressure required to bring about liquefaction at critical temperature.

  ii. Critical point

  Critical point consists of the temperature and pressure beyond which the liquid and

  gas phases cannot be distinguished. Every substance has a critical temperature

  above which the gas cannot be liquefied, regardless the applied pressure.

  iii.Sublimation point

  The line AT is the sublimation curve which represent the variation in the vapor

  pressure of the solid as it sublimes into gas at different temperatures. The reverse

  process is deposition of the gas as a solid. Sublimation point is the temperature at

  which the solid turns to gas at a constant pressure.

  iv. Melting point

  The line TB is the melting point curve which represent the change in melting point

  of the solid with increasing pressure.The line usually slopes slightly to the right as

  pressure increases. For most substances, the solid is denser than the liquid, therefore,

  an increase in pressure favors the more compact solid. Thus, higher temperatures

  are required to melt the solid at higher pressures. The temperature at which the

  solid melts at a pressure of 1atm is the “normal melting point”.

  v. Triple point

  The triple point T is a point where the three curves intersect. All the three phases

  exist at equilibrium at this temperature and pressure. The triple point is unique for

  each substance.

  vi. Supercritical fluid

  Supercritical fluid of a substance is the temperature and pressure above its own

  thermodynamic critical point that can diffuse through solids like a gas and dissolved

  materials like a liquid.

  Any point on the diagram that does not fall on a line corresponds to conditions

  under which one phase is present. Any other point on the three curves represents

  equilibrium between two phases.

  The gas phase is stable phase at low pressures and elevated temperatures.The

  conditions under which the solid phase is stable extend to low temperatures and

  high pressures.The stability range for liquids lie between the other two regions. That

  is between solid and liquid regions.

  10.3. 1. Phase diagram of water

  Water is a unique substance in many ways due to its properties. One of these special

  properties is the fact that solid water (ice) is less dense than liquid water just above

  the freezing point. The phase diagram for water is shown in the Figure 10.2.

  

  Water can turn into vapor at any temperature that falls on the vapor pressure curve

  depending on the conditions of pressure, but the temperature at which water liquid

  turns into vapor at normal pressure (1atm) is called the normal boiling point of

  water, 100 °C (Figure 10.2)

  Point E in the Figure 10.2 is the critical point of water where the pressure is equal

  to 218 atm and the temperature is about 374 °C. At 374°C, particles of water in the

  gas phase are moving rapidly.At any other temperature above the critical point of

  water, the physical nature of water liquid and steam cannot be distinguished; the

  gas phase cannot be made to liquefy, no matter how much pressure is applied to

  the gas.

  The phase diagram of water is not a typical example of a one component system

  because the line AD (melting point curve) slopes upward from right to left. It has a

  negative slope and its melting point decreases as the pressure increases. This occurs

  only for substances that expand on freezing. Therefore, liquid water is denser than

  solid water (ice), the reason why ice floats on water.

  10.3.2. Phase diagram of carbon dioxide

  Compared to the phase diagram of water, in the phase diagram of carbon dioxide

  the solid-liquid curve exhibits a positive slope, indicating that the melting point for

  CO2 increases with pressure as it does for most substances. The increase of pressure

  causes the equilibrium between dry ice and carbon dioxide liquid to shift in the

  direction of formation of dry ice that is freezing. Carbon dioxide contracts on freezing

  and this implies that dry ice has higher density than that of liquid carbon dioxide.

  The Figure 10.3 shows the phase diagram of carbon dioxide.

  

  The triple point is observed at the pressure above 1atm, indicating that carbon

  dioxide cannot exist as a liquid under normal conditions of pressure. Instead,

  cooling gaseous carbon dioxide at 1atm results in its deposition into the solid state.

  Likewise, solid carbon dioxide does not melt at 1atm pressure but instead sublimes

  to yield gaseous CO2.

  

  10.4. Comparison of phase diagrams of substances that expand and those that contract on freezing

  

  For the phase diagrams, some materials contract on freezing while others expand

  on freezing. The main differences between substances that expand and those

  that contract on freezing can be highlighted by comparing the phase diagrams

  of carbon dioxide and that of water. In the phase diagram of carbon dioxide, the

  substance contracts on freezing and that of water expands on freezing.

  Both phase diagrams for water and carbon dioxide have the same general Y-shape,

  just shifted relative to one another. This shift occurs because the liquid phase in

  the dry ice can only occur at higher temperatures and pressures, whereas, in ice

  the liquid phase occurs at lower temperatures and pressures. There are two more

  significant differences between the phase diagram of carbon dioxide and that of

  water:

  10.4.1. Melting point curve

  The melting point curve of carbon dioxide slopes upwards to right (Figure 10.3)

  whereas that of water slopes upward to left (Figure 10.2). This means that for carbon

  dioxide the melting point increases as the pressure increases, a characteristic

  behavior of substances that contract on freezing. Further, water expands on freezing

  (Figure 1.4) and this unusual behavior is caused by the open structure of the regular

  packing of water molecules in ice due to the network of hydrogen bonding in ice

  which is more extensive than in liquid.

  

  Ice floats on liquid water (Figure 10.5), this unusual behavior is caused by the open

  structure of the regular packing of water molecules in ice due to the network of

  hydrogen bonding in ice which is more extensive than in liquid.The ice is less dense

  than water reason why it floats in water.

  
  

  
  

  10.4.2. Triple point

  The triple point of carbon dioxide is above atmospheric pressure. This means that

  the state of liquid carbon dioxide does not exist at ordinary atmospheric pressure.

  Dry ice remains as a solid below -78ºC and changes to fog (gas) above -78ºC. It

  sublimes without forming liquid at normal atmospheric pressure (Figure 10.6). The

  sublimation of carbon dioxide results in a low temperature which causes water

  vapors in the air to form moist. 

  
  

  Ice is stable below 0 ºC and water is stable between 0ºC and 100 ºC while water

  vapor is stable above 100 ºC. At normal atmospheric pressure, ice can first melts and

  ultimately boils as the temperature increases.

  

  10.5. Applied aspect of phase diagrams

  
  

  
  

  The applications of phase diagrams are useful for engineer’s materials and material

  applications. The scientists and engineers understand the behavior of a system

  which may contain more than one component. Multicomponent phase’s diagrams

  show the conditions for the formation of solutions and new compounds. The

  phase diagrams are applied in solidification and casting problems. Many materials

  and alloy system exist in more than one phase depending on the conditions of

  temperature, pressure and compositions. In the area of alloy development, phase

  diagrams have proved invaluable for tailoring existing alloys to avoid over design

  in current applications, each phase has different microstructure which is related

  to mechanical properties. The development of microstructure is related to the

  characteristics of phase diagrams. Proper knowledge and understanding of phase

  diagrams lead to the design and control of heating procedures for developing the

  required microstructure and properties.

  
  

  Phase diagrams are consulted when materials are attacked by corrosion. They

  predict the temperature at which freezing or melting begins or ends. Phase

  diagrams differentiate the critical point, triple point, normal boiling point, etc of

  some substances.

  

  In general the industrial applications of phase diagrams include alloy design,

  processing, and performance.

  

  

  

  

• UNIT 11: SOLUTIONS AND TITRATION

  UNIT 11: SOLUTIONS AND TITRATION

  Key unit competency:

   Be able to prepare standard solutions and use them to determine concentration of

  other solutions by titration.

  Learning objectives:

  •  Define the terms standard solution and primary standard solution.

  •  Explain the properties of a standard primary solution.

  •  Explain the titration process, emphasising the need for precise measurements.

  •  Prepare solutions with different concentrations.

  •  Properly use the burettes, pipettes during titration.

  •  Interpret the experimental data obtained by titration and report.

  •  Carry out acid-base, redox titrations and do calculations involved.

  •  Develop a team approach and a sense of responsibility in performing the

  experiments of titration.

  •  Respect of procedure in practical experiment.

  •  Develop a culture of orderliness in performing practical experiments.

  •  Appreciate the use of appropriate measurements in daily life.

  

  11.1. Definition of standard solution and primary standard solution.

  
  

  In analytical chemistry, a standard solution is a solution containing a precisely known

  concentration of an element or a substance and used to determine the unknown

  concentration of other solutions. A known weight of solute is dissolved to make a

  specific volume. It is prepared using a standard substance, such as a primary standard.

  A primary standard is defined as a substance or compound used to prepare

  standard solutions by actually weighing a known mass, dissolving it, and diluting to a

  definite volume.

  

  

  A good primary standard meets the following criteria:

  •  High level of purity

  •  High stability

  •  Be readily soluble in water

  •  High equivalent weight (to reduce error from mass measurements)

  •  Not hygroscopic (to reduce changes in mass in humid versus dry environments)

  •  Non-toxic

  •  Inexpensive and readily available

  •  React instantaneously, stoichiometrically and irreversibly with other substances

  i.e. should not have interfering products during titration.

  •  It should not get affected by carbon dioxide in air

  Molar concentrations are the most useful in chemical reaction calculations because

  they directly relate the moles of solute to the volume of solution.

  The formula for molarity is:

  

  The preparation of the solution requires a reagent that is so to say the quantity to be

  weighed for mass if the reagent is in solid state ; The volume to be pipetted using

  pipette if the solute is a liquid and then after dissolve it in water, So the solution can be

  prepared by two methods such as dissolution method and dilution method.

  In the preparation of solution, glasses, volumetric flask, pipette, glass rod, measuring

  cylinder, analytical balance, spatula, beakers, magnetic stirrer and other laboratory

  devices are used.

  11.3.1. Preparation of standard solution by dissolution of solids

  

  Scope: This method is applied for solute in solid state and you should be able to

  determine the mass required from calculation to be weighed and provide distilled

  water to dissolve the solute.

  Examples:

  1. Describe in details how you can prepare the following solution: 50 mL of

  NaOH, 10%.

  

  

  

  11.3.2. Preparation of standard solution by dilution
  

  

  

  

  

  

  

  

  
  

  11.4. Simple acid-base titrations
  

  

  

  Titration is the controlled addition and measurement of the amount of a solution

  of known concentration required to react completely with a measured amount of a

  solution of unknown concentration.

  Acid-base titration

  It is the determination of the concentration of an acid or base by exactly neutralizing

  the acid or base of known concentration

  Alkalimetry and acidimetry

  •  Alkalimetry is the specialized analytic use of acid-base titration to determine

  the concentration of a basic substance.

  •  Acidimetry is the same concept of specialized analytic use of acid-base titration

  to determine the concentration of an acidic substance.

  Equivalence point

  The point at which the two solutions used in a titration are present in chemically

  equivalent amount is the equivalence point. At this point the moles of two solutions

  will be equal.

  

  Indicators and pH-meters can be used to determine the equivalence point. The

  point in a titration at which an indicator changes color is called the end-point of the

  titration.

  Equipments and set up (Figure 11.1) of materials for Titration

  The common equipment used in a titration are:

  •  Burette

  •  Pipette

  •  pH-indicator/acid-base indicator

  •  White tile: used to see a color change in the solution(a white paper can also be

  used)

  •  Conical flask (Erlenmeyer flask)

  •  Titrant: a standard solution of known concentration

  •  Analyte: a solution of unknown concentration

  

  How to perform titrations

  Knowing the use of pipette and burettes and how to handle them, the following

  points are useful in order for a correct titration to be done:

  1. The apparatus should be arranged as shown in the above Figure.

  2. The burette tap is opened with the left hand and the right hand is used to

  shake the conical flask.

  3. The equivalence-point is reached when the indicator just changes

  permanently the colour.

  4. At the end-point, the level of the titrant is read on the burette

  5. The titration is now repeated, three more times are recommended. Towards

  the end-point, the titrant is added dropwise to avoid overshooting.

  Notice: Before titration, check if the tip of the burette is filled with the titrant, and doesn’t

  contain bulb of air. If there is a bulb of air, a quick opening and closing of the tap will expel

  the air out of the burette.

  Choice of indicators in acid-base titrations

  When the technique of acid-base titration is extended to a wide variety of acidic and

  alkaline solution, care needs to be taken about the choice of indicator for any given

  reaction.

  The choice of an inappropriate indicator would lead to incorrect results, and it is

  therefore extremely important that the indicator is chosen carefully.

  The principle on which a choice of indicator is made concerns the strength of the acid

  or base involved in the reaction. Note that the strength of an acid or base is not to be

  confused with the concentration of its solution. Example of strong and weak acids and

  bases and choice of indicator are given in the Table below.

  

  
  

  
  

  11.5. Titration involving redox reactions

  

  11.5.1. Titrations with potassium manganate (VII); KMnO₄

  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  

  76-104
  

  
  

  
  
• UNIT 12: CONDUCTIVITY OF SOLUTIONS

  Key unity competence:

  To be able to: Explain the effect of different factors on the molar conductivity of different electrolytes and the applications of conductivity measurements.

  Learning objectives:

  • Explain the conductivity of solutions.

  • State and explain the factors that affect molar conductivity of solutions.

  • State Kohlrausch’s law of individual molar conductivity.

  •Use Kohlrausch’s law to calculate the molar conductivity of an electrolyte.

  • Interpret a graph of molar conductivity against concentration for both weak and strong electrolytes.

  • Compare and contrast metallic conductivity and electrolytic conductivity.

  • Develop a team approach and responsibility in performing experiments.

  • Appreciate the contributions of other scientists like Kohlrausch’s law in calculation of molar conductivity of solutions.

  • Respect the procedure in performing experiment.

  2.1. Conductance of electrolytic solutions
  

  

  

  
  

  

  

  

  

  

  Conductivity: Definition and description

  
  

  Conductivity of a substance is defined as the ability or power to conduct or transmit heat, electricity, or sound›. Its units are Siemens per meter [S/m] in SI and milliohms per centimeter [m mho/cm] in U.S. customary units.

  12.2. Measurement of conductivity of solutions

  

  The conductivity is the reciprocal of the resistance (1/R) and is measured in Siemens or mhos.Conductivity measurements are used routinely in many industrial and environmental applications as a fast, inexpensive and reliable way of measuring the ionic content in a solution. For example, the measurement of conductivity is a typical way to monitor and continuously trend the performance of water purification systems.

  Electrical conductivity meter

  

  Principle of the measurement

  The electrical conductivity of a solution of an electrolyte is measured by determining the resistance of the solution between two flat or cylindrical electrodes separated by a fixed distance. An alternating voltage is used in order to avoid electrolysis. The resistance is measured by a conductivity meter. Typical frequencies used are in the range 1–3kHz. The dependence on the frequency is usually small, but may become appreciable at very high frequencies.

  A wide variety of instrumentation is commercially available. There are two types of cell, the classical type with flat or cylindrical electrodes and a second type based on induction. Many commercial systems offer automatic temperature correction. Tables of reference conductivities are available for many common solutions.

  The conductivity of an electrolyte is the conductance of a volume of solution containing one mole of dissolved electrolyte placed between two parallel electrodes 1dm apart and large enough to contain between them all the solution; the conductivity is affected by temperature.

  Checking up 12.2

  Describe the functioning of conductivity meter and derive the formula of calculation of conductivity

  12.3. Specific conductivity of solutions

  Activity 12.3:

  1. Define resistivity

  2. Establish a relation between conductivity and resistivity and among the following substances, which ones are conductors and non-conductors, for each you have to explain why they are or not conductors: pure water, sugar, iron plate, clothes, plastic bags, ammonia solution, salt solution, etc...

  Specific Conductivity (better known as specific conductance) is the measure of the ability of that material to conduct electricity. It is represented by the symbol “К”. Hence, by definition, the specific conductance (specific conductivity), κ (kappa) is the reciprocal of the specific resistance. The SI unit of conductivity is Siemens per meter (S/m).

  

  

  

  12.4. Molar conductivity of solutions

  

  The molar conductivity of a solution at any given concentration is the conductance of the volume of solution containing one mole of electrolyte kept between two electrodes with the unit area of cross section and distance of unit length. In general terms, it is defined as the ratio of specific conductivity and the concentration of the electrolyte.

  

  

  12.4.1. Strong electrolytes

  For strong electrolyte, molar conductivity increases steadily with dilution until it reaches the maximum value at infinite dilution (at high concentration, the lower conductivity values are due to ionic interference. The formation of ionic pairs or triplet and symmetrical spheres greatly reduces the mobility of ions however as the dilution increases, there is reduced ionic interference as result of many solvent molecules surrounding the oppositely charged ions thus an increase in molar conductivity.

  At infinite, there is independent migration of ions that is ions experience negligible ionic interference and move independent of each other.

  The molar conductivities of strong electrolytes are high. This is because, by nature, strong electrolytes are highly dissociated when molten or when in solution into large number of ions. These ions are mobile, hence they migrate to the electrodes, resulting in the high conduction of electricity: the higher the number of ions are free in solution, the higher the conductivity.

  

  This graph can be obtained by extrapolation of the graph to zero concentration.

  12.4.2. Weak electrolytes

  Weak electrolytes show partial dissociation in solution, producing few ions, which results in low conduction of electricity.

  A weak electrolyte dissociates to a much lesser extent so its conductance is lower than that of a strong electrolyte at the same concentration.

  

  Summary:

  • The higher the number of ions per unit volume in solution, the greater the conductivity of the electrolytic solution. This means that the conductivity increases with concentration of ions in solution up to an optimum level over which it starts decreasing.

  • On the other hand when the conductivity has decreased due to very high concentration of ions, it can be increased with dilution (i.e. lower concentrations) up to its optimum, beyond which further dilution will decrease conductivity.

  • The decrease or increase of conductivity by concentrating or diluting the solution is sharp in strong electrolytes while it is gradual in weak electrolytes. The following graph shows

  

  The table below shows the trend in conductivity with dilution for a strong and a weak acid.

  

  Explanation of Increase in Conductivity with Dilution:

  With increase in dilution (decrease in concentration), ions become farther apart, and inter-ionic forces (i.e. forces of attraction between unlike ions and forces of repulsion between like ions) decrease considerably, so that greater number of ions are able to migrate to the electrodes. In addition, due to change in equilibrium, the electrolyte undergoes further ionization from the same mass in solution (in order to balance the effect). Hence, more ions (conducting species) are introduced into the solution

  12.5. Molar conductivity at infinite dilution

  Kohlrausch’s law of independent migration of ions states that “at infinite dilution, where ionization of all electrolytes is complete and where all interionic effects are absent, the molar conductivity of an electrolyte is the sum of the molar conductivities of its constituent ions at constant temperature”

  

  

  

  

  12.5. Factors that affect molar conductivity of solutions

  

  

  

  

  

  12.6. Kohlrausch’s law of individual molar conductivity

  

  

  

  

  12.7. Use of conductivity measurement in titration and solu-bility product

  

  12.7.1. Using conductivity to find the end point of a titration

  The end-point in titration experiment can be determined using conductivity. The procedure of the technique is:At the start of this titration the conical flask contains a strong alkali that is fully ionized in water. If electrodes are placed inside the conical flask the ions in the water will conduct electricity and a current will flow.The more ions there are the better the conductivity and the higher the current will be. The current can be measured using an ammeter. As acid is added to the alkali hydrogen ions and hydroxide ions react together to form water molecules.The number of ions in the conical flask starts to decrease and the current flowing through the solution will decrease. At neutralization all of the hydrogen ions and hydroxide ions have reacted together to form water molecules.The neutral solution contains only salt ions dissolved in water molecules. The solution will still conduct electricity because of the salt ions but the current will be at a minimum.As more acid is added the current will start to increase because there will now be unreactedhydrogen ions in the solution as well as the salt ions. The solution is now no longer neutral but has become acidic.

  If you draw a graph of current against the amount of acid added you can see where the minimum is. This is the end point of the titration at neutralization.

  

  12.7.2. Determination of solubility product by conductivity measure-ment.

  Solubility product, Ksp, is the mathematical product of its dissolved ion concentrations raised to the power of their stoichiometric coefficients. Solubility products are relevant when a sparingly soluble ionic compound releases ions into solution.That is the product of the concentration of ions in the solution which are in equilibrium with the solid ion. These concentrations can be determined via conductivity measurements, consider the following examples :

  

  The measurement of conductivity will depend on the value of Ksp for the sparingly soluble substances. The measurement of the specific conductivity, K of the saturated solution leads to a value of the concentration.

  

  

  12.8. Difference between metallic conductivity and electrolyt-ic conductivity

  

  The substances, which allow the passage of electric current, are called conductors. The best metal conductors are such as copper, silver, tin, etc. On the other hand, the substances, which do not allow the passage of electric current through them, are called non-conductors or insulators. Some common examples of insulators are rubber, wood, wax, etc.

  The conductors are broadly classified into two types, Metallic and electrolytic conductors.

  

  The electrolyte may, therefore, be defined as the substance whose aqueous solution or fused state conducts electricity accompanied by chemical decomposition. The conduction of current through electrolyte is due to the movement of ions.On the contrary, substances, which in the form of their solutions or in their molten state do not conduct electricity, are called non-electrolytes.

  

  

  

  

  

  
  

  
  

  
  

  
  

  
  

• UNIT 13 : ELECTROLYSIS

  Key unit competence:

  Predict the products of given electrolytes during electrolysis and work out quantitatively to determine how much is liberated at a given electrode using Faraday’s law.

  Learning Objectives:

  • Define electrolysis, cathode and anode.

  • Explain the electrolysis of different substances.

  • State Faraday’s laws and define the Faraday’s constant.

  • Develop practical experimental skills related to electrolysis, interpret results, and draw valid conclusions.

  • Carry out a practical activity to explain the phenomenon of electrolysis.

  • Compare the electrolysis of dilute solutions and concentrated solutions.

  • Calculate the masses and volumes of substances liberated during electrolysis.

  • Relate the nature of electrode, reactivity of metal ion in solution to the products of electrolysis.

  • Perform electroplating of graphite by copper metal

  Introductory activity

  Observe carefully the figure below and answer the following question. Record your answer and discuss them.

  

  1. Label the set up and give the name of this Experiment.

  2. Suggest how water can be decomposed into hydrogen and oxygen.

  13.1. Definition of electrolysis and Description of electrolytic cells.

  Activity 13.1:
  

  A.1. In one case, you have a source of water at the top of a hill and you want to supply water to a community in the valley down the hill

  2. In another case, you have a community at the top of a hill and you want to supply water to the community from a source located in the valley down. Students in groups discuss how they would proceed to supply water to the communities in the above two cases.

  B. Why do we cook food by heating?

  C. What is the difference between a spontaneous reaction and a non-spontaneous reaction?

  D. Have you heard about electrolysis? If yes, can you say what it is about?

  1.Definition of electrolysis.

  A spontaneous reaction is a reaction that favors the formation of products without external energy. It is a process that will occur on its own. For example, a ball will roll down an incline, water will flow downhill, radioisotopes will decay, and iron will rust. No intervention is required because these processes are thermodynamically favorable.

  A nonspontaneous reaction (also called an unfavorable reaction) is a chemical reaction that necessitates external energy to occur. For example, without an external energy source, water will remain water forever. Under the right conditions, using electricity (direct current) will help to produce hydrogen gas and oxygen gas from water. Cooking foods is not spontaneous reaction that is why heat is used.

  Electrolyte: Sodium chloride is an ionic compound in which ions arrange themselves in a rigid cubic lattice when in solid state. In this state, it cannot allow electric current to pass through it. However, when it is melted, or dissolved in water, the rigid lattice is broken, ions are free to move and electric current can pass. Therefore, it is classified as an electrolyte.

  Substances which cannot allow the flow of electriccurrent when in molten or in solution are referred to as non-electrolytes. When electric current (direct current) flows through an electrolyte, it decomposes it. This phenomenon is called electrolysis.

  Thus electrolysis is the decomposition of an electrolyte by passage of an electric current through it. Therefore for electrolysis to take place, there must be a source of direct current. The direct current is conveyed from its source to the electrolyte by means of a metallic conductor and electrodes. The electrode connected to the positive terminal of the direct current is called the anode and the one connected to the negative terminal is the cathode. By convention, the electric current enters the electrolyte by the anode and leaves by the cathode.

  When the current passes through an electrolytic solution, ions migrate and electrons are gained or lost by ions on the electrodes surface. Electrode that is positively charged has deficit of electrons is called anode and the other electrode negatively charged has excess of electrons and is called cathode. Chemical changes at the electrodes due to the passage of electric current are called electrolysis.

  2. Description of electrolytic cells.

  An electrolytic cell is an electrochemical cell that drives a non-spontaneous redox reaction through the application of externalelectrical energy. They are often used to decompose chemical compounds, in a process called electrolysis . The Greek word lysis means to break up.

  Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, production of sodium metal, Na, from molten NaCl, production of aluminium and other chemicals. Electroplating (e.g. of copper, silver, nickel or chromium) is done using an electrolytic cell.

  An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes. When driven by an external voltage applied to the electrodes, the ions in the electrolyte are attracted to an electrode with the opposite charge, where charge-transferring (also called faradaic or redox) reactions can take place. Only with an external electrical potential (i.e. voltage) of correct polarity and sufficient magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. The electrical energy provided can produce a chemical reaction which would not occur spontaneously otherwise.

  The main components required to achieve electrolysis are:

  • An electrolyte is substance containing free ions which are the carriers of electric current in the electrolyte. If the ions are not mobile, as in a solid salt then electrolysis cannot occur.

  • A direct current (DC) supply: provides the energy necessary to create or discharge the ions in the electrolyte. Electric current is carried by electronds in the external circuit.

  • Electrolysis depends on controlling the voltage and current.

  • Alternating current (AC) would not be appropriate for electrolysis. Because the “cathode” and “anode” are constantly switching places, AC produces explosive mixtures of hydrogen and oxygen.

  • Two electrodes: an electrical conductor which provides the physical interface between the electrical circuit providing the energy and the electrolyte.

  

  The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons from the external circuit. The products of electrolysis are in some different physical states from the electrolyte and can be removed by some physical processes.

  Electrodes of metal, graphite and semi-conductor material are widely used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and the cost of manufacture.

  Note:

  The suitable electrode in electrolysis should be inert (Cu, Pt, etc.) therefore it will not participate in the chemical reaction.

  It is very easy to be confused about the names CATHODE and ANODE and what their properties are, both with electrochemical cells and electrolytic cells.

  (To help you to remember, Cathode is the site of reduction, or, if you prefer, CCC = Cathode Collects Cations. Anode is the site of oxidation, or, AAA = Anode Attracts Anions.)

  Checking up 13.1:

  Choose the correct answer from the options given below each of the following questions:

  1) Which of the following substances is an electrolyte?

  a) Mercury

  b) Copper

  c) Sodium sulphate

  d) Aluminium

  2) Which of the following substances is a weak electrolyte?

  a) Dilute hydrochloric acid

  b) Dilute sulphuric acid

  c) A solution of potassium bromide.

  d) Carbonic acid

  3) Which of the following statements is true for the formation of sodiumchloride by the direct combination of sodium with chlorine?

  a) Sodium is reduced

  (b) Chlorine is oxidized.

  c) Chlorine is the oxidising agent

  (d) Sodium is the oxidizing agent.

  4) Which of the following species will be deposited at the cathode on theelectrolysis of an aqueous solution of potassium bromide?

  a) K

  b) H₂

  c) Br₂

  d) O₂

  5) If you want to electrolyse concentrated HCI, which of the following will youchoose for making the anode?

  a) Graphite

  b) Aluminium

  c) Iron

  d) Copper

  13.2. Electrolysis of sodium chloride

  Activity 13.2

  Investigating the effect of concentration on the products formed during electrolysis of concentrated sodium chloride solution.

  Materials: Carbon or graphite rods, connecting wire, U-tube, dry cell, glass syringes, concentrated sodium chloride, cork and switch.

  

  Procedure:

  2. Add 10g of sodium chloride to 100cm3 of distilled water.

  3. Warm the mixture and continue adding sodium chloride until a saturated solution is formed.

  4. Put the saturated solution in U-tube and fit it with carbon rods and glass syringes.

  5. Level the brine solution in the two arms and switch on the circuit. Record any observations made after some time. Identify any gases collected in the syringe.

  Questions:

  a. Identify the gases formed by testing them using litmus papers.

  b. Using ionic equations, explain how the products are formed.

  Sodium chloride may be in different forms that can be electrolytes. It may be in its molten state, dilute solution or concentrated solution. In each case, the products of electrolysis differ because of different factors.

  13.2.1. Electrolysis of molten sodium chloride

  The molten salt is introduced in a container called electrolytic cell (or electrolysis cell) in which there are two inert electrodes (platinum or graphite). Electrodes are connected to a DC generator.

  • Cations (Na⁺) move toward the cathode (negative electrode), where they take electrons and are reduced. On cathode metallic sodium is deposited:

  

  • Anions (Cl^-) move towards the anode (positive electrode), where they give up electrons and are oxidized. On the anode Cl₂ is released:

  The overall reaction is the addition of half reactions at electrodes:

  

  The cathode provides electrons so it is a reducing site.

  The anode takes electron, so it is an oxidizing site.

  

  • Another important thing to note is that twice as much hydrogen is produced as oxygen. Thus the volume of hydrogen produced is twice that of oxygen. Refer to the equations above and note the number of electrons involved to help you

  13.2.2. Electrolysis of Dilute Sodium Chloride Solution

  An aqueous solution of sodium chloride contains four different types of ions. They are ions from sodium chloride: Na⁺ (aq) and Cl^- (aq) Ions from water: H⁺ (aq) and OH^- (aq)When dilute sodium chloride solution is electrolysed using inert electrodes, the Na⁺ and H⁺ ions are attracted to the cathode. The Cl^- and OH^- ions are attracted to the anode.

  

  The table shown below is simply a table of standard reduction potentials in decreasing order. The species at the top have a greater likelihood of being reduced while the ones at the bottom have a greater likelihood of being oxidized. Therefore, when a species at the top is coupled with a species at the bottom, the one at the top will be easily reduced while the one at the bottom will be oxidized.

  • At the cathode:

  The H⁺ and Na+ ions are attracted to the platinum cathode. H⁺ ions gains electrons from the cathode to form hydrogen gas. (The hydrogen ions accept electrons more readily than the sodium ions. As a result, H⁺ ions are discharged as hydrogen gas, which bubbles off. Explanation why H+ ions are preferentially discharged will be given later.)

  2H⁺(aq) + 2e^-→ H₂(g) ,Na⁺ ions remain in solution.

  • At the anode:

  OH^- and Cl^- are attracted to the platinum anode. OH^- ions give up electrons to the anode to form water and oxygen gas.

  

  Note:

  • Since water is being removed (by decomposition into hydrogen and oxygen), the concentration of sodium chloride solution increases gradually. The overall reaction shows that the electrolysis of dilute sodium chloride solution is equivalent to the electrolysis of water.

  understand.

  

  13.2.3. Electrolysis of Concentrated Sodium Chloride Solution

  The only difference with dilute NaCl solution is that at the anode, Cl^- ions are more numerous than OH^- ions. Consequently, Cl^- ions are discharged as chlorine gas, which bubbles off. A half-equation shows you what happens at one of the electrodes during electrolysis.

  

  Sodium ions Na⁺ and hydroxide OH^– are also present in the sodium chloride solution. They are not discharged at the electrodes. Instead, they make sodium hydroxide solution.

  These products are reactive, so it is important to use inert (unreactive) materials for the electrodes.

  One volume of hydrogen gas is given off at the cathode and one volume of chlorine gas is produced at the anode. The resulting solution becomes alkaline because there are more OH^- than H⁺ ions left in the solution.

  

  Checking up 13.2:

  With the help of equations of reactions which occur at each electrode, outline what happens during electrolysis of dilute aqueous sodium chloride. What happens to the pH of the solution as electrolysis continues?

  13.3. Electrolysis of water

  Activity 13.3: Investigate the products formed during the electrolysis of water

  Materials:

  • Distilled water

  • Tap water

  • 2 silver-colored thumb tacks

  • 9V battery

  • Small, clear plastic container

  • 2 test tubes

  • Stopwatch

  • Baking soda

  • Table salt

  • Lemon

  • Dish washing detergent

  

  Procedure:

  3. Insert the thumb tacks into the bottom of the plastic container so that the points push up into the container. Space them so that they’re the same distance apart as the two terminals of the 9V battery. Be careful not to harm yourself!

  4. Place the plastic container with the terminals of the battery. If the cup is too large to balance on the battery, be sure thumb tacks are connect to positive and negative pushpins and do no touch each other.

  5. Slowly fill the container with distilled water. If the tacks move, go ahead and use this opportunity to fix them before you proceed. Will distilled water conduct electricity on its own? Try it!

  6. Add a pinch of baking soda.

  7. Hold two test tubes above each push pin to collect the gas being formed. Record your observations. What happens? Does one tube have more gas than the other? What gases do you think are forming?

  8. Discard the solution, and repeat the procedure with a different combination:

  • Distilled water and lemon juice

  • Distilled water and table salt

  • Distilled water and dish detergent

  • Distilled water (no additive)

  • Tap water (Does tap water works? If so, why?)

  Question: During the electrolysis of water, which electrolyte conducts electricity the best?

  Water can be decomposed by passing an electric current through it. When this happens, the electrons from the electric current cause an oxidation-reduction reaction. At one electrode, called the cathode, electrons cause a reduction. At the other electrode, called the anode, electrons leave their ions completing the circuit, and cause an oxidation.In order to carry out electrolysis the solution must conduct electric current. Pure water is a very poor conductor.

  To make the water conduct better we can add an electrolyte (NaCl) to the water. The electrolyte added must not be more electrolyzable than water. Many electrolytes that we add electrolyze more easily than water. Sulfate ions do not electrolyse as easily as water, so sulfates are often used to enhance the conductivity of the water

  Water may be electrolyzed in the apparatus shown below. Pure water is however a very poor conductor of electricity and one has to add dilute sulphuric acid in order to have a significant current flow.

  

  The electrodes consist of platinum foil. The electrolyte is dilute sulphuric acid.

  Hydrogen gas is evolved at the cathode, and oxygen at the anode.

  The ratio, by volume, of hydrogen to oxygen, is exactly 2:1.

  Remember that electron flow in the circuit is opposite to the conventional current flow.

  The reaction at the cathode (tube A) is the reduction of protons:

  

  Oxidation takes place at the anode (tube B). There are two anions competing to give up their electrons, namely sulphate (SO4₂^-) from sulphuric acid, and hydroxide (OH^-) from the ionization of water. Here, the activity series is used to know the ion to be discharged.

  The oxidation of OH^- according to the reaction:

  In pure water at the negatively charged cathode, a reduction reaction takes place, with electrons (e⁻) from the cathode being given to hydrogen cations to form hydrogen gas. The half reaction, balanced with acid, is:

  

  Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen:

  

  The number of hydrogen molecules produced is thus twice the number of oxygen molecules. Assuming equal temperature and pressure for both gases, the produced hydrogen gas has therefore twice the volume of the produced oxygen gas.

  Checking up 13.3

  I understand the process of water electrolysis, that water as an electrolyte can be decomposed into hydrogen and oxygen via an external energy source (an electrical current). I know that the reduction of hydrogen takes place on the cathode and the oxidation reaction takes place on the anode. I also know that water, already partially split into H⁺ and OH^- (though there are very few of these ions in pure water).

  1. Electric current (direct current) electrolyzes water. Discuss this statement.

  2. Why alternative current are not used for the same process?

  13.4. Electrolysis of concentrated copper (II) sulphate solution using inert electrode

  Activity 13.4:

  Investigating what happens when a solution of copper (II) sulphate is electrolysed using carbon and copper electrodes.

  Apparatus and chemicals: Glass cell, Carbon rod, 2M copper (II) sulphate solution, connecting wires, dry cells, copper plates, propanone and litmus paper.

  

  Procedure:

  3. Determine the mass of the graphite rods and record it.

  4. Put 0.5M of copper (II) sulphate solution in a glass cell with the carbon (graphite) rods and set up the apparatus. Carefully observe all the changes taking place at the electrodes and the solution. Test the resulting solution with blue litmus paper.

  5. After some time, switch off the current, remove the electrodes, wash them in propanone, dry then and then weigh them.

  6. Repeat the experiment using clean strips of copper metal as electrodes. Weigh them and then complete the circuit using freshly prepared copper (II) sulphate solution. Record your observations.

  Questions:

  1. Explain the changes observed during the electrolysis of copper (II) sulphate using :

  a.Carbon electrodes

  b.Copper electrodes

  2. Outline the changes that occur in the solution from the beginning to the end of the experiment.

  The products of electrolysing of copper sulphate solution with inert electrodes (carbon/graphite or platinum) are copper metal and oxygen gas.

  

  Using the simple apparatus (diagram above) and inert carbon (graphite) electrodes, you can observe the products of the electrolysis of copper sulfate solution are a copper deposit on the negative cathode electrode and oxygen gas at the positive anode electrode. This anode reaction differs if you use copper electrodes. You have to fill the little test tubes with the electrolyte (dilute copper sulphate solution), hold the liquid in with your finger and carefully invert them over the nearly full electrolysis cell. The simple apparatus (above) can be used with two inert wire electrodes.

  The blue colour fades as more and more copper is deposited, depleting the concentration of blue copper ion Cu²⁺ in solution.

  

  The electrode reactions and products of the electrolysis of the electrolyte copper sulphate solution (with inert carbon-graphite electrodes) are illustrated by the diagram above

  (a) The electrode products from the electrolysis of copper sulphate with inert graphite (carbon) electrodes

  The negative cathode electrode attracts Cu²⁺ ions (from copper sulphate) and H+ions (from water). Only the copper ion is discharged, being reduced into copper metal. The less reactive a metal, the more readily its ion is reduced on the electrode surface.

  A copper deposit forms as the positive copper ions are attracted to the negative electrode (cathode)

  

  The traces of hydrogen ions are not discharged, so you do not see any gas collected above the negative electrode. The blue colour of the copper ion will fade as the copper ions are converted into the copper deposit on the cathode

  At the positive anode reaction with graphite electrodes

  Oxygen gas is formed at the positive electrode, an oxidation reaction (electron loss).

  The negative sulphate ions (SO₄^2-) or the traces of hydroxide ions (OH^–) are attracted to the positive electrode. But the sulphate ion is too stable and nothing happens. Instead hydroxide ions are discharged and oxidised to form oxygen.

  

  Checking up 13.4

  1. Name the product at the cathode and anode during electrolysis of:

  a. Molten lead bromide with inert electrode.

  b. Acidified copper sulphate solution with inert electrodes.

  c. Acidified water with inert electrode.

  d. Dilute hydrochloric acid with inert electrode.

  e. Concentrated hydrochloric acid with inert electrode.

  2.Predict the products formed when the following molten compounds are electrolysed using carbon electrodes;

  a. Lead(II) bromide

  b. Magnesium oxide

  13.5. Electrolysis of concentrated copper (II) sulphate solution using copper electrodes

  The products of electrolysis of copper sulfate solution with copper electrodes are copper metal and copper ions (the copper anode dissolves).

  The electrolysis of copper (II) sulphate solution using copper electrodes is shown below.

  

  Using the simple apparatus and two copper electrodes the products of the electrolysis of copper sulphate solution are a copper deposit on the negative cathode electrode and copper dissolves, Cu⁺², at the positive anode electrode. This copper anode reaction differs from the one when you use an inert graphite electrode for the anode.

  When Copper (II) sulphate is electrolysed with a copper anode electrode (the cathode can be carbon or copper), the copper deposit on the cathode (–) equals the copper dissolves at the anode (+). Therefore the blue colour of the Cu²⁺ ions stays constant because Cu deposited = Cu dissolved. Both involve a two electron transfer so it means mass of Cu deposited = mass of Cu dissolving for the same quantity of current flowing (flow of electrons). You can check this out by weighing the dryelectrodes before and after the electrolysis has taken place.

  The experiment works with a carbon anode and you see the blackness of the graphite change to the orange-brown colour of the copper deposit.

  

  The electrode reactions and products of the electrolysis of copper sulphate solution (with a copper anode) are illustrated by thediagram above.

  (a) The electrode products from the electrolysis of copper sulphate with copper electrodes

  The negative cathode electrode attracts Cu²⁺ ions (from copper sulphate) and H⁺ions (from water).

  Only the copper ion is discharged, being reduced to copper metal.A reduction electrode reaction at the negative cathode: (copper deposit, reduction 2 electrons gained) reduction by electron gain

  The positive anode reaction with copper electrodes

  
  

  Copper atoms oxidised to copper (II) ions: dissolving of copper in its electrolytic purification or electroplating (must have positive copper anode).

  Copper (II) ionsreduced to copper atoms: deposition of copper in its electrolytic purification or electroplating using copper (II) sulphate solution, so the electrode can be copper or other metal to be plated or any other conducting material.

  This means for every copper atom that gets oxidised, one copper ion is reduced, therefore
  

  when copper electrodes are used in the electrolysis of copper sulphate solution, the mass loss of copper from the positive anode electrode should equal tothe mass of copper gained and deposited on the negative cathode electrode.

  You can show this by weighing both electrodes at the start of the experiment. After the current has passed for some time, carefully extract the electrodes from the solution, wash them, dry them and reweigh them. The gain in mass of the cathode should be the same as the loss of mass from the anode.
  

  

  13.6. Faraday’s Laws
  

  Activity 13.6
  

  Comparison of the amounts of different substances liberated by the same quantity of electricity.

  

  
  

  Set up the circuit containing a copper voltmeter and a silver voltmeter (a voltmeter is a vessel containing two electrodes immersed in a solution of ions through which a current is to be passed.)

  
  

  Identify the copper and silver cathodes, clean and dry them, and after weighing them return them in their respective voltmeters. Pass a current of about 0.5A for 20 or 30 minutes, after which the cathodes should be removed, cleaned and dried, and reweighed.

  Compare the masses of copper and silver deposited. Note that care must be taken in removing the silver cathode from the solution as the metal does not always adhere well to the cathode.

  

  

  

  

  13.7. Factors affecting Electrolysis

  
  

  

  Procedure: let the current flow for 5 minutes

  Observe what happens at each electrode

  Questions:

  1. What product is formed on the cathode?

  2. What product is formed at the anode?

  3. Has the colour of the solution changed?

  4. Explain the observations in 3.

  

   Electrolysing of copper(II) sulphate solution using Copper electrodes Procedure: let the current flow for 5 minutes Observe what happens at each electrode

  Questions:

  1. What product is formed on the cathode?

  2. What product is formed at the anode?

  3. Has the colour of the solution changed?

  4. Explain the observations in 3.

  In an electrolysis where there are more than one species which can be discharged at the same electrode, only one of them is discharged at a time; for example, in an aqueous sodium chloride solution, we have four species that is,Na⁺ and Cl^- ions from sodium chloride and H⁺ and OH^- ions from water.

  During electrolysis Na⁺ ions and H⁺ ions migrate to the cathode while Cl^-ions and OH^- migrate to the anode.

  Now the question is, which species of ions will be discharged at the cathode and which ones will be discharged at the anode first?

  The factors which decide the selective discharge of ions are:

  • Nature of electrodes

  • Position of the ion in electrochemical series

  • Concentration

  • The state of the electrolyte

  

  13.7.2. Position of ion in Electrochemical Series

  When solving for the standard cell potential, the species oxidized and the species reduced must be identified. This can be done using the activity series. The table 13.2 is simply a table of standard reduction potentials in decreasing order. The species at the top have a greater likelihood of being reduced while the ones at the bottom have a greater likelihood of being oxidized. Therefore, when a species at the top is coupled with a species at the bottom, the one at the top will become reduced while the one at the bottom will become oxidized.

  During electrolysis of solution containing a mixture of ions, the ion lower in electrochemical series is discharged first in preference to the one high in the series.

  

  

  13.7.3. Concentration of electrolyte solution

  Increase of concentration of an ion tends to promote its discharge, for example in concentrated hydrochloric acid, containing OH^-(from water) and Cl^- as negative ions, the highly concentrated Cl^- is discharged in preference.

  However, if the acid is very dilute, some discharge of OH^- will also occur. It is important to know that as the acid is diluted, there will not be a point at which chlorine ceases to be produced and oxygen replaces it. Instead a mixture of the two gases will come off, with the proportion of oxygen gradually increasing.

  Another case in which the order of discharge according to the electrochemical series is reversed by a concentration effect is that of sodium chloride solution.

  In concentrated solution of sodium chloride called brine, the following reactions occur.

  

  Question 3

  A university student set up three different electrolytic cells. The substances that were electrolysed were NaCl (l), 0.05 M NaCl (aq) and 5.0 M NaCl (aq). Which of the following statements correctly describes the results of the experiment?

  a. The reactions occurring for the aqueous solutions will produce the same products at the anode and cathode.

  b. Chlorine gas is the major product when molten NaCl (aq) and 0.05 M NaCl (aq) are electrolysed.

  c. The pH at the cathode increases when solutions of NaCl are electrolysed.

  d. The only means by which different products can be produced for varying concentrations of NaCl is to alter the voltage.

  13.7.4. The state of the electrolyte

  The half reactions taking place at the electrode depends on whether the electrolyte is in a molten or an aqueous state, and if in aqueous state its concentration. For example, the electrode reactions that take place during the electrolysis of molten potassium iodide are:

  

  

  13.8. Application of electrolysis

  Activity 13.8 Copper-Plated Key

  Materials:

  • 1.5-volt D batterywith battery holder

  • Two alligator clip leads or insulated wire

  • Beakeror glass

  • Copper sulphate

  • Copper electrode(or coil of copper wire)

  • Brass key

  • Safety equipment

  Procedure:

  5. Prepare the key for copper-plating by cleaning it with toothpaste or soap and water. Dry it off on a paper towel.

  6. Stir copper sulphate into some hot water in a beaker until no more will dissolve. Your solution should be dark blue. Let it cool.

  3. Use one alligator clip to attach the copper electrode to the positive terminal of the battery (this is now theanode) and the other to attach the key to the negative terminal (now called the cathode).

  4. Partially suspend the key in the solution by wrapping the wire lead loosely around a pencil and placing the pencil across the mouth of the beaker. The alligator clip should not touch the solution.

  5. Place the copper strip into the solution, making sure it doesn’t touch the key and the solution level is below the alligator clip. An electrical circuit has now formed and current is flowing.

  6. Leave the circuit running for 20-30 minutes, or until you are happy with the amount of copper on the key.

  Question: Observe carefully electrolysis process and records what happened during the electrolysis process.

  Electrolysis has a number of important industrial applications. These include the extraction and purification of metals, electroplating and anodizing and the manufacture of other chemicals for example sodium hydroxide (NaOH).

  Extraction of metals Metals in group I and II of the periodic table cannot be reduced by chemical reducing agents; they are extracted from their fused halides by electrolysis. Sodium is obtained by electrolysis of molten sodium chloride in the Dawncell.Magnesium is obtained by electrolysis of MgCl₂, generated from dolomite and sea water.

  

  

  The electrolytic cell is an iron tank lined with carbon, which acts as the cathode. The anodes are blocks of carbon dipped into the electrolyte. The electrolyte is a solution of molten aluminum oxide in molten cryolite. Cryolite acts as a solvent to dissolve aluminium oxide and as an impurity to lower the melting point of aluminium oxide. The electrolytic cell is maintained at around 900°C

  

  

  

  

  

• UNIT 14: ENTHALPY CHANGE OF REACTIONS

  Key unit competency:

  To be able to design an experimental procedure to verify the enthalpy chang-es in a chemical reaction

  Learning objectives

  • Define heat of reaction, standard enthalpy change of combustion, enthalpy of neutralisation, enthalpy of solution, enthalpy of hydration and lattice enthalpy

  • Describe an experimental procedure in determination of heat of combustion

  • Explain the relationship between quantity of heat produced and mass of substance in combustion reaction

  • State Hess’s law of constant heat summation

  • State and explain the factors that affect the magnitude of lattice energy

  • Describe bond breaking as endothermic and bond making as exothermic

  • Develop practical experimental skills about enthalpy changes of reactions, interpreting results and drawing valid conclusions.

  • Carry out practical activities to determine enthalpy change of reactions (enthalpy change of combustion of ethanol, enthalpy change of neutralization).

  • Calculate the enthalpy change of combustion, neutralization and dissolution from experimental data

  • Deduce how Hess’s law is applied to Born-Haber cycle.

  • Construct Hess’s energy cycles and Born-Haber cycles from data obtained experimentally or provided.

  • Calculate the enthalpy changes of reactions using Hess’s law.

  • Use the standard bond energy to determine the standard enthalpy of reaction.

  • Relate the heat of hydration and lattice energy to heat of solution.

  • Respect of procedure during experiments of combustion and neutralization.

  • Appreciate the contributions of other scientists such as Hess, Born and Haber’s work.

  

  

  14.1. Definition of standard enthalpy of different reactions

  
  

  In thermodynamics, it is shown how energy, work, and heat are related. Every chemical reaction occurs with a concurrent change in energy. Before to embark the explanation of these chemical changes, some key terms have to be defined as follows.

  (i) Enthalpy change (ΔH)

  In thermodynamics, the heat of reaction also known as enthalpy of reaction is the change in the enthalpy (H) of a chemical reaction that occurs at a constant pressure. Enthalpy, H is a state function used to describe the heat changes that occur in a reaction under constant pressure. It is a state function as it is derived from pressure, volume, and internal energy, all of which are state functions. The enthalpy is a measurement of the amount of energy per mole either released or absorbed in a reaction.

  When a reaction is taking place in an open container, a quantity of heat which is proportional to the quantity of the matter present, will be released or absorbed.

  The flow of heat is the enthalpy change noted ΔH. The units of ΔH are kJ/mol or kcal/mol.

  (ii) Thermochemical equation

  A thermochemical equation is a balanced equation that includes the amount of heat exchanged (produced or absorbed).

  

  The rules of enthalpy change of reaction:

  a) The enthalpy change of a reaction is proportional to the amount of reactants that are involved in the reaction.

  

  

  

  

  b) Enthalpy of combustion

  The enthalpy of the combustion of a substance (element or compound) ΔHoc, is the enthalpy change which occurs when one mole of a substance undergoes complete combustion with oxygen in excess at 298 K and 1 atm.

  

  c) Enthalpy of neutralization

  The standard enthalpy of neutralization, ΔHon is the enthalpy change which occurs when one gram equivalent of an acid is neutralized by one gram equivalent of a base to produce a salt and water under the standard conditions of temperature and pressure.

  The equation of the neutralization reaction is: H+(aq) + OH- (aq)→H2O(l)

  

  

  

  14.2. Relationship between temperature and heat

  
  

  
  

  
  

  
  

  

  

  

  

  

  14.3. Experimental methods for finding the standard enthalpy of combustion reactions

  

  

  

  

  

  

  

  

  

  14.4. Experimental methods for finding the standard enthalpy of neutralisation reactions

  

  

  

  

  

  

  

  14.5. Hess’s law or Law of constant heat summation

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  14.12. Calculating enthalpy change of reaction using average bond enthalpies

  
  

  
  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  
  

• UNIT 15 : ENTROPY AND FREE ENERGY

  

  

  

  

  

  

  

• ASSESSMENT