Section: UNIT 1: PROPERTIES AND USES OF TRANSITION METALS | Chemistry

Section outline

UNIT 1: PROPERTIES AND USES OF TRANSITION
METALS
Key unit competence:
The learner should be able to explain the properties and uses of transition metals.
Learning objectives
At the end of this unit , students will be able to:
• Discuss qualitatively the propertie of transition elements;
• Explain the principle of ligan exchange;
• State the rules of naming complex ions and stereoisomers;
• Describe reactions of transition metals;
• State the use of transition metals;
• Relate the electronic configurations to special properties of transition metals;
• Predict the shape of the complex compounds of transition metal cations;
• Perform the confirmatory tests for transition metal ions.
Introductory Activity
The following photos show how some elements play a big role in our daily lives.
Observe these objects carefully.

Most of the metals in the periodic table belong in the d-block of transition metals.
They are hard and strong, and many of them are very familiar to us. For instance, zinc
is in brass instruments like trumpets and tubas. Have you ever heard of the element
“scandium” before? But you’ve interacted with it if you have ever ridden a bicycle.
1.1. Definition and electronic configuration of transition metals.
Activity 1.1
1. Write the electronic configuration of the following atoms and ions:
a. Ca(Z=20) b. Ca²⁺ c. Na(Z=11) d. Na⁺
2. Referring to the portion of periodic table in this book,
a. Write the electronic configuration of the elements from Sc to Zn.
b. Point out any difference between the electronic configuration of the
above elements and that of other elements in s and p blocks
3. Define the term transition metal.
According to IUPAC system, a transition metal is “an element whose atom has a
partially filled d sub-shell, or which can give rise to cations with an incomplete
d-orbitrals”.
Transition metals are located between groups 1& 2 (s-block) and group 13 (p-block)
on the periodic table. The elements are also called d-block elements because their
valence electrons are in d-orbitals.

The properties of transition elements are between the highly reactive metallic
elements of the s-block which generally form ionic compounds and the less reactive
elements of the p-block which form covalent compounds. Transition metals form
ionic compounds as well as covalent compounds.

The first 3 rows, i.e. period 4, period 5 and period 6, are called first transition series,
second transition series and third transition series respectively. The metals of the
first series are all hard and dense, good conductors of heat and electricity.
This block is known as the transition metals because some of their properties show a
gradual change between the active metals in s-block and p-block where non-metals
are found.
Electronic configuration is the arrangement of electrons in orbitals around the
nucleus. The electronic structure of the first transition series is shown in the table
below:

When building electronic structure of transition metals, 4s orbital is filled before 3d
orbitals.
The transition elements are stable when their d-orbitals are filled (d¹⁰) or when their
d-orbitals are half filled (d⁵).This explains the electronic structure of copper, [Ar]
4s¹3d¹⁰instead of
[Ar] 4s²3d⁹. The same applies for Cr: [Ar] 4s¹3d⁵ and not [Ar] 4s²3d⁴.
In order to attain that stability an electron can jump from 4s orbital to 3d orbital
because those two orbitals are close in energy.
This also explains why Fe²⁺ with 3d⁶ is easily oxidized to Fe³⁺ with 3d⁵ and Mn²⁺ with
3d⁵ is resistant to oxidation to Mn³⁺with 3d^4.
Transition metals form ions by losing electrons first from the 4s sub-shell rather than
the 3d sub-shell. Hence electronic configuration of Fe, Fe²+ and Fe³⁺ are the following:

The 4s electrons are removed before 3d electrons. This is because the 3d electrons
are inner while the 4s electrons are outer therefore the outer electrons (4s) have to
be removed before the inner electrons.

Checking up 1.1
1. Explain the difference between the electronic configuration of transition
elements and that of main group elements.
2. 2.Why d-block metals are so called transition metals?
1.2. Properties of the transition metals
1.2.1. Melting and boiling points
Activity 1.2 (a)
Experiment: Investigation of the melting point of transition metals compared
with s-block elements
Materials: Potassium or Rubidium metal and copper or iron metal, pair of tongs,
spatulas, bunsen burner and match box.
Procedure:
1. Take a half filled spatula of
a. Potassium (K) or Na, Rb, or Cs
b. Iron turnings or very small piece of copper sheet (which can fit on a
spatula)
2. Heat both spatulas on the Bunsen burner flame
3. Write down the observations
4. What can you conclude about your findings?
The melting points and the molar enthalpies of fusion of the transition metals are
both high in comparison to main group elements. Most of the transition metals have
melting points above 1000^oC; mercury is liquid at room temperature.
This is due to the high number of valence electrons that increases the electrostatic
attraction force between those electrons and the metallic cations, hence increasing
the strength of the metallic bond and the melting point.
Table 1.1: Melting and boiling points of the 1^stseries of Transition Metals

Checking up 1.2 (a)
Compare and comment on the melting points of transition metals and those of
s-block metals.
1.2.2. Densities and atomic/metallic radii
Activity 1.2 (b)
Procedure for practical:
1. a. Take a magnesium ribbon and a copper foil of the same size (if
possible you may use their turnings)
b. You weigh those two samples using an electronic balance. And record
their masses
2. a. Take aluminum foil and copper foil of the same size (if possible you
may use their turnings)
b. You weigh those two samples using an electronic balance. Record their
masses
3. Comment on your observations by explaining why their masses are
different and yet they have the same size.
4. Use the internet or any book or even this one to interpret the data
given about metallic radii of the first series transition metals. From your
research, compare metallic radii of transition metals and those of main
group elements.
The transition elements are much denser than the s-block elements and show in
general a gradual increase in density from left to right in a period as you can see
below from scandium to copper. This trend in density can be explained by a decrease
in metallic radii coupled with the relative increase in atomic mass
Table 1.2: Density/g cm ^-3of the first transition series

Table1.3: Metallic radii of the first transition series

Checking up 1.2 (b)
The metallic radius of vanadium is smaller than that of titanium. Explain this
statement.
1.2.3. Ionization energies
Activity 1.2 (c)
Use this book or any other source (textbook or search engine) to interpret/ analyze
the summary about ionization energies of the transition metals (first series). From
your findings, compare

a. Ionization energies of those transition metals.
b. Ionization energies of transition metals and those of main group elements

The ionization energy of transition metals is related to the energies of its d orbitals,
its ease of oxidation, and its basicity.In simplest terms, the greater the ionization
energy of a metal, the harder it is to pull an electron from it.

As the number of protons increases across a period (or row) from left to right of
the periodic table, the first ionization energies of the transition-metal elements are
relatively the same, while that for the main-group elements increases.

In moving across the series of metals from scandium to zinc, a small change in the
values of the first and second ionization energies is observed. This is due to the buildup
of electrons in the immediately underlying d-sub-shells that efficiently shields
the 4s electrons from the nucleus and minimizing the increase in effective nuclear
charge from element to element.
Table 1.4: First, second and third ionization energies of 1st Series Transition
metals /kJ mol^-1
The figure 1.3 below shows the first iazonisation energies for transition metals of 1st,
2nd and 3rd rows (series).

In general, ionization energy increases as we move from left to right across the
period. Notable dips occur at row 1, group 10 (Ni) and row 3, group 7 (Re).

Checking up 1.2 (c)
Briefly explain the following observations:
a. The first ionization energy of cobalt is only slightly larger than the first
ionization energy of iron.
b. The third ionization energy of iron is much lower than the 3rd ionization
energy of Mn.
1.2.4. Transition elements have variable oxidation states
Activity 1.2 (d)
Use this book or any other source (textbook or search engine) to
a. Explain the term oxidation number
b. Compare the oxidation numbers of transition metals (first series) and those
of main group elements.
c. Analyze the stability of ions formed by transition metals (first series).

Oxidation state is a number assigned to an element in chemical combination which
represents the number of electrons lost or gained.The transition elements from
Titanium to Copper all form ions with two or more oxidation states. In most cases,
this is the result of losing the two electrons of 4s orbital and electeons in 3d orbitals.
The 4s electrons are lost first because they are in the highest energy level. However,
because the 3d and 4s energy level are so close in energy, the 3d electrons can also
be lost when an atom forms a stable ion. The common oxidation states shown by the
first transition series are:
Table 1.5: The oxidation states shown by the transition metals (series)

• The common stable oxidation states for those transition metals with variable
oxidation states are bolded and underlined.
• The oxidation state corresponding to a full or half-filled d-orbital is energetically
stable. For example, Fe³⁺ is more stable than Fe²⁺and Mn²⁺ is more stable than
Mn^3+.
• However, in most compounds and solutions, copper exist as Cu²⁺ ion rather
than Cu⁺ ion. Meaning that the former is more stable than the latter. The
explanation of this is beyond this level.
Checking up 1.2 (d)
Which gaseous ion is more stable, Mn²⁺ or Mn³⁺? Explain why.
1.2.5. Most transition metals and their compounds have high ability of being
catalyst
Activity 1.2 (e)
Practicals:
1. Preparation of oxygen using hydrogen peroxide, H₂O₂, without a catalyst
a. Put 10 mL of H₂O₂ in a conical flask (Pyrex preferably)
b. Heat the conical flask for about 5 min
c. Write down the observation in A.
2. Preparation of oxygen using hydrogen peroxide, H₂O₂, with MnO₂ as a
catalyst
a. Put10 mL of H₂O₂ in a conical flask(Pyrex preferably)
b. Put a very small amount of MnO₂ in the conical flask
c. Heat the conical flask for about 5 min
d. Write down the observation in (B)
Question: What is the role of MnO₂ in the above experiment?
A catalyst is a substance that can speed up (positive catalyst) or that can slow down
(negative catalyst) the rate of reaction and is found unchanged at the end of the
reaction. But generally the term catalyst is used for the substance that helps in
accelerating the rate of the reaction. A catalyst that speeds up the reaction provides
another pathway with lower activation energy.
In some catalytic process, transition metal ions undergo changes in their oxidation
states but are regenerated at the end of the reaction.
The reasons for transition metals to work as catalysts:
• Presence of empty d orbitals which enable transition metal ions (or atoms) to
form temporary bonds with reactant molecules at the surface of a catalyst and
weakens the bond in the reactant molecules
• Variable oxidation states which allow them to work as catalysts in the reactions
involving the transfer of electrons.
Table 1.6: Reactions catalysed by transition metals

Checking up 1.2 (e)
Explain why s-block metals and their compounds are not used as catalysts
1.2.6. Most transition metal ions are paramagnetic
Activity 1.2(f)
Given the following materials:

1. Organize yourself in to group to find the objects shown in the photo
above.
2. Using a magnet, classify the above materials into two groups as shown in
the table below.
Objects attracted by a magnet
Objects not attracted by a magnet
3. Research, using any relevant source (textbook or internet), to identify in
which metal the objects A to E are made
4. Research to know why some objects are attracted by a magnet while
others are not
Paramagnetism is a property of substances to be attracted in a magnetic field.
Substances which are not attracted (i.e slightly repelled) in a magnetic field are
said to be diamagnetic.Transition metal ions show paramagnetism because of the
presence of unpaired electrons in their 3d arbitrals.
The greater the number of unpaired electrons, the stronger the paramagnetism;
that is the reason why:
• Fe³⁺ is more paramagnetic than Fe²⁺because Fe³⁺ has five unpaired electrons
while Fe²⁺ has four unpaired.
• Sc³⁺ and Zn²⁺ are not paramagnetic, they are diamagnetic because they do not
have unpaired electrons.
Other examples of paramagnetic substances are: Cr, Mn, CuSO_4, Fe, Co, Ni, Pt.
Examples of diamagnetic substances are: Zn, Cu⁺, Au⁺, TiO_2.

Checking up 1.2 (f)
Predict whether the following substances are paramagnetic or not. Explain
a. CuSO₄
b. Co
c. Ca
d. Cr
1.2.7. Formation of alloys
Activity 1.2 (g)
Observe the trophies/or other objects made in the materials below and compare
their appearances with the elements from which they are derived.
a. Bronze with copper
b. Stainless steel with iron
c. Pewter and copper
You can use the internet, books (including this one) or any other relevant source
to find the figures of the above objects, the elements they are made from and
their uses.
An alloy is a homogenious solid mixture (solid solution) made by combining two or
more elements where at least one is a metal.

Importance of alloying:
• Increase of the strength of a metal,
• Resistance to corrosion,
• Gives to the metal a good appearance
Generally, alloys are needed and used to improve the quality of the required
material. For example, brass (alloy of zinc and copper) is much stronger than either
pure copper or pure zinc. Pure gold is too soft to be used in some applications.
Table 1.7: The properties and uses of some common alloys formed by transition
metals (first series)

Checking up 1.2 (g)
1. Explain why alloys are said to be solid solutions.
2. Give the importance of alloying
1.2.8. Formation of complex ions
Activity 1.2
Use this book or any other source (library textbook or internet) to analyze and
discuss on the following. You have to take note on what to be presented to share
with your colleagues and teacher.
a. What is a ligand?
b. State the types of ligands
c. The geometry of complexes

A complex or coordination compound is a chemical species made of a central metal
(cation or neutral) bonded to other chemical species called ligands by coordination
or dative bonds. A complex may be neutral, positively or negatively charged.
Transition metal form complexes because of:
• Their small and highly charged ions,
• The presence of vacant (empty) d-orbitals which can accommodate lone pair
of electrons donated by other groups (ligands)
The general formula of a complex is: [ML_n]^y
Where:
• M-metal ion or atom
• L-Ligand
• n-the number of ligands surrounding the metal
• y-the charge of the complex; [ML_n] indicates a neutral complex.
-Coordination number of a complex: is the number of coordinate bonds on the
central metal in a complex.
-Ligand: It is a species (anion or a molecule) that is bonded to a central metal ion
or atom in a complex. A ligand should have at least one lone pair of electrons
to form a coordinate bond.
Ligands are classified depending on the number of sites at which one molecule
of a ligand is coordinated to the central metallic atom; the ligands are classified
as monodentate (or unidentate), ambidentate and polydentate (or multidentate)
lingards.

a. Monodentate ligands
The ligands which have only one donor atom or are coordinated through one
electron pair are called monodentate ligands because they have only one tooth with
which to attach themselves to the central cation or atom. Such ligands are coordinated
to the central metal at one site or by one metal-ligand bond only. These ligands may
be neutral molecules or in anionic form.
The table below provides examples of some monodentate ligands.
Table 1.8: some monodentate ligands

Ligands that can use different sites to coordinate to the central metal are called
“ambidentate”: e.g. CN- and NC-(see table above).

Notice that a ligand with a donor atom that possesses 2 lone pairs of electrons, such
as H₂Ö:, is not bidentate, since it cannot use both lone pairs simultaneously to bind
to the metal because of the steric effect.

b. Polydentate ligands
These may be bidentate, tridentate, tetradentate, pentadentate, and hexadentate
ligands if the number of donor atoms present in one molecule of the ligand attached
with the central metallic atom is 2, 3, 4, 5, and 6 respectively. Thus one molecule of
these ligands is coordinated to the central metallic atom at 2, 3, 4, 5, and 6 sites
respectively. In other words, we can say that one molecule of these ligands makes 2,
3, 4, 5, and 6 metal-ligand coordinate bonds respectively.

• Tetradentate

• Hexadentate
The structure shows that it has two neutral N- atoms and four negatively charged Oatoms
as its donor atoms which can form coordinate bonds with a transition metal
ion.

The complex ions which form between polydentate ligands and cations are known
as chelates or chelated complexes.

In general, polydentate ligands form more stable complexes than monodentate
ligands. The stability of complex is much enhanced by chelation. A polydentate
ligand can hold the central cation more strongly.
Examples of complexes:
• Copper (II) ions have a coordination number of four in most of its complexes:
[Cu(H₂O)₄]^2+, [Cu(NH₃)₄]²⁺, [CuCl₄]²⁺, [Cu(NH₂-(CH₂)₂-NH₂)₂]²⁺, …

• Most of ions have coordination number of 6.
[Cr(H₂O)₆]³⁺ , [Cr(NH₃)₆]³⁺ , [Cr(H₂O)₄Cl₂]- , …
• Very few ions have a coordination number of 2: [Ag(NH₃)₂]+, [Ag(CN)₂]-, [CuCl₂]-,
…
Geometry of complexes
Complexes have a variety of geometries or shapes, but the most common geometries
are the following:
• Complexes with coordination number 2 adopt a linear shape. Example:
[Ag(NH₃]₂+: [H₃N-Ag-NH₃]+

The complexes having coordination number of 2 are linear since minimises ligand
repulsion.

• Complexes with coordination number 4 generall adopt a tetrahedral shape.
But few of them can form a square planar shape.
Examples:
[Zn(NH₃)₄]²⁺, [NiCl₄]^2- and some few others adopt a square planar shapes, examples:
[Cu(NH₃)4]²⁺ , [Ni(CN)₄]^2-,[CuCl₄]^2-,[CoCl₄]^2-,…

The square plannar geometry is characteristic of transition metal ions with eight d
electrons in the valence shell, such as platinum(II)and gold(III).
Copper (II) and cobalt (II) ions have four chloride ions bonded to them rather than
six, because the chloride ions are too big to fit any more around the central metal
ion.

• Complexes with coordination number 6 adopt an octahedral shape.
Example: [Cr(NH₃)₆]^3+.
These ions have four of the ligands in one plane, with the fifth one above the plane,
and the sixth one below the plane.

Checking up 1.2
1. What do you understand by :
a. Coordination number.
b. Ligand.
2. Give the main types of ligands and give an example for each
3. Say if the following statement is correct or wrong and justify: The
coordination number equals the number of ligands bonded to the central
metal.

1.2.9. Many transition metal ions and their compounds are coloured
Activity 1.2 (i)
Experiment 1: Observation of the colors of transition elements
Apparatus: Test tubes, droppers, spatula, test tube holders.
Chemicals: NaCl, CaCl₂, FeSO₄, Fe₂ (SO₄)₃, KMnO₄, K₂Cr₂O₇ ,distilled water, Cr₂(SO₄)₃.
1. What are the colours of the compounds above?
2. Determine the oxidation states of each metal in the above compounds?
3. a. Take an endful spatula of each product given above and put each in a test
tube.
b. Put 10 mL of distilled water in each test tube.
c. Write down the colours of solutions formed and conclude.

Experiment 2: Investigation of ligand exchange reactions involving copper (II) ions, Cu2+
Apparatus: Test tubes, droppers, spatula, test tube holders.
Chemicals: Copper (II) sulphate, concentrated hydrochloric acid, concentrated ammonia
solution and distilled water.
Procedure:
1. Use a spatula to place a small amount of anhydrous copper (II) sulphate in a test
tube.
2. Add 10 drops of distilled water to the anhydrous copper (II) sulphate and shake
3. To the test tube in step 2, add concentrated ammonia solution drop by
drop while shaking the test tube until there is no further change. Record all
observations.
4. Repeat steps 1 and 2
5. To the test tube from step 4, add concentrated hydrochloric acid drop by drop
while shaking until there is no further change. Record all observations.

Points for discussion:
1. What happens when anhydrous copper (II) sulphate is dissolved in water?
2. Describe what is observed when concentrated ammonia is added dropwise to
an aqueous solution of copper (II) sulphate.Write balanced equations for each
observation if possible
3. Describe what happens when concentrated hydrochloric acid is added to an
aqueous solution of copper (II) sulphate. Write balanced equation(s) for the
observation(s) made.
4. State any other possible observation(s) for this experiment.

The formation of colored ions by transition elements is associated with the presence
of incompletely filled 3d orbitals.

This property has its origin in the excitation of d electrons from lower energy
d-orbitals to higher energy. In fact, when the central metal is surrounded by ligands,
these cause d orbitals to be split into groups of higher and lower energy orbitals.
When electrons fill d-orbitals, they fill first of all the lower energy orbitals; if there is
free space in higher energy d-orbitals, an electron can be excited from lower energy
d-orbitals to higher energy d-orbitals by absorbing a portion of light corresponding
to a given colour, the remaining color light is the white light minus the absorbed
colour.

When a coloured object is hit by white light, the object absorbs some colour and
the colour transmitted or reflected by the object is the colour which has not been
absorbed. The observed colour is called complementary colour.
When a metal cation has full d-orbitals, such as Cu⁺or Zn²⁺or no electron in d orbital,
such as Sc³⁺.

Table1.9: Complementarities of colors observed and absorbed when light is
emitted

The colour of a particular transition metal ion depends upon two factors:
• The nature of the ligand

The principle of ligand exchange
Complexing reactions involve competitions between different ligands for central
metal. A more powerful ligand displaces a less powerful ligand from a complex.
During the process there is a change in colour.
Here below is a list of some ligands in increasing order of strength.

The above series are called the spectrochemical series and shows that cynide ion
and carbon monoxide are very strong ligands
The stability of a complex ion is measured by its stability constant. The higher the
stability constant of a complex, the more stable is the complex.
Checking up 1.2 (i)
Predict whether each of the following ion forms coloured compounds and explain
your reasoning: Fe²⁺, Mn⁷⁺, K⁺
1.3. The anomalous properties of Zinc and Scandiu
Activity 1.3
From the information you have learnt about the properties of transition metals,
Suggest the difference between the properties of Zn and Sc and other transition
metals. You can consult different sources (books or internet) to provide enough
information.

On the basis of the properties of transition metals, zinc and scandium are not
considered as typical transition metals even though they are members of the d-block.
Zinc:
• It has a complete d-orbital.
• Zinc forms only the colourless Zn²⁺ ion, isoelectronic with the Ga³⁺ion, with 10
electrons in the 3d subshell.
• Zinc and its compounds are not paramagnetic
Scandium:
• Has one oxidation state,+3
• Sc forms only the colourless Sc³⁺ion, isoelectronic with the Ca²⁺ ion, with no
electrons in the 3d subshell.
• Its compounds are diamagnetic
• It forms compounds containing ions with a completely empty 3d subshell.

Checking up 1.3
Give any one property by which Zn differs from Sc

1.4. Naming of complex ions and isomerism in of transition
metal complexes
1.4.1. Naming of complex ions
Activity 1.4 (a):
1. Name the following molecules and explain the basis /principle used to
name them.
a. CaBr₂
b. CCl₄
c. SF₆
2. Analyze the IUPAC rules for naming complex ions in the summary in this
book or using any other source (textbook or search engine) and apply
them by naming the following:
a. [CuCl₄]^2-
b. [Cu(H₂O)₆]²⁺
c. [Cr(NH₃)₃(H₂O)₃]Cl₃
d. [Pt(NH₃)₂Cl₂]
e. (NH₄)₂[Ni(C₂O₄)₂(H₂O)₂]

Naming molecules requires the knowledge of certain rules, such as how to name
cations, anions, where to start from when both a cation and an anion are combined
in an ionic molecule or when two non metals are combined in a covalent molecule.
Like other compounds, complex compounds/ions are named by following a set of
rules. You are familiar with some of them and the new ones can be understood and
applied easily.
1. In simple metal compounds, the metal is named first then the anion.
Example: CaCl₂: calcium chloride
2. In naming the complex:
a. Name the ligands first, in alphabetical order, then the metal atom or cation,
followed by its oxidation state written between brackets as Roman number,
though the metal atom or cation is written before the ligands in the chemical
formula.
Example: [CuBr₄]^2-: Tetrabromocuprate (II) ion
The names of some common ligands are listed in the table below:
Table 1.10: Names of common ligands

b. Greek prefixes are used to indicate the number of each type of ligand in the
complex:
The numerical greek prefixes are listed in the following table:
Table 1.11: Greek numerical prefixes

c. After naming the ligands, name the central metal.
• If the complex bears a positive charge (cationic complex), the metal is named
by its usual name.
Example: Cu: Copper Pt: Platinum

If the complex bears a negative charge (anionic complex), the name of the metal
ends with the suffix –ate
Example: Co in a complex anion is called cobaltate and Pt is called platinate.
For some metals, the Latin names are used in the complex anions e.g. Fe is called
ferrate (not ironate). See table below:
Table 1.12: Latin names of some transition metals in anionic complexes

1. For historic reasons, some coordination compounds are called by their common
names.
Example: Fe(CN)₆³- and Fe(CN)₆^4- are named ferricyanide and ferrocyanide
respectively, and Fe(CO)₅ is called iron carbonyl.
2. To name a neutral complex molecule, follow the rules of naming a complex
cation. Example: [Cr(NH₃)₃Cl₃]: triamminetrichlorochromium (III)
You can have a compound where both the cation and the anion are complex ions.
Notice how the name of the metal differs even though they are the same metal ions.
Remember: Name the cation before the anion.
Example: [Ag(NH₃)₂][Ag(CN)₂] is diamminesilver(I)dicyanoargentate(I)
Note that:
• The names are written as a one word: Tetraamminecopper (II), not Tetraammine
copper (II).
• Complex ions formula is written between square brackets and the charge of the
ion as superscript outside the brackets: [Cu(NH₃)₄]²⁺. When oppositely charged
ions approach the complex ion, a neutral molecule can be obtained:
[Cu(NH₃)₄]²⁺2Cl^- or simply, [Cu(NH₃)₄]Cl₂: tetraamminecopper(II)chloride.
The ions outside the square brackets are known as “counter ions”.

Checking Up 1.4 (a):
1. Complete the table below using the names of the given metals when they
are in anionic complexes
Element
Name in an anionic complex

2. Give the systematic names for the following complex ions/compounds:
a. [Cr(NH₃)₃(H₂O)₃]³⁺
b. [Co(H₂NCH₂CH₂NH₂)₃]₂(SO₄)₃
c. K₄[Fe(CN)₆]
d. Fe(CO)₅
1.4.2. Isomerism in complexes
Activity 1.4 (b):
1. Discuss on the following questions:
a. What do you understand by the term “isomerism”?
b. Is there any relationship between isomers and isomerism?
c. Give examples of molecules that can exist as isomers and explain their
isomerism
2. Read and discuss the summary below to understand how complex ions/
compounds exhibit isomers
3. Present your findings to your colleagues and teacher to share your
understanding.

Isomers are chemical species that have the same molecular formulal, but different
molecular structures or different arrangements of atoms or groups of atoms in
space. Isomerism among transition metal complexes arises as a result of different
arrangements of their constituent ligands around the metal.
The diagram below shows the different categories of isomerism in transition metal
complexes.

In this unit, we are specifically concerned with ‘stereoisomerism’ which gives rise to
isomers known as “stereoisomers”. Stereoisomers have the same structural formulal
but different arrangements of ligands in space.
They are classified in two categories: geometrical isomers and optical isomers.

1. Geometrical isomers
Coordination complexes, with two different ligands in the cis and trans positions
from a ligand of interest, form isomers.
For example, the square planar, diammine dichloroplatinum (II) Pt(NH₃)₂Cl₂),can be
presented as follows:

The octahedral [Co(NH₃)₄Cl₂]⁺ ion can also have geometrical isomers.

Different geometrical isomers are different chemical compounds. They exhibit
different properties, even though they have the same formula. For example, the two
isomers of [Co(NH₃)₄Cl₂]NO₃ differ in color; the cis form is violet, and the trans form is
green. Furthermore, these isomers have different dipole moments, solubilities and
reactivities.

2. Optical isomers (enantiomers)
Optical isomers are non-superimposable mirror images of each other. A classic
example of this is your two hands (left and right); hold them face-to-face: one is the
mirror image of the other. Now try to superimpose them one over another: they
are non-superimposable (only the middle fingers superimpose one over the other.
Chemical compounds that behave like the hands are called “chiral”, in reference to
the Greek word for hands.

Optical isomers are very important in organic and biochemistry because living
systems often incorporate one specific optical isomer and not the other.

Unlike geometric isomers, optical isomers have identical physical properties (boiling
point, polarity, solubility, etc.). Optical isomers differ only in the way they affect
polarized light and how they react with other optical isomers.

1. For coordination complexes, many coordination compounds such as
[M(en)₃]ⁿ⁺ [in which Mⁿ⁺ is a central metal ion such as iron(III) or cobalt(II)]
form enantiomers, as shown in figure below.These two isomers will react
differently with other optical isomers. For example, DNA helices are optical
isomers, and the form that occurs in nature (right-handed DNA) will bind
to only one isomer of [M(en)₃]ⁿ⁺ and not the other.

Checking up 1.4 (b):
1. The geometric isomer of [Pt(NH₃)₂Cl₂] is shown in the figure below. Draw
the other geometric isomer and give its full name.

2. Draw the ion trans-diaqua-trans-dibromo-trans-dichlorocobaltate (II).
3. Sketch the arrangement of bonds in the complexes
a. Hexaaquacobalt(III) ion
b. Hexacyanoferrate (III) ion
c. Diamminesilver (I) ion
d. The complex compound tetracarbonylnickel (0).
4. The compound [NiCl₂(NH₃)₂] has cis-trans isomers. These have a complex
non-ionic structure.
a. Does [NiCl₂(NH₃)₂] have a tetrahedral or a square-planar structure?
Explain your answer.
b. Draw the cis and trans isomers for [NiCl₂(NH₃)₂].
5. Early in the 20^thcentury, the German scientist Werner succeeded in
clarifying the situation concerning the five compounds of PtCl₄^- and
ammonia. The properties of these compounds are listed in the table
below.

a. What is the oxidation state of Pt in each of the compounds A-E?
b. The co-ordination number of Pt in each compound is six. Write a right formula for each
of the five compounds. Show the complex ion and the other ions and/or molecules
present.
c. Each of the compounds forms an octahedral complex ion. Draw the structures for the
complex ions in A, B, C, and D.
d. Which of the complex ions in (c) have isomers?
1.5. The Chemistry of individual transition metals
Activity 1.5
Using the library and internet or other textbooks, make your own research and make
presentation of the results of your research:
1. On how each of the transition metals (first series) reacts with each of the
following substances
a. Oxygen
a. Water
c. Hydrochloric acid
d. Sodium hydroxide
e. Chlorine
2. On the uses and their corresponding properties for each of the above
transition metals.
1.5.1. Scandium
Scandium is a silvery-white solid. It melts at 1539^oC and boils at 2748^oC. Its density
is about 3.0.
1. Chemical reactions
a. Reaction of scandium with air
Scandium tarnishes in air, and burns readily, forming scandium (III) oxide, Sc₂O₃.
4 Sc(s) + 3 O₂(g) ——→ 2 Sc₂O₃(s)

b. Reaction of scandium with water
When finely divided, or heated, scandium dissolves in water, forming Sc (III)
hydroxide and hydrogen gas, H₂.
2 Sc(s) + 6 H₂O(l) ——→2 Sc(OH)₃(aq) + 3 H₂(g)

c. Reaction of scandium with acids
Scandium dissolves readily in dilute hydrochloric acid, forming Sc(III) ions and
hydrogen gas, H_2.
2 Sc(s) + 6 HCl(aq) ——→2 Sc³⁺(aq) + 6 Cl−(aq) + 3 H₂(g)

d. Reaction of scandium with halogens
Scandium reacts with the halogens, forming the corresponding Sc(III) halides
2 Sc(s) + 3 F₂(g)——→ 2 ScF₃(s)
2 Sc(s) + 3 Cl₂(g) ——→2 ScCl₃(s)
2 Sc(s) + 3 Br₂(g)——→ 2 ScBr₃(s)
2 Sc(s) + 3 I₂(g)——→ 2 ScI₃(s)

2. Uses
• Scandium has as low density (2.99 g/cm³) asaluminium (2.7 g/cm³) but a much
higher melting point.
• An aluminium-scandium alloy has been used in fighter planes, high-end
bicycle frames and baseball bats.
• Scandium iodide is added to mercury vapour lamps to produce a highly
efficient light source resembling sunlight. These lamps help TV cameras to
reproduce colour well when filming indoors or at night-time.

1.5.2. Titanium
Titanium is a gray, solid with a density of about 4.50. It melts at 1667^oC and boils at
3285^oC.
1. Chemical reactions
a. Reaction of titanium with air
Titanium does not react with air under normal conditions. If brought to burn,
titanium will react with both oxygen, O₂, and nitrogen, N₂.

b. Reaction of titanium with water
Titanium does not react with water, under normal conditions. If the water is heated
to steam, it will react with titanium, forming titanium(IV) oxide, TiO₂, and hydrogen,
H²

c. Reaction of titanium with acids
Titanium does not react with most acids, under normal conditions. It will react with
hot hydrochloric acid, and it reacts with HF, forming Ti(III) complexes and hydrogen
gas, H₂.

d. Reaction of titanium with bases
Titanium does not appear to react with alkalis, under normal conditions, even when
heated.
e. Reaction of titanium with halogens
Titanium reacts with halogens, when heated, forming the corresponding titanium(IV)
halides

2. Uses
• Titanium is as strong as steel but much less dense. It is therefore important as
an alloying agent with many metals including aluminium, molybdenum and
iron. These alloys are mainly used in aircraft, spacecraft and missiles because
of their low density and ability to withstand extremes of temperature. They are
also used in golf clubs, laptops, bicycles and crutches.
• Power plant condensers use titanium pipes because of their resistance to
corrosion. Because titanium has excellent resistance to corrosion in seawater,
it is used in desalination plants and to protect the hulls of ships, submarines
and other structures exposed to seawater.
• Titanium metal connects well with bone, so it has found surgical applications
such as in joint replacements (especially hip joints) and tooth implants.
• The largest use of titanium is in the form of titanium (IV) oxide. It is extensively
used as a pigment in house paint, artists’ paint, plastics, enamels and paper.
It is a bright white pigment with excellent covering power. It is also a good
reflector of infrared radiation and so is used in solar observatories where heat
causes poor visibility.

1.5.3. Vanadium
Vanadium is a grey, solid with a density of about 6.11. It melts at 1915^oC and boils at
3350^oC. It is insoluble in water at room temperature.

1. Chemical reactions
a. Reaction of vanadium with air
Vanadium metal reacts with excess oxygen, O₂, upon heating to form vanadium (V)
oxide, V₂O₅. When prepared in this way, V₂O₅ is sometimes contaminated by other
vanadium oxides.

b. Reaction of vanadium with water
Vanadium does not react with water, under normal conditions.

c. Reaction of vanadium with bases
Vanadium metal is resistant to attack by molten alkali.
In strong alkaline solutions (pH > 13), Vanadium (V) exists as colourless
orthovanadate ions, VO₄³⁻.

d. Reaction of vanadium with halogens
Vanadium reacts with fluorine, F₂, when heated, forming vanadium (V) fluoride

2. Uses
• About 80% of the vanadium produced is used as a steel additive. Vanadium-
steel alloys are very tough and are used for spanners, armour plate, axles,

piston rods and crankshafts. Less than 1% of vanadium, and as little chromium,
makes steel shock resistant and vibration resistant. Vanadium alloys are used
in nuclear reactors because of vanadium’s low neutron-absorbing properties.
• Vanadium (V) oxide is used as a pigment for ceramics and glass, as a catalyst
and in producing superconducting magnets.

1.5.4. Chromium
Chromium is a silver gray metal with density of about 7.14. It melts at 1900^oC and
boils at 2690^oC. Chromium is insoluble in water at room temperature.

1. Chemical reactions
a. Reaction of chromium with air
Chromium metal does not react with air at room temperature. Heated clean
chromium is oxidized superficially in air to green solid, chromium (II) oxide.

b. Reaction of chromium with water
Normally, Chromium metal does not react with water at room temperature. When
red hot, it reacts with steam to form chromium (II) oxide.

c. Reaction of chromium with acids
Metallic chromium dissolves in dilute hydrochloric acid forming Cr(II) and hydrogen
gas, H₂. In aqueous solution, Cr(II) is present as the complex ion [Cr(OH₂)₆]²⁺.

Similar results are seen for sulphuric acid but pure samples of chromium may be
resistant to attack.
Chromium metal is not dissolved by nitric acid, HNO₃ but is passivated instead.

d. Reaction of chromium with hydroxide ions
Chromium dissolves rapidly in hot concentrated aqueous alkali forming a blue
solution containing chromium (II) ion and hydrogen gas is evolved.

Similar results are seen for sulphuric acid but pure samples of chromium may be
resistant to attack.
Chromium metal is not dissolved by nitric acid, HNO₃ but is passivated instead.

d. Reaction of chromium with hydroxide ions
Chromium dissolves rapidly in hot concentrated aqueous alkali forming a blue
solution containing chromium (II) ion and hydrogen gas is evolved.

e. Reaction of chromium with halogens
Chromium reacts directly with fluorine, F₂, at 400°C and 200-300 atmospheres to
form chromium (VI) fluoride, CrF₆.

Under milder conditions, chromium (V) fluoride, CrF₅, is formed.

Under milder conditions, chromium metal reacts with the halogens to form
chromium tri halides or chromium (III) halides:
2. Uses
• Chromium is used to harden steel, to manufacture stainless steel (resists to
corrosion) and to produce several alloys.
• Chromium plating can be used to give a polished mirror finish to steel.
Chromium-plated car and lorry parts, such as bumpers, were once very
common. It is also possible to chromium plate plastics, which are often used
in bathroom fittings.
• About 90% of all leather is tanned using chromium. However, the waste
effluent is toxic so alternatives are being investigated.
• Chromium compounds are used as industrial catalysts and pigments (in
bright green, yellow, red and orange colours). Rubies get their red colour from
chromium, and glass treated with chromium has an emerald green colour.
• Chromium (IV) oxide is used in magnetic tapes for sound/video recording.
• Chromium is used in the control of cholesterol and help insulin sugar control
in blood.

1.5.5. Manganese
Manganese is a grey-white solid with a slightly red colour. Its density is about
7.44^oC. Manganese melts at 1244^oC and boils at 2060^oC. It is insoluble in water but
soluble in diluted acids, at room temperature.

1. Chemical reactions
a. Reaction of manganese with air
Manganese is not very reactive with air. The surface of manganese lumps oxidizes
a little. Finely divided manganese metal burns in air. In oxygen the oxide Mn₃O₄ is
formed and in nitrogen the nitride Mn₃N₂ is formed.

b. Reaction of manganese with water
Manganese reacts slowly with water to form manganese (IV) oxide:

c. Reaction of manganese with acids
Manganese dissolves readily in dilute sulphuric acid, forming a colorless solution of
Mn(II) ions and hydrogen gas, H₂.

d. Reaction of manganese with halogens
Manganese reacts with the halogens, forming the corresponding manganese (II)
halides. For fluoride, manganese (III) fluoride is also formed.

2. Uses
• Manganese is too brittle to be of much use as a pure metal. It is mainly used
in alloys, such as steel. Steel contains about 1% manganese, to increase the
strength and also improve workability and resistance to wear. Manganese
steel contains about 13% manganese. This is extremely strong and is used for
railway tracks, safes, rifle barrels and prison bars.
• Drinks cans are made of an alloy of aluminium with 1.5% manganese, to improve
resistance to corrosion. With aluminium, antimony and copper it forms
highly magnetic alloys.
• Manganese (IV) oxide is used as a catalyst, a rubber additive and to decolourise
glass that is coloured green by iron impurities. Manganese (IV) oxide is a powerful
oxidising agent and is used in quantitative analysis. It is also used to make
Fertilizers and ceramics.
• Manganese sulphate is used to make a fungicide.
1.5.6. Iron
Iron is a grey to black, odourless metal with density 7.874. It melts at 1535 ^oC and
boils at 2750 ^oC.
1. Chemical reactions
a. Reaction of iron with air
Iron reacts with oxygen, O_2, forming Fe (II) and Fe(III) oxides. The oxide layer does not
passivate the surface. Finely divided iron, e.g. powder or iron wool, can burn:

b. Reaction of iron with water
Air-free water has little effect upon iron metal. However, iron metal reacts in moist
air by oxidation to give a hydrated iron oxide. This does not protect the iron surface
to further reaction since it flakes off, exposing more iron metal to oxidation. This
process is called rusting.

c. Reaction of iron with acids
Iron metal dissolves readily in dilute sulphuric acid in the absence of oxygen forming
Fe(II) ions and H₂. In aqueous solution Fe(II) is present as the complex [Fe(H₂O)₆]²⁺.

Concentrated nitric acid, HNO₃, reacts on the surface of iron and passivates the
surface (makes it unreactive).
d. Reaction of iron with halogens
Iron reacts with excess of the halogens, F₂, Cl₂, and Br₂, to form Fe(III) halides.

2. Uses
• Iron is an enigma – it rusts easily, yet it is the most important of all metals. 90%
of all metal that is refined today is iron. Most is used to manufacture steel, used
in civil engineering (reinforced concrete, girders etc) and in manufacturing.
• Alloy steels are carbon steels with other additives such as nickel, chromium,
vanadium, tungsten and manganese. These are stronger and tougher than
carbon steels and have a huge variety of applications including bridges, electricity
pylons, bicycle chains, cutting tools and rifle barrels.
• Stainless steel is very resistant to corrosion. It contains at least 10.5% chromi
um. Other metals such as nickel, molybdenum, titanium and copper are added
to enhance its strength and workability. It is used in architecture, bearings,
cutlery, surgical instruments and jewellery.
• Cast iron contains 3–5% carbon. It is used for pipes, valves and pumps. It is not
as tough as steel but it is cheaper.
• Magnets can be made of iron and its alloys and compounds.
• Iron catalysts are used in the Haber process for producing ammonia, and in the
Fischer–Tropsch process for converting syngas (hydrogen and carbon monoxide)
into liquid fuels.
• Iron plays an important role in the transfer of oxygen by hemoglobin. Each
hemoglobin binds four iron atoms. Iron in hemoglobin binds with oxygen as
it passes through the blood vessels in the lungs and releases it in the tissues.

1.5.7. Cobalt
Cobalt is a dark grey metal with a density of 8.90. It is insoluble in water at room
temperature.
1. Chemical reactions
a. Reaction of cobalt with air
Cobalt does not react readily with air. Upon heating the oxide Co₃O₄ is formed, and if
the reaction is carried out above 900°C, the result is cobalt (II) oxide, CoO.

Cobalt does not react directly with nitrogen, N₂.

b. Reaction of cobalt with water
Cold water has little effect upon cobalt metal. The reaction between red hot cobalt
metal and steam produces cobalt (II) oxide, CoO.

c. Reaction of cobalt with acids
Cobalt metal dissolves slowly in dilute sulphuric acid to form solutions containing
the hydrated Co(II) ion together with hydrogen gas, H2. The actual occurrence of Co
(II) in aqueous solution is as the complex ion [Co(OH₂)₆]²⁺.

It also dissolves in dilute nitric acid to form cobalt (II) nitrate and oxides of nitrogen.

(where NO_x stands for any oxide of nitrogen, i.e, NO, NO₂, …)

Concentrated nitric acid renders it passive due to the formation of oxide layer Co₃O₄
which is insoluble in the acid.
d. Reaction of cobalt with halogens
Metallic cobalt reacts with halogens, forming cobalt (II) halides.

2. Uses
• Cobalt, like iron, can be magnetized and so is used to make magnets. It is alloyed
with aluminium and nickel to make particularly powerful magnets.
• Other alloys of cobalt are used in jet turbines and gas turbine generators at
high temperature.
• Cobalt metal is sometimes used in electroplating because of its attractive appearance,
hardness and resistance to corrosion.
• Cobalt salts have been used for centuries to produce brilliant blue colours in
paint, porcelain, glass, pottery and enamels.
• Cobalt is an essential trace element and found at the centre of the vitamin
B₁₂ (cobalmin, C₆₃H₈₈CoN₁₄O₁₄P). It contains a cobalt(III) ion and is necessary for
the prevention of pernicious anaemia and the formation of red blood corpuscles,
but it is involved many other functions too.

1.5.8. Nickel
Nickel is a grey solid metal with density of about 8.9. It melts at 1455^oC and boils at
2920^oC.
1. Chemical reactions
a. Reaction of nickel with air
Nickel does not react with oxygen, O₂, at room temperature, under normal conditions.
Finely divided nickel can burn in oxygen, forming nickel (II) oxide, NiO.

b. Reaction of nickel with water
Nickel metal does not react with water under normal conditions. Nickel (II) ion
complexes with water under acidic and neutral conditions forming a light green
hexaqua nickel ion: [Ni(H₂O)₆]²⁺(_aq)

In basic condition, nickel hydroxide precipitates:

c. Reaction of nickel with acids
Nickel metal dissolves slowly in dilute sulphuric acid to form the aquated Ni(II) ion
and hydrogen, H₂.

The strongly oxidizing concentrated nitric acid, HNO₃, reacts on the surface of
nickel and passivates the surface.

d. Reaction of nickel with hydroxide ions
Metallic nickel does not react with aqueous sodium hydroxide.

e. eaction of nickel with halogens
Nickel reacts slowly with halogens, forming the corresponding dihalides.

2. Uses
• Nickel resists corrosion and is used to plate other metals to protect them. It is,
however, mainly used in making alloys such as stainless steel. Nichrome is an
alloy of nickel and chromium with small amounts of silicon, manganese and
iron. It resists corrosion, even when red hot, so is used in toasters and electric
ovens. A copper-nickel alloy is commonly used in desalination plants, which
convert seawater into fresh water. Nickel steel is used for armour plating. Other
alloys of nickel are used in boat propeller shafts and turbine blades.
• Nickel is used in batteries, including rechargeable nickel-cadmium batteries
and nickel-metal hydride batteries used in hybrid vehicles.
• Nickel has a long history of being used in coins. The US five-cent piece (known
as a ‘nickel’) is 25% nickel and 75% copper.
• Finely divided nickel is used as a catalyst for hydrogenating vegetable oils.
Adding nickel to glass gives it a green colour.

1.5.9. Copper
Copper is a light pink to red (shiny-reddish) metal of density 8.95 g/cm³. It melts at
1083^oC and boils at 2570^oC.

1. Chemical reactions
a. Reaction of copper with air
Copper metal is stable in air under normal conditions. When heated until red hot,
copper metal and oxygen react to form Cu₂O.

b. Reaction of copper with water
Copper does not react with water in all conditions.

c. Reaction of copper with acids
Copper is not dissolved by non-oxidizing dilute acids such as dilute
H₂SO₄and HCl to produce hydrogen gas. This is why it is called a
‘noble metal’. Other noble metals include gold, silver and platinum.
But copper metal dissolves in dilute and concentrated nitric acid, HNO₃ to form
copper (II) nitrate and oxides of nitrogen. Here nitric acid acts as an oxidising agent.

It also reacts with hot concentrated sulphuric acid to form copper (II) sulfate, sulphur
dioxide gas and water. But normally, sulphuric acid is not an oxidising acid!

d. Reaction of copper with halogens
Metallic copper metal reacts with the halogens forming corresponding dihalides.

2. Uses
• Historically, copper was the first metal to be worked by people. The discovery
that it could be hardened with a little tin to form the alloy bronze gave the
name to the Bronze Age.
• Traditionally it has been one of the metals used to make coins, along with silver
and gold. However, it is the most common of the three and therefore the
least valued. All US coins are now copper alloys, and gun metals also contain
copper.
• Most copper is used in electrical equipment such as wiring and motors. This is
because it conducts both heat and electricity very well, and can be drawn into
wires. It also has uses in construction (for example roofing and plumbing), and
industrial machinery (such as heat exchangers).
• Copper sulphate is used widely as an agricultural poison and as an algaecide
in water purification.
• Copper compounds, such as Fehling’s solution, are used in chemical tests for
sugar detection.
• Copper helps in storing iron, is involved in production of pigments for colouring
hair, skin and eyes.

1.5.10. Zinc
Zinc is a grey solid with a density of 7.14 g/cm³. It melts at 419.5 ^oC and boils at 907
^oC.

1. Chemical reactions
a. Reaction of zinc with air
Zinc reacts with oxygen in moist air. The metal burns in air with a blue-green flame to
form zinc (II) oxide, a material that goes from white to yellow on prolonged heating.
b. Reaction of zinc with water
Zinc is unaffected with cold water. However, elemental zinc will reduce steam at
high temperatures:

c. Reaction of zinc with acids
Zinc metal dissolves slowly in dilute sulphuric acid to form Zn(II) ions and hydrogen,
H₂. In aqueous solution the Zn (II) ion is present as the complex ion [Zn(H₂O)₆]²⁺.

When zinc reacts with oxidizing acids like HNO₃, no hydrogen gas is evolved.

2. Uses
• Mostly, zinc is used to galvanise other metals, such as ironsheets (amabati), to
prevent corrosion. Galvanised steel is used for car bodies, street lamp posts,
safety barriers and suspension bridges.Many houses in Rwanda are covered by
galvanized iron sheets (amabati).
• Large quantities of zinc are used to produce die-castings, which are important
in the automobile, electrical and hardware industries.
• Zinc is also used in alloys such as brass, nickel silver and aluminium solder.
• Zinc oxide is widely used in the manufacture of many products such as paints,
rubber, cosmetics, pharmaceuticals, plastics, inks, soaps, batteries, textiles
and electrical equipment. Zinc sulphide is used in making luminous paints,
fluorescent lights and x-ray screens.
• It is a component of insulin.

Checking up 1.5
1. State what is observed and write an equation, for the reaction that would
take place when
a. Copper is added to hot concentrated sulphuric acid.
b. Chromium is dissolved rapidly in hot concentrated aqueous alkali
c. Nickel (II) ions complexes react with water under acidic and neutral
conditions.
d. Powdered zinc is dissolved in hot aqueous alkali.
2. State at least one property that makes that:
a. An aluminum - scandium alloy be used in fighter planes, high-end bicycle
frames and baseball bats.
b. Many alloys of titanium with aluminium, molybdenum and iron be mainly
used in aircraft, spacecraft and missiles.
c. Vanadium-steel alloys be used for armour plate, axles, piston rods and
crankshafts.
d. Alternatives of tanning leather using chromium be investigated.
e. Manganese steel be used for railway tracks, safes, rifle barrels and prison
bars.
f. Iron be considered as an enigma.
g. Cobalt be necessary for the prevention of pernicious anaemia and the
formation of red blood corpuscles.
h. Nichrome be used in toasters and electric ovens.
i. Most copper be used in electrical equipment such as wiring and motors.
j. Galvanised steel be used for car bodies, street lamp posts, safety barriers
and suspension bridges.

Assignment
Question 3 is given to you as an assignment. You can use any source to carry
out research in order to gain and provide relevant information to be presented
comfortably.
3. The following figures show objects made in different transition metals.
Observe them and complete the table with the main transition metal
which forms the objects, its two important properties and other two uses
(apart from that shown by the figure).

1.6. Identification of transition metal ions
Activity 1.6
Given a substance Y which contains one cation (from transition metal) and one
anion,identify the cation and anion in Y. Carry out the following tests on Y , record
your observations and deductions in the table below. Identify any gas evolves.

• The cation in Y is …………
• The anion in Y is ……………
• Write the ionic equations for the reactions in test (i) and test (ii)
……………………
Different transition metals have different colors. Also, different charges, or cations
of one transition metal can give different colors. Another factor is the natural of
the ligand. The same cations of a transition metal can produce a different color
depending on the ligand it binds to.
Many compounds containing transition metals have certain characteristic colours
and thus, by observing a compound, we can not identify it.
• Appearance or colour of different solid compounds containing transition
metals
Table 1.13: Colours of different solid compounds containing transition metals
(first series)

• Colours of aqueous solutions of some transition metal ions
In aqueous solutions where water molecules are the ligands, the colours of some
metal ions observed are listed in the table below:
Table 1.14: Colours of different aqueous solutions containing some transition
metals (first series)
• Action of heat on solid compounds containing transition metal ions
Table 1.15: Colours of different solid compounds containing transition metals
(first series) due to action of heat.
Note: On heating the following temporary colour changes may also occur:

• Effect of aqueous sodium hydroxide and aqueous ammonia on solutions
containing transition metal ions
The hydroxides of transition metals are precipitated from solutions of the metal ions
by the addition of hydroxide ions or ammonia. The colour of the precipitate can
often be used to identify the metal present. The precipitates formed are gelatinous
and often coloured and some form soluble complex ions with excess ammonia.
a. To about 1cm³ of the solution containing the positive ion (cation), add
2M aqueous sodium hydroxide dropwise until in excess

b. To about 1cm³of the solution containing the positive ion (cation), add
2M aqueous ammonia dropwise until in excess

Confirmatory tests for some transition metal ions
Confirmatory tests are the tests required to confirm the analysis. Generally, a
confirmatory test is used after other tests have been carried out to isolate/identify
the ion. In order to confirm the ion without any dought

a. Zinc ions
i. Addition of little solid ammonium chloride followed by disodium
hydrogen phosphate solution to a solution of zinc ions gives a white
precipitate. The precipitate dissolves in excess ammonia or dilute
mineral acids.
ii. Addition of potassium ferrocyanide solution to a solution of zinc ions
gives a white precipitate.

b. Chromium ions
To a solution of chromium (III) ions, add excess aqueous sodium hydroxide followed
by little hydrogen peroxide and boil the resultant mixture. A yellow solution of a
chromate is formed.

Treatment of the yellow solution with:
i. Lead (II) ethanoate or Lead(II)nitrate solution gives a yellow precipitate of
Lead(II) chromate. Pb²⁺(aq) + CrO₄^2-(aq) →PbCrO4(s)
ii. Barium nitrate (or chloride) solution gives a yellow precipitate of barium
chromate.

c. Manganese (II) ions
To the solution of manganese (II) ions, add little concentrated nitric acid followed
by little solid lead(IV) oxide or solid sodium bismuthate(V) and boil the mixture. A
purple solution is formed due to MnO₄^- ion.

d. Iron (II) ions
i. Addition of potassium hexacyanoferrate (III) solution to a solution of iron
(II) ions gives a dark blue precipitate .
ii. Addition of few drops of concentrated nitric acid to a solution of iron (II)
ions gives a yellowish solution due to iron (III) ions formed. The solution
gives positive test for iron (III) ions.
e. Iron(III) ions
i. Addition of potassium hexacyanoferrate (II) solution to a solution of
iron(III) ions gives a dark blue precipitate
ii. Addition of potassium thiocyanate or ammonium thiocyanate
solution to a solution of iron (III) ions gives a blood red coloration.
f. Cobalt (II) ions
Addition of potassium thiocyanate or ammonium thiocyanate solution to a solution
of cobalt(II) ions gives a blue colored product of potassium cobalt(II) tetrathiocyanate.

g. Nickel (II) ions
i. Addition of potassium cyanide solution gives a yellow-green
precipitate of Nickel(II) cyanide. The precipitate dissolves in excess
reagent to form a dark yellow solution tetracyanonickel (II) ion.

ii. Addition of aqueous ammonia followed by 2 to 3 drops of
dimethylglyoxime solution to a solution of nickel (II) ions gives a red
precipitate. The formation of this precipitate may sometimes require
that the solution mixture would be warmed.
h. Copper (II) ions
In addition to use of aqueous ammonia, the copper(II) ions can be confirmed by
addition of the following reagents to an aqueous solution of copper(II) ions:
i. Potassium iodide solution: A white precipitate of copper (I) iodide
stained brown with free iodine.

Brown color fades on addition of sodium thiosulphate solution due to the reaction
below:
Checking up 1.6 (a)
Given a substance K which contains one cation and one anion, carry out the
following tests on K and record your observations and deductions in the table
below. Identify any gas evolved.

• The cation present in the compound K is ……………
• The anion present in the compound K is ……………
Checking up 1.6 (b)
You are provided with substance D which contains one cation and one anion.
You are required to identify the cation and anion in D. Carry out the following
tests, record your observations and deductions in the table below. Identify any
gas evolved.

Checking up 1.6 (c)
Aqueous sodium hydroxide is added separately to solutions of salts of the
transition metals A, B and C. Identify A, B and C from the following observations.
A: The white precipitate which appears is soluble in an excess of aqueous sodium
hydroxide and also in aqueous ammonia.
B: The blue precipitate which appears is insoluble in an excess of aqueous sodium
hydroxide but dissolves in aqueous ammonia to form a deep blue solution.
C: The green precipitate which appears is insoluble in an excess of aqueous
sodium hydroxide and also in aqueous ammonia.
END UNIT ASSESSMENT
a. Multiple choice questions: Write the Roman number corresponding to the
correct answer.
1. Which of the following elements is not a transition metal?
i. Copper
ii. Nickel
iii. Iron
iv. Magnesium
2. Which of the following complexes is linear?
i. [Ag(NH₃)_2]+
ii. [CoCl₄]^2-
iii. [Pt(NH3)₂Cl²]
iv. [CuCl₄]^2-
3. Which of the following ions does not form coloured solutions?
i. Cu⁺
ii. Mn²⁺
iii. Cr³⁺
iv. Co²⁺
4. Which of the following reactions of Cu²⁺ is an example of a chelation
reaction?
i. [Cu(H₂O)₆]²⁺ + 2OH- → [Cu(H₂O)₄(OH)₂] + 2H₂O

2. What is the characteristic of electron configurations of transition metals?
3. Which electrons, 3d or 4s, have the lowest ionization energies in a
transition metal?
4. a. Name any three transition metals that are essential to the biological
system.
b. Why do you think transition metals form coordination compounds that
have covalent bonds?
5. Name the following coordination compounds using systematic
nomenclature.
a. [Co(H₂O)₆]Cl₂
b. [Cr(NH₃)₆](NO₃)₃
c. K₄[Fe(CN)₆]
d. Na[Au(CN)₄]
e. [Co(H₂O)₂(en)₂]Cl₃
6. a. (i) What is meant by the term co-ordinate bond?
(ii) Explain why co-ordinate bonds can be formed between transition
metal ions and water molecules.
b. What name is given to any ligand that can form two co-ordinate bonds
to one metal ion? Give an example of such a ligand.
7. In order to determine the concentration of a solution of cobalt(II) chloride,
a 25.0 cm³ sample was titrated with a 0.0168 M solution of EDTA^4-; 36.2
cm3 were required to reach the end-point. The reaction occurring in the
titration is:
[Co(H₂O)₆]²⁺ + EDTA⁴–——→ [Co(EDTA)]²– + 6 H₂O
a. What type of ligand is EDTA4–?
b. Calculate the molar concentration of the cobalt (II) chloride solution.
8. The ethanedioate (oxalate) ion,C₂O₄² , acts as a bidentate ligand. This ligand
forms an octahedral complex with iron (III) ions.
a. Deduce the formula of this complex and draw its structure showing all
the coordinate bonds present.
b. Give the name of a naturally-occurring in human body complex
compound which contains iron.
c. What is theimportant function of this complex compound?
9. The compound [Co(NH₃)₄Cl₂]Cl contains both chloride ions and ammonia
molecules as ligands.
a. State why chloride ions and ammonia molecules can behave as
ligands.
b. What is the oxidation state and the co-ordination number of cobalt in
this complex compound?
10. a. Suggest why the compound [Co(NH₃)₆]Cl₃ has a different colour from
that of [Co(NH₃)₄Cl₂]Cl.
b. Name and give the formula of an ammonia complex used to distinguish
between aldehydes and ketones.
11. Chloride ions form the tetrahedral complex ion [AlCl4]– but fluoride
ions form the octahedral complex ion [AlF6]3-. Suggest a reason for this
difference.
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  UNIT 12. ELECTROCHEMICAL CELL AND APPLICATIONS
  Key unit competency:
  To be able to explain the working and industrial applications of electrochemical and electrolytic cells
  Learning objectives
  At the end of this unit, students will be able to:
  • Define the term electrochemical cell.
  • Construct a simple galvanic cell.
  • Describe the standard hydrogen electrode.
  • Explain the working of galvanic cells using the fully labelled diagram.
  • Use the e.m.f. of the galvanic cell to predict if the cell will generate current or not.
  • Record the results of a measurement accurately using a voltmeter.
  • Calculate standard cell potentials from standard electrode potentials of two half cells.
  • Use standard electrode potentials of cells to determine the direction of electron flow and feasibility of a reaction.
  • Predict how the value of an electrode potential varies with concentration using Le Chatelier Principle.
  • Apply the principles of redox processes to energy storage devices.
  • Compare electrochemical cell with electrolytic cell.
  • Properly use electrolytic cell to carry out electroplating of graphite by copper.
  • Describe industrial applications of electrochemical cells.
  • Appreciate contributions of electrochemistry to the social and economic development of the society.
  • Develop a culture of team work, sense of responsibility in group activities and experiments.
  Introductory Activity
  Activity 1
  Set up the circuit shown in the following figure using 1.25 V, 0.25 A lamp bulb, and a beaker two-thirds full of a 1 molar solution of sulphuric acid.
  
  Into the beaker place strips of magnesium ribbon and copper foil, thoroughly
  cleaned by abrasive paper, and connect them up to the circuit using the crocodile
  clips on the ends of the connecting wires.
  1. What do you observe?
  2. From your observations, suggest the importance of this apparatus.
  3. Describe what is happening to allow these observations.
  Activity 2
  Do research in the library, textbooks, and if you have access to Internet try to access the video on the following link:
  Wk which shows how an electrochemical cell work to help you answer the discussion questions.
  
  Discuss each question with your partner(s) and try to write down your best answers
  while watching this video.
  Discussion questions:
  1. The half-cell at anode consist of _______________
  2. The half-cell at cathode consist of _______________
  3. What is the salt bridge used in the experiment?
  4. What is the function of salt bridge?
  5. Which direction does the electron flow?
  6. Why Zn is likely to lose electron?
  7. Which metal undergoes oxidation?
  8. Which metal undergoes reduction?
  9. What can you observe during the oxidation process?
  10. Which metal undergoes reduction?
  11. What can you observe during the reduction process?
  12. How are the excessive SO₄^2- ions neutralised?
  13. How are the excessive Zn²⁺ ions neutralised?
  14. Identify the important characteristics of a cell.
  15. Suggest any use of this electrochemical cell.
  12.1. Definition of electrochemical cell
  Activity 12.1
  Basing on the observations and the answers provided in the introductory activity above, suggest a definition of an electrochemical cell.
  An electrochemical cell is a device which is capable of either producing electrical energy from chemical reactions or causes chemical reactions to take place through the introduction of electrical energy.
  There are two types of electrochemical cells: galvanic (voltaic) cells and electrolytic cells.
  A galvanic (voltaic) cell is a device used to convert chemical energy of a redox reaction into electrical energy.
  Electrolytic cell is a type of chemical cell in which the flow of electric energy from an external source causes a redox reaction to occur.
  A galvanic cell is named after Luigi Galvani, an Italian physicist (1780). It is also called Voltaic cell, after an Italian physicist, Alessandro Volta (1800).
  Both L. Galvani and A. Volta contributed greatly in the existence of this type of electrochemical cells.
  Both types of cells contain electrodes where the oxidation and reduction reactions occur. Oxidation occurs at the electrode called the anode and reduction occurs at the electrode called the cathode.
  12.2. Description of electrochemical cells
  Activity 12.2 (a)
  A Daniell cell is an electrochemical cell consisting of a zinc electrode in zinc sulphate solution and a copper electrode in copper sulphate solution, linked by a salt bridge. Use the method below to make one.
  Equipment and materials
  • A piece of zinc rod or plate
  • Two (2) beakers
  • A piece of copper rod or plate
  • One (1) voltmeter
  • 1 mol dm^-3 copper sulphate solution
  • Connecting wires and crocodile clips (or similar)
  • 1 mol dm^-3 zinc sulphate solution
  • A plastic tube (or a U-tube), full of 2 M potassium nitrate (or 2 M sodium chloride)
  solution stoppered on both sides by a cotton wool soaked in that potassium
  nitrate (or sodium chloride) solution, acting as a salt bridge.
  Procedure
  Set up the apparatus as shown below.
  
  Connect the voltmeter leads to the two electrodes. Read the voltmeter.
  With the voltmeter connected correctly (with the voltage positive), remove the salt bridge. Note what happens
  Discussion questions
  1. The cell may be described as being made up of two half cells. Each has a metal (the electrode) in a solution of its ions. In this cell, zinc is the negative electrode and copper is the positive electrode. Describe, using ionic equations, what happens in the cell.
  2. Draw the above electrochemical cell. Add the labels ‘negative terminal’ and ‘positive terminal’. Show the direction of:
  a. The flow of electrons
  b. The (conventional) current.
  Oxidation-reduction (or redox) reactions take place in electrochemical cells. Spontaneous reactions occur in galvanic (voltaic) cells; non-spontaneous reactions occur in electrolytic cells. Both types of cells contain electrodes where the oxidation and reduction reactions occur. Oxidation occurs at the anode and reduction occurs at the cathode.
  The anode of an electrolytic cell is positive, since it attracts anions from the solution, whereas the cathode is negative and attracts positive ions. In a galvanic cell, the anode is negatively charged, since the spontaneous oxidation at the anode is the source of the cell’s electrons or negative charge. The cathode of a galvanic cell is its positive terminal.
  In both galvanic and electrolytic cells, oxidation takes place at the anode and electrons
  flow from the anode to the cathode.
  In this unit, we will deal with galvanic cells especially because electrolytic cells were dealt with in Senior 5. As said above, the redox reaction in a galvanic cell is a spontaneous reaction. For this reason, galvanic cells are commonly used as batteries or sources of energy. Galvanic cell reactions supply energy which is used to perform work. Galvanic cells are electrochemical cells which use the transfer of electrons in redox reactions to supply an electric current.
  The energy is harnessed by putting the oxidation and reduction reactions in separate containers, connected by a salt bridge, or separated by a porous membrane (an apparatus that allows electrons to flow).
  A galvanic cell generally consists of two different metal rods called electrodes. Each electrode is immersed in a solution containing its own ions (in the separate containers) and these form a half cell. The solutions in which the electrodes are immersed are called electrolytes i.e. made of ions.
  One electrode acts as anode in which oxidation takes place and the other acts as the cathode in which reduction takes place.
  12.2.1 Half cells and Redox reactions in half cells
  Each chamber of an electrochemical cell constitutes a half-cell containing an electrode and an electrolyte. Half-cells are sometimes known as redox electrodes or redox couples.
  3. Using your observations, explain the purpose of the salt bridge and use any relevant document to explain how it works.
  Oxidation-reduction (or redox) reactions take place in electrochemical cells.
  Spontaneous reactions occur in galvanic (voltaic) cells; non-spontaneous reactions occur in electrolytic cells. Both types of cells contain electrodes where the oxidation and reduction reactions occur. Oxidation occurs at the anode and reduction occurs at the cathode.
  The anode of an electrolytic cell is positive, since it attracts anions from the solution, whereas the cathode is negative and attracts positive ions. In a galvanic cell, the anode is negatively charged, since the spontaneous oxidation at the anode is the source of the cell’s electrons or negative charge. The cathode of a galvanic cell is its positive terminal.
  In both galvanic and electrolytic cells, oxidation takes place at the anode and electrons flow from the anode to the cathode.
  In this unit, we will deal with galvanic cells especially because electrolytic cells were dealt with in Senior 5. As said above, the redox reaction in a galvanic cell is a spontaneous reaction. For this reason, galvanic cells are commonly used as batteries or sources of energy. Galvanic cell reactions supply energy which is used to perform work. Galvanic cells are electrochemical cells which use the transfer of electrons in redox reactions to supply an electric current.
  The energy is harnessed by putting the oxidation and reduction reactions in separate containers, connected by a salt bridge, or separated by a porous membrane (an apparatus that allows electrons to flow).
  A galvanic cell generally consists of two different metal rods called electrodes.
  Each electrode is immersed in a solution containing its own ions (in the separate containers) and these form a half cell. The solutions in which the electrodes are immersed are called electrolytes i.e. made of ions.
  One electrode acts as anode in which oxidation takes place and the other acts as the cathode in which reduction takes place.
  12.2.1 Half cells and Redox reactions in half cells
  Each chamber of an electrochemical cell constitutes a half-cell containing an electrode and an electrolyte. Half-cells are sometimes known as redox electrodes or redox couples.
  Types of half cells
  The half cells may be categorized into three types: Metal, non-metal and ion half cells.
  1. Metal half cells
  This half cell is made of metal and its aqueous ions. It consists of a metal (electrode)
  dipped into an aqueous solution containing its own ions. For example: Zn/Zn2+ half cell.
  2. Non-metal half cells
  This half cell is made from a non-metal and its aqueous ions. For example, hydrogen half cell comprises hydrogen gas in contact with hydrogen ions: H2/2H+
  3. Ion half cell
  This type of half cell consists of an inert electrode such as platinum electrode dipping into a solution containing ions of the same metal in two different oxidation states.
  For example, a half cell containing Fe³⁺(aq) and Fe²⁺(aq) ions: Fe³⁺(aq), Fe²⁺(aq)/Pt
  Working of a half cell
  The simplest half-cell consists of a metal placed in solution of its ions.
  If a piece of metal is put into a solution of its own ions, atoms from the metallic structure pass into a solution and form hydrated metal ions. At the same time hydrated, metal ions gain electrons from the metal structure to form metal atoms.
  Eventually, an equilibrium is established.
  For example, when a strip of zinc metal is dipped in aqueous solution of zinc
  sulphate, some zinc atoms are oxidized. Each zinc atom that is oxidized leaves
  behind 2 electrons and enters the solution as Zn²⁺ ion.
  Oxidation: Zn(S) → Zn²⁺(aq) + 2e
  At the same time, some Zn²⁺ ions in the solution gain 2 electrons from the zinc strip and deposit as zinc atoms. They are reduced.
  Reduction: Zn²⁺(aq) + 2e → Zn(s)
  The opposing oxidation and reduction processes quickly come to an equilibrium:
  A strip of metal used in an electrochemical experiment is called an Electrode. By convention, the equilibrium is written with the electrons on the left-hand side. The electrode potential of the half-cell indicates its tendency to lose or gain electrons in the equilibrium.
  12.2.2 Constructing cells from half cells
  A simple electrochemical cell can be made by connecting together two half cells with different electrode potentials.
  
  Figure 12.1: A general cell constructed from its half-cells
  • One half cell releases electrons (oxidation at the anode).
  • The other half cell gains electrons (reduction at the cathode).
  Note that, in the electrochemical cell, the electrons flow from the negative terminal
  (anode) to the positive terminal (cathode) and the current flows in the opposite
  direction i.e. from cathode to anode.
  The salt bridge is usually an inverted U-tube filled with a concentrated solution
  of an inert electrolyte. The inert electrolyte is neither involved in any chemical
  change, nor does it react with the solutions in the two half cells. Generally salts like
  KCl, KNO₃, Na₂SO₄ and NH₄NO₃ are used as the electrolytes.
  To prepare salt bridge, agar-agar or gelatin is mixed with a hot concentrated solution of electrolyte and is filled in the U-tube. On cooling, the solution sets in the form of a gel inside the U-tube and thus prevents the inter mixing of the fluids.
  The two ends of the U-tube are then plugged or stoppered with cotton wool to minimise diffusion.
  a. The Significance of Salt Bridge
  Its main function is to prevent the potential difference that arises between the two solutions when they are in contact with each other. This potential difference is called the liquid junction potential.
  • It completes the electrical circuit by connecting the electrolytes in the two half cells.
  • It prevents the diffusion of solutions from one half-cell to the other.
  • It maintains the electrical neutrality of the solutions in the two half cells.
  b. How is the electrical neutrality of the solutions in the two half cells maintained using a salt bridge?
  In the anodic half-cell, there will be accumulation of positive charge when the positive ions that are formed pass into the solution. To maintain the electrical neutrality, salt bridge provides negative ions.
  For example, in Daniell cell, zinc oxidizes at the anode and passes into the solutions as Zn²⁺ ions, so there will be accumulation of positive charge in the solution. To maintain the electrical neutrality of the solution, the salt bridge provides negative ions (may be, NO₃^- or Cl^-).
  In the cathodic half-cell, there will be accumulation of negative ions formed due to the reduction of positive ions. To maintain the electrical neutrality, salt bridge provides positive ions.
  For example, in Daniell cell, Cu²⁺ ions from the CuSO₄ solution is reduced by the electrons formed by the oxidation of zinc, and deposited on the copper cathode.
  As a result, the concentration of Cu²⁺ ions decreases in the solution and that of SO₄ ^2- ions (sulphate ions) increases. So there will be an accumulation of negatively charged sulphate ions around the cathode. To maintain the electrical neutrality, salt bridge provides positive ions (may be, K⁺ or NH₄ ⁺).
  12.2.3. Daniell Cell
  The Daniell cell was invented by a British chemist, John Frederic Daniell. In the
  Daniell cell, copper and zinc electrodes are immersed in a solution of copper (II)
  sulphate (CuSO_4(aq)) and zinc (II) sulfate (ZnSO₄ (aq)) respectively. The two half cells
  are connected through a salt bridge. Here, zinc acts as anode and copper acts as cathode.
  At the anode, zinc undergoes oxidation to form zinc ions and electrons. The zinc ions pass into the solution. If the two electrodes are connected using an external wire, the electrons produced by the oxidation of zinc travel through the wire and enter into the copper cathode, where they reduce the copper ions present in the solution and form copper atoms that are deposited on the cathode.
  Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
  The half-reactions are:
  • At Anode: Anode: is an electrode at which oxidation occurs. Because it is the
  source of electrons, in voltaic cell, it is also called the negative electrode.
  Zn(s)→ Zn²⁺(aq) + 2e-
  • At Cathode: Cathode: is an electrode at which reduction occurs. Because it is the receiver of electrons, in voltaic cell, it is also called the positive electrode.
  Cu²⁺(aq) + 2e^- → Cu(s)
  Total cell reaction is the sum of the two half-cell reactions:
  Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
  • Zinc is oxidized to Zn²⁺ and releases 2 electrons; it is the anode of the cell.
  • Cu²⁺ is reduced to copper and captures two electrons; it is the cathode where copper metal is deposited.
  • In half-Zn cell, Zn²⁺ ions are produced, which means that the solution is positively charged.
  • In half-cell of Cu, Cu²⁺ ions are deposited and there is an excess of SO₄^2-ions, so that the solution becomes negatively charged.
  The flow of electrons takes place through the external circuit. As electrons leave one half of a galvanic cell and flow to the other, a difference in charge is established. If no salt bridge were used, this charge difference would prevent further flow of electrons.
  The salt bridge and wire join the two half cells:
  • The salt bridge connects the two solutions, allowing ions to be transferred between the half cells.
  • The wire connects the two metals, allowing the electrons to be transferred between the two half cells.
  The two equilibria which are set up in the half cells are:
  The negative sign of the zinc E⁰ value shows that it releases electrons more readily than hydrogen does while the positive sign of the copper E⁰ shows that it releases electrons less readily than hydrogen.
  Taking the apparatus as a whole, there is a chemical reaction going on, in which zinc is going into solution as zinc ions, and is giving electrons to copper (II) ions to turn them into metallic copper. This is exactly the same reaction that occurs when you drop a piece of copper into some copper (II) sulphate solution. The blue colour of the solution fades as the copper (II) ions are converted into brown copper metal. The final solution contains zinc sulphate (where the sulphate ions are spectator ions). You can add the two electron-half-equations above to give the overall ionic equation for the reaction.
  Zn(s) → Zn²⁺(aq) + 2e^- Oxidation to anode electrode
  Cu²⁺(aq) + 2e^- → Cu(s) Reduction to cathode electrode
  Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) Overall reaction
  12.2.4. Description of standard hydrogen electrode as used to determine standard electrode potentials
  Activity 12.2 (b)
  Some electrode potentials experimentally measured are found in the table 12.3. Use
  any document to:
  1. Suggest a reason why these potentials are called “standard”.
  2. Describe the way the electrode potential given for an atom can be measured.
  When a metal electrode is dipped in a solution containing its metal ions, a potential difference is developed at the metal /solution interface. This potential difference is called the electrode potential.
  For example, when a copper rod is dipped in a solution containing Cu²⁺ ions, the Cu²⁺ ions gain electrons from the copper rod leaving positive charge on the copper rod. As a result, a potential difference is set up between the copper rod and the solution and is called the electrode potential of copper.
  In a galvanic cell, the anode has a negative potential and cathode has a positive
  The negative sign of the zinc E⁰ value shows that it releases electrons more readily than hydrogen does while the positive sign of the copper E⁰ shows that it releases electrons less readily than hydrogen.
  Taking the apparatus as a whole, there is a chemical reaction going on, in which zinc is going into solution as zinc ions, and is giving electrons to copper (II) ions to turn them into metallic copper. This is exactly the same reaction that occurs when you drop a piece of copper into some copper (II) sulphate solution. The blue colour of the solution fades as the copper (II) ions are converted into brown copper metal. The final solution contains zinc sulphate (where the sulphate ions are spectator ions). You can add the two electron-half-equations above to give the overall ionic equation for the reaction.
  Zn(s) → Zn²⁺(aq) + 2e^- Oxidation to anode electrode
  Cu²⁺(aq) + 2e^- → Cu(s) Reduction to cathode electrode
  Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) Overall reaction
  12.2.4. Description of standard hydrogen electrode as used to determine standard electrode potentials
  Activity 12.2 (b)
  Some electrode potentials experimentally measured are found in the table 12.3. Use
  any document to:
  1. Suggest a reason why these potentials are called “standard”.
  2. Describe the way the electrode potential given for an atom can be measured.
  When a metal electrode is dipped in a solution containing its metal ions, a potential difference is developed at the metal /solution interface. This potential difference is called the electrode potential.
  For example, when a copper rod is dipped in a solution containing Cu²⁺ ions, the Cu²⁺ ions gain electrons from the copper rod leaving positive charge on the copper rod. As a result, a potential difference is set up between the copper rod and the solution and is called the electrode potential of copper.
  In a galvanic cell, the anode has a negative potential and cathode has a positive potential. The potential of each individual half-cell cannot be measured. We can
  measure only the difference between the potential of the two half cells. While it is
  impossible to determine the electrical potential of a single electrode, we can assign an electrode the value of zero and then use it as a reference.
  The electrode chosen as having the value of “zero” is called the Standard Hydrogen
  Electrode (SHE). The SHE consists of 1 atm of hydrogen gas bubbled through a 1 M
  strong acid solution, usually at room temperature. Platinum or graphite, which is
  chemically inert, is used as the electrode.
  Thus, electrode potential at standard conditions (temperature of 25°C, pressure
  of 1 atm and concentration of 1 M concentration for the electrolyte) is called
  the “standard electrode potential”. It is denoted by the symbol E⁰ where the
  superscript “⁰” on the E denotes standard conditions.
  Other standard electrode potentials can be determined using the SHE. The
  standard reduction (electrode) potential can be determined by subtracting
  the standard reduction potential for the reaction occurring at the anode from
  the standard reduction potential for the reaction occurring at the cathode.
  Here, the minus sign is necessary because oxidation is the reverse of reduction.
  The diagram below shows how the standard potential, E⁰ of copper can be determined.
  The copper electrode contains Cu²⁺ ions in equilibrium with copper metal. The hydrogen electrode is linked via a salt bridge to the solution in which the copper electrode is immersed. This permits charge transfer and potential measurement but not mass transfer of the acid solution in the electrode.
  The oxidation and reduction at the electrodes leads to a standard electrode potential between the electrode and ions in solution. The standard electrode potential is expressed in Volts (V).
  Volt is the measure of electromotive force (e.m.f.) produced by a cell. Standard electrode potentials refer to redox potentials as well. A voltmeter measures a difference in electric potential between two points in an electrical circuit.
  If the two points are the electrodes in voltaic cells, the potential difference is the driving force that propels electrons from the anode to the cathode. It is also called the cell potential (E_cell). Since the measurements are in volts, the cell potential is also called the “Cell voltage”.
  When E⁰ are measured relative to the SHE (or some other reference electrode), a voltmeter is used. In order to resist any current flowing between the electrode and the SHE, the voltmeter must have high impedance and if a current were allowed to flow, the electrodes would become polarised and would no longer be at equilibrium.
  The hydrogen electrode is not a convenient reference electrode to use in measurements because maintaining a stream of hydrogen at 1 atm takes careful management. It is usual to employ a secondary standard such as the saturated calomel electrode.
  This electrode contains mercury and solid mercury (I) chloride (calomel) in contact with a saturated with a solution of potassium chloride. An equilibrium is set up between Hg⁺ ions and Hg(l). The emf of a cell composed of a saturated calomel electrode and a standard hydrogen electrode is 0.244 V at 25 ^oC. The saturated calomel electrode is easier to use than the standard hydrogen electrode. It can be
  combined with electrodes of unknown potential, and from the emf of the cell, the unknown electrode potential can be found.
  Table 12.1: Standard Reduction Potentials in aqueous solution (at 25 ^oC).
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  UNIT 13:FACTORS THAT AFFECT THE RATE OF REACTIONS
  Key unit competency:
  Explain the factors that affect the rate of chemical reaction and use Arrhenius
  equation to calculate the ratio of rate constant and activation energy with change
  in the temperature.
  Learning objectives
  By the end of this unit the learners should be able to:
  • Explain the concept of reaction kinetics
  • Explain the effect of different conditions on the rate of reaction
  • Carry out experiments to show how different factors affect the rate of chemical
  reactions
  • Predict the effect of changing conditions on the rate of reactions
  • Appreciate the importance of reaction kinetics.
  • Appreciate the importance of different conditions on the reaction rates.
  Introductory Activity
  Observe the following processes and answer the questions below.
  
  A: RUSTING PROCESS (Oxidation of iron)
  B: PRECIPITATION REACTIONS
  
  C: FERMENTATION
  
  D: COMBUSTION
  
  1. Classify the above processes or reactions as slow or fast in the following table.
  
  2. Which conditions are necessary for rusting process (corrosion) to take place?
  3. During wood combustion, describe what will happen if the amount of air increases.
  4. What is the role of yeast in the fermentation process?
  
  Chemistry, by its nature, is concerned with change. It is the study of matter, its
  composition, properties, and reactivity. Substances having well-defined properties
  are converted by chemical reactions into other substances with different properties.
  One of the features of a chemical reaction is the rate or speed of the reaction and
  the factors controlling the rate of a chemical reaction. For example, which factors
  determine how rapidly food gets spoiled? Or what controls the rate at which
  fuel burns in an auto engine? All these questions can be answered by the branch
  of chemistry, which deals with the study of reaction rates and their mechanisms,
  called chemical kinetics.
  
  13.1. Concept of reaction kinetics
  Activity 13.1
  a. Burn a small amount (about 2 cm3) of ethanol and record the time taken for
  this amount of ethanol to be completely burned.
  a. Put 4ml of lead (II) nitrate aqueous solution in a test tube and add three
  drops of aqueous potassium iodide solution, then record the time taken
  for the yellow precipitate to be formed.
  a. Put two or 3 iron nails in an open beaker containing water and predict
  what will happen to these nails after three weeks or one month.
  
  Questions:
  2. Suggest the name of a reaction that takes place in each of the above activities.
  3. In the above three activities, indicate the slowest and fastest reaction.
  4. Using books or internet describe your understanding on the rate of chemical reaction.
  
  Chemical kinetics is a branch of chemistry mainly dealing with the rates of reaction.
  The goal of chemical kinetics is to investigate how fast the reactants are consumed
  or the products are formed. It is important to predict the rate of processes and to
  find the influencing factors that promote a desired reaction or inhibit an undesired one.
  
  During a chemical reaction, reactants are converted to products by breaking old
  bonds and formation of new bonds. Some chemical reactions may be slow or fast.
  For example, the reaction between a solution of lead (II) nitrate and potassium iodide is a fast reaction.
  
  The precipitation of silver chloride is also fast. It occurs instantaneously by mixing
  aqueous solutions of silver nitrate and sodium chloride.
  However, the rusting of iron is a slow process that takes place over many days, even
  over the years. Similarly, the formation of diamond and other minerals in the earth’s crust takes millions of years.
  
  Chemical kinetics is the study of rates and mechanisms of chemical reactions. Any
  chemical process may be broken down into a sequence of one or more single-step
  processes known either as elementary processes, elementary reactions, or elementary steps.
  
  The hydrolysis of sucrose was the first chemical reaction whose rate was quantitatively
  studied by L. Wilhemly in 1850. Since that period, advanced research has
  been made in the field of chemical kinetics, enabling chemists to study very fast
  reactions (reactions that occur within less than a second). When we talk about the
  rate of a chemical reaction, we mean the rate at which reactants are used up, or
  equivalently the rate at which products are formed.
  
  The rate of a reaction can be defined as the change in concentration of a reactant
  or product in a given period of time. The reaction rate tells us how fast the reaction
  is taking place by indicating how much of a reactant is consumed or how much of a product is yielded in a given time.
  
  For example, if we consider the speed of an automobile, it can be expressed in terms
  of change in the distance covered in a certain period of time.
  In chemical kinetics, the rate of a reaction can be expressed as the initial rate, average
  rate or instantaneous rate.
  
  The initial rate is the rate of a reaction measured at the beginning of the reaction. It
  is the change in concentration of a reactant, or product, per unit time at the start of
  reaction. The initial rate is simply the instantaneous rate at time equal to zero, or at
  the beginning of the reaction.
  
  The average rate is defined as the change in concentration of a reactant or product
  of a chemical reaction in a given interval of time.
  
  For example, consider the following reaction:
  
  In this reaction, if one mole of the reactant R produces one mole of the product P. If
  [R]₁ and [P]₁ are the concentrations of R and P respectively at time t₁ and [R]₂, [P]₂ are
  
  the concentrations of R and P respectively at time t₂ then,
  
  By convention, the rate is always expressed as positive quantity. Because the
  concentration of reactants is decreasing with time, Δ[R] is a negative number. We
  use the negative sign to make the rate of reaction a positive quantity.
  
  Average rate depends upon the change in concentration of reactants or products
  and the time taken for that change to occur.
  For a better understanding, consider the following examples:
  
  Consider the following reaction at 300 oC: 2 NO2(g) →2 NO(g) + O2(g)
  If in the first 150 seconds, the concentration of nitrogen dioxide has decreased from
  0.0100 mol/L to 0.0055 mol/L, the average rate for the disappearance of nitrogen
  dioxide for the first 150 seconds is:
  
  Suppose that in the next 150 seconds the concentration of nitrogen dioxide
  decreases from 0.0055 mol/L to 0.0038 mol/L. The average rate becomes:
  
  Consider the reaction of the acidified hydrogen peroxide (H2O2) added to a solution of potassium iodide (KI) and iodine is liberated.
  
  In this reaction, initially the concentration of iodine is zero. But with the passage of
  time, it increases and the reaction solution becomes brownish. The concentration
  of iodine can be measured at different intervals of time by titration against sodium
  thiosulphate. If the concentration of iodine rises from 0 to 10-5 mol L-1 in 10 seconds,
  the average rate becomes:
  
  The instantaneous rate is the change in concentration of a reactant or product of a chemical reaction at a given instant. This rate is measured when the reaction is in progress at any point of time.
  
  It is obtained when we consider the average rate at the smallest time interval dt (i.e. when Δt approaches zero).
  
  Mathematically, for an infinitesimally small dt, the instantaneous rate is given by:
  
  Where d[R] = small change in concentration of R and dt = small change in time
  For the general reaction:
  
  Where a, b, c and d are stoichiometric coefficients of reactants and products.
  Example:
  The decomposition of dinitrogen pentoxide: 2 N₂O₅ (g) → 4 NO₂(g) + O₂(g) was
  performed in the laboratory and the rate of formation of NO₂ was found to be 0.53 M/s.
  
  a. What was the rate of formation of O₂(g)?
  b. What was the rate of consumption N₂O₅(g)?
  
  Answer:
  a. First determine the relationship between NO₂(g) and O₂(g) in terms of rate using
  the coefficients of the balanced equation:
  
  Next substitute in the given values and solve the rate of formation of O2(g):
  Rate of formation of O2(g) = (1/4) (0.53 M/s) = 0.13 M/s
  b. First determine the relationship between NO₂(g) and N₂O₅(g) in terms of rate using
  the coefficients of the balanced equation:
  
  Next substitute the given values and solve the rate of consumption of N₂O₅(g):
  
  The rate of a chemical reaction is important in our daily life as well as in industry.
  At home, it is interesting to know the rate at which a kind of food is cooked (i.e.
  boiling an egg or baking a cake). It is interesting to know the rate at which the seeds
  of maize or beans are growing till the time of harvesting. In the body, chemical
  reactions must take place at the correct rate to supply exactly the needs of the cells.
  The rate of reaction dictates the rate of production of our daily products. In order
  to meet the demand and safety standards, optimization of rate of reaction is
  nonetheless the most important subject to control and study. Kinetic studies are
  important in understanding reactions, and they also have practical implications. In
  industry, the reactions are conducted in reactors in which compounds are mixed
  together, possibly heated and stirred for a time, and then moved to the next phase
  of the process. It will be important to know how long to hold the reaction at one
  stage before moving on, to make sure that reaction has finished before starting
  the next one. By understanding how a reaction takes place, many processes can be
  improved.
  Another importance of kinetics is to understand the biological processes, especially enzyme-catalyzed reactions.
  
  Checking up 13.1
  1. What is meant by the rate of a reaction?
  2. In the following reaction:
  A Products,
  The concentration of A decreases from 0.5 mol/L to 0.4 mol/L in 10 minutes.
  Calculate the rate during this interval?
  
  13.2. Factors that change the rates of reactions
  Activity 13.2 (a)
  1. Using textbooks and internet (where possible) find out at least six factors
  that affect the rate of reaction and explain how they affect the reaction rate.
  2. Why does bread grow mold more quickly at room temperature than in the refrigerator?
  
  In chemical kinetics, the determination of how quickly or slowly reactants turn into
  products is a measure of the reaction rate. A reaction that takes long time has low
  reaction rate while a reaction that occurs quickly has a high reaction rate. Some
  conditions can be used to speed or to slow down the rate of a chemical reaction.
  The rate of reaction is defined as the rate of increase in the molar concentration
  of the products of a reaction per unit time or the rate of decrease in the molar
  concentrations of the reactants per unit time.
  The factors that influence rate of chemical reaction are:
  
  a. Concentration
  b. Temperature
  c. Pressure
  d. Catalyst
  e. Light
  f. Surface area
  In order to understand how these factors affect the rate of chemical reactions, it
  is essential to understand how reactions take place based on the collision theory.
  According to this theory, a chemical reaction takes place only by collisions between the reacting particles; it means that, for a chemical reaction to occur it is necessary for the reacting species (atoms or molecules) to collide with energy equal to or greater than the activation energy and with proper orientation.
  
  In general, a factor that increases the number of collisions between particles will increase the reaction rate and a factor that decreases the number of collisions between particles will decrease the chemical reaction rate.
  
  Activity 13.2 (b)
  Chemicals and Apparatus: 0.2M sodium thiosulphate, 2.0M hydrochloric acid,
  distilled water, measuring cylinder, stop clock (or stop watch), conical flask and
  white paper.
  Procedure:
  1. Prepare sodium thiosulphate solution (about 0.2M) by adding 2g of sodium
  thiosulphate solid to 60cm³ of water in a conical flask.
  2. Divide the solution into three equal portions, put them in three different
  conical flasks and place each conical flask on white paper marked with a cross (X).
  3. To one conical flask add 1cm³ of 2M HCl solution and swirl the mixture
  gently for uniform mixing, you must start the stop clock immediately you
  add the acid into the sodium thiosulphate solution. Record the time it takes
  for sulphur to be precipitated enough to obscure the cross when viewed from the top.
  4. To the remaining two conical flasks, add 2cm³ and 3cm³of the 2M HCl
  respectively. Record the time taken in each case for the cross to be obscured.
  Questions:
  1. State, in which conical flask, the reaction was fast. Give a reason for your answer.
  2. Write balanced equation for the reaction that takes place.
  
  Concentration refers to the number of particles dissolved in a given volume of a
  solution. The more the reacting particles are present in a given volume, the more
  opportunities of the collisions involving those particles will occur.
  The rates of many reactions depend on the concentrations of the reactants. It has
  been observed that the rate of chemical reaction is directly proportional to the
  concentration of the reactants. Thus, increasing the concentration of the reactants
  usually results in a higher reaction rate. At lower concentrations, there is less chance
  for collisions between particles. Thus, decreasing the concentrations of reactants
  results in a lower reaction rate.
  At the beginning of a reaction, when the concentration of reactants is maximum,
  the rate of reaction is also maximum. However, as the reaction proceeds the
  concentration of the reactants decreases as some amount of reactants is converted
  
  into products. The rate of the reaction decreases as the reactants are consumed and
  it increases when the products are formed (Figure 13.1). The reaction rate is directly
  proportional to the concentration, i.e it increases with increasing concentration of the reactants.
  
  Figure 13.1: Effect of concentration on the rate of reaction
  
  The rate of this reaction doubles when the concentration of oxygen is doubled.
  On the other hand, when the concentration of nitric oxide is doubled, the rate of
  reaction increases four times.
  
  13.2.2. Effect of pressure
  Pressure is defined as the force per unit area. The units of pressure are Pascals (N/
  m2) or atmospheres (1 atm); 1 atm = 101,325 Pascals (Pa). The pressure affects the
  rate of chemical reactions when the reactants or products are in gaseous state. This
  is because solids and liquids cannot be compressed, they do not change in volume
  when pressure is reduced or increased.
  Pressure has a similar effect as concentration on the reaction rate. When the pressure
  increases in the gaseous system, the number of collisions between reactants also
  increases. Consequently the rate of reaction is increased. At high pressure, the gas
  particles are closer together which can increase the collision.
  
  13.2.3. Effect of temperature
  
  The temperature is a factor that plays a big role in chemical reactions. In general,
  the rates of most chemical reactions increase as the temperature rises. When the
  temperature of a reaction is increased, the heat is supplied to the particles involved
  in the reaction. Since heat is a form of energy, these particles acquire more energy
  which enables them to move more quickly and they collide with each other more
  frequently and with more energy. Therefore more particles will overcome the
  activation energy barrier to form products. If the temperature is decreased, the
  particles will move more slowly, therefore decreasing the rate of reaction.
  Increasing the temperature of the reactants sometimes provides the activation
  energy needed to initiate a chemical reaction.
  In general, an increase in the temperature increases the rate of reaction or a decrease
  in temperature decreases the rate of reaction.
  Examples:
  (i) Some types of food such as meat can spoil quickly when kept out of a refrigerator.
  However, when they are kept in a refrigerator, the lower temperature inside the
  refrigerator retards the process of developing microorganisms that can destroy the meat.
  
  (ii) Consider the chemical reaction: CaCO₃(s) + 2 HCl(aq) ⇌ CaCl₂(aq) + CO₂(g) + H₂O(l)
  
  At 273 K (0 °C), the reaction is extremely slow and the amount of CO₂ produced is
  very small. When the temperature is raised to 323 K (50 °C), the reaction becomes
  fast and more quantity of CO2 is obtained.
  
  (iii) The reaction between zinc and hydrochloric acid can be performed at two
  different temperatures: 293 K and 308 K. It is found that the volume of hydrogen
  formed during the same interval time is higher at 308 K than at 298 K.
  The temperature dependence of the rate of a chemical reaction can be explained by
  the following Arrhenius equation.
  
  Where:
  K is Rate constant
  A is Arrhenius constant
  Ea is Activation energy
  R is Gas constant
  T is Temperature
  The logarithmic form of the Arrhenius equation is:
  
  The Arrhenius equation at the temperature T₁ and T₂ can be written as:
  
  Using the above equations, we can calculate the ratio of rate constants or activation
  energy if we know the rate constant of a reaction at two or more temperatures.
  
  Example:
  The rate constant for the reaction: H₂(g) + I₂(g) ® 2HI(g) is 2.7 x 10^-4 at 600 K and
  3.5x 10^-3 at 650 K. Calculate the activation energy of the reaction. (R = 8.31Jmol^-1K^-1).
  
  Solution:
  
  With this, Ea = 1.03 x105 J mol-1
  
  Note: The relationship between temperature and rate constant will be developed in
  more detail in unit 14.
  
  13.2.4. Effect of Catalyst on rate of reaction
  
  Activity 13.2 (c)
  Apparatuses and chemicals: Hydrogen peroxide, test tubes, wooden splint,
  manganese (IV) oxide, test tube racks, ammonium iron (II) sulphate crystals or
  iron (II) sulphate crystals, propane-1,2,3-triol.
  Procedure
  1. Put 2 cm3 of dilute hydrogen peroxide in a clean test tube and place it
  in the test tube rack. Observe the test tube from time to time and record
  your observations.
  2. Put another fresh 2cm3 of dilute hydrogen peroxide in a clean test tube
  and add a half-spatula of manganese (IV) oxide. Test the gas produced
  with a glowing splint.
  3. Put 50 cm3 of dilute hydrogen peroxide in a clean beaker and put a very
  small amount of manganese (IV) oxide. When the reaction starts add
  about 1cm3 of propane-1,2,3-triol and shake the contents of the beaker.
  Question:
  Compare the rate of the progress of gas when peroxide has both manganese (IV)
  oxide and propane-1,2,3-triol with the peroxide containing (IV) oxide only
  
  For a reaction to take place, the molecules must possess a certain energy called
  activation energy. Activation energy is the minimum amount of energy required
  for a chemical reaction to take place. If the activation energy is high, the reaction will
  be slow and a catalyst is required for lowering this activation energy.
  
  A catalyst is a substance that increases or decreases the rate of a chemical reaction
  by lowering the activation energy without itself being consumed by the reaction.
  Catalysts reduce the amount of energy required to break and form bonds during
  a chemical reaction. These substances that slow down the rate of reaction; they
  are often called negative catalysts or inhibitors. When the reaction is complete,
  catalysts remain chemically unchanged and they can be reused several times. For
  example, enzymes are catalysts that allow chemical reactions to occur at relatively
  low temperatures within the body.
  
  Other examples of catalysts:
  The decomposition of hydrogen peroxide (H₂O₂) is relatively slow; however,
  exposure to light accelerates this process and in the presence of MnO₂ as a catalyst,
  the reaction goes very fast.
  
  In the manufacture of ammonia, iron is used as a catalyst to increase the rate of
  reaction. N₂ (g) + 3H₂ (g) ® 2 NH₃ (g)
  Similarly, reaction of Sulphur dioxide and oxygen to produce Sulphur trioxide takes
  place in the presence of vanadium (V) pentoxide(V₂O₅) catalyst
  2SO₂ (g) + O₂ (g) ® 2 SO₃ (g)
  A catalyst increases the reaction rate by providing an alternative pathway (route) or
  mechanism for the reaction to follow. The effect of adding a catalyst on a reaction
  can be demonstrated on a Potential Energy Diagram (Figure 13.2)
  
  .Figure 13.2: Effect of catalyst on the rate of reaction
  Most catalysts are highly selective; they often determine the product of a reaction
  by accelerating only one of several possible reactions that could occur.
  
  Note: Except a catalyst that can lower the activation energy (Ea), increasing the
  temperature, concentration or surface area has no effect on the activation energy (Ea).
  
  Since a catalyst lowers the activation energy barrier, the forward and reverse
  reactions are both accelerated to the same degree.
  
  Types of Catalysts
  There are two main types of catalysis: homogeneous or heterogeneous catalysis.
  a. Homogeneous Catalysis
  In a homogeneous catalysis the reactants and catalysts are in the same physical
  state (phase).
  Example
  • Oxidation of iodide ions by peroxodisulphate ions:
  
  b. Heterogeneous Catalysis
  Catalysts may be classified as positive and negative catalysts. In heterogeneous
  catalysis, the reactants and catalyst are in different phases. This process is also known
  as surface catalysis or contact catalysis.
  
  Examples:
  
  a. Positive catalyst
  The substance which increases the rate of reaction is known as positive catalyst. It
  acts by decreasing the activation energy for reaction.
  
  An example of positive catalyst is the V₂O₅ used in Contact Process:
  
  b. Negative catalyst (Inhibitor or retarder)
  The substance which decreases the rate of reaction is known as negative catalyst. It
  acts by increasing the activation energy for reaction.
  Example:
  Acetanilide prevents oxidation of Na₂SO₃ by air, it acts as negative catalyst. Similarly,
  H₃PO₄ prevents the decomposition of H₂O₂ by acting as a negative catalyst for the reaction.
  
  c. Auto catalyst
  In this type of catalysis, one of the products of the reaction acts as catalyst.
  
  Example:
  In the oxidation of ethanedioic acid (oxalic acid) by potassium permanganate
  (KMnO₄), Mn²+ ion formed acts as catalyst and increases the rate of reaction.
  The initial stages of the reaction are slow but as the reaction proceeds, the reaction
  rate increases. The increase in the rate is caused by the accumulation of the catalyst
  Mn2⁺ which is formed in the initial stages of the reaction.
  
  Induced catalyst
  When a chemical reaction increases the rate of another chemical reaction, it is called
  induced catalysis.
  
  Example:
  Sodium arsenite solution is not oxidized by air but when air is passed through a
  mixture of the solution of sodium arsenite and sodium sulphite, simultaneous
  oxidation of both takes place. Thus the oxidation of sodium arsenite is induced by
  the oxidation of sodium sulphite.
  
  d. Enzymes
  Enzymes are biological catalysts which catalyze the chemical reactions occurring
  in living organisms. Enzymes have high relative molecular masses and are protein
  molecules. The enzymes are specific in catalyzing only a particular set of reactions.
  Enzyme activity depends on pH and the temperature.
  
  Examples:
  Oxidoreductases: Responsible for catalysis of reduction or oxidation of molecules.
  Transferases: Catalyze transfer of functional group from one chemical compound to the other.
  Hydrolases: Catalyze the hydrolysis reactions and the reverse reactions.
  Note: Enzymes work most effectively at body temperature (37^oC). At higher
  temperatures (above 50-60^oC), they are denatured and become ineffective. As the
  activity of enzyme decreases, the temperature of the body also decreased.
  
  13.2.5. Effect of light on rate of reaction
  Activity 13.2 (d)
  Procedure:
  1. To approximately 10 cm3 sodium chloride solution in a large test tube
  add a few cm3 of silver nitrate solution. An immediate precipitate of silver
  chloride will be formed.
  2. Quickly divide the precipitate formed into three parts, and put them into
  three separate test tubes.
  3. Put one of the test tubes immediately into a dark cupboard, the second
  may be left out on the bench and the third may be placed near a source
  of strong light (example: in direct sunshine or near to a lamp).
  Question:
  Examine the colour of the precipitates at fairly regular intervals and interpret
  your observations.
  Light is a form of energy that can affect the rate of a reaction. The rate of some
  photochemical reactions, increases with increase in the intensity of suitable light
  used. With the increase in the intensity, the number of photons in light also increases.
  Hence more number of reactant molecules gets energy by absorbing more number
  of photons and undergo chemical change.
  
  Examples:
  1. When methane reacts with chlorine in dark, the reaction rate is very low.
  It can be speeded up when the mixture is put under diffused light. In
  bright sunlight, the reaction is explosive.
  2. The rate of photosynthesis reaction is more on brighter days. At low light
  intensities, as light intensity increases, the rate of the light-dependent
  reaction (photosynthesis) generally, increases. The more photons of light
  that fall on a leaf, the greater the number of chlorophyll molecules that
  are ionized and the more ATP and NADPH are generated.
  
  However, some photochemical reactions involving the free radicals, generated in a
  chain process, are not greatly affected by the intensity of the light. Only one photon
  is sufficient to trigger the formation a free radical. This, in turn, initiates a chain
  process in which more free radicals are formed repeatedly in each cycle without the
  need of extra photons.
  When photons strike the reactant molecules, they provide necessary activation
  energy to the reactant molecules.
  There are many reactions influenced by light. For example: photosynthesis and
  photography. Such reactions are known as photochemical reactions. Other
  examples are:
  
  13.2.6. Effect of surface area on rate of reaction
  Surface area is the measure of how much area of an object is exposed. For the same
  mass, many small particles have a greater surface area than one large particle. The
  more surface contact between reactants, the higher the rate of reaction. The less
  surface contact, the lower the reaction rate.
  The rate of chemical reactions increases by increasing the surface area of reactants.
  For example, a log of wood burns slowly but if it is cut into small wooden chips, the
  burning takes place rapidly.
  The rate of a reaction between two phases depends to a great extent on the surface
  contact between them. A finely divided solid has more surface area available for
  reaction than does one large piece of the same substance. Thus a liquid will react
  more rapidly with a finely divided solid than with a large piece of the same solid.
  
  Examples:
  
  1. Large pieces of iron react slowly with acids; finely divided iron reacts
  much more rapidly.
  2. Large pieces of wood smolder, smaller pieces burn rapidly, and sawdust
  burns explosively.
  3. Iron powder reacts rapidly with dilute hydrochloric acid and produces
  bubbles of hydrogen gas because the powder has a large total surface
  area:2Fe(s) + 6HCl(aq) ⟶ 2FeCl₃(aq) + 3H₂(g)
  
  4. An iron nail reacts more slowly.
  5. Powdered calcium carbonate reacts much faster with dilute hydrochloric
  acid than if the same mass was present as lumps of marble chips or limestone.
  
  6. Solid manganese (IV) oxide is often used as the catalyst. Oxygen is given
  off much faster if the catalyst is present as a powder than as the same
  mass of granules.
  
  If one of the reactants is solid, it can be broken into smaller pieces. This increases the surface area for the other reactant to collide with it. In general, the more finely divided a solid reactant is, the greater the surface area per unit volume and the more the number of collisions between reactants. Increasing the collisions increases the rate of reaction.
  
  Checking up 13.2
  1. Explain how the following factors affect the reaction rate:
  a. Concentration
  b. Pressure
  c. Temperature
  d. Catalyst
  e. Light
  f. Surface area or particle size
  
  2. The rate constants of a reaction at 500K and 700K are 0.02s-1 and 0.07s-1
  respectively. Calculate the values of activation energy (Ea) and frequency
  factor (A).
  3. The rate constants of reaction at 700 K and 760 K are 0.011M-1s-1 and 0.105
  M-1s-1 respectively. What are the values of ‘A’ and ‘Ea’?
  
  END UNIT ASSESSMENT
  
  1. For the following multiple choice questions on the rate of reactions, choose
  the correct answer.
  Study of rate of chemical reactions is called ………………………
  a. reaction rate
  b. reaction kinetics
  c. reaction speed
  d. reaction power
  
  2. As the temperature of a reaction is increased, the rate of the reaction increases because the…………………………..
  a. reactant molecules collide less frequently
  b. reactant molecules collide more frequently and with greater energy per collision
  c. activation energy is lowered
  d. reactant molecules collide less frequently and with greater energy per collision
  
  3. Rate of reaction is change in amount of reactants or products in specific
  a. volume
  b. density
  c. time
  d. area
  
  4. Effect of pressure (in case of gases) is similar to effect of
  a. concentration
  b. molarities
  c. temperature
  d. humidity
  
  5. Factors which affect rate of chemical reaction are
  a. surface area
  b. temperature, catalysts
  c. concentration
  d. All of above
  
  6 The minimum amount of energy needed to start a reaction is called the
  A) Activation eenergy.
  B) energy of reaction.
  C) entropy of reaction
  D) reaction mechanism energy
  
  7. Of the following, __________ will lower the activation energy for a reaction.
  A) increasing the concentrations of reactants
  B raising the temperature of the reaction
  C) adding a catalyst for the reaction
  D) removing products as the reaction proceeds
  
  8. If a piece of beef liver (which acts as manganese (IV) oxide) is dipped in
  hydrogen peroxide, gas given off is:
  A) Hydrogen
  B) Ice
  C) Carbon monoxide
  D) Oxygen
  
  9. Which of the following would NOT increase the rate of reaction.
  A) raising the temperature
  B) adding catalyst
  C) increasing the concentration of the reactants
  D) increasing the volume of the container
  
  10. If the temperature of a reaction is increased by 10 oC, the new rate of
  reaction will be approximately
  A) ten times greater
  B) doubled
  C) unchanged
  D) four times greater
  
  11. Aluminium reacts with warm dilute hydrochloric acid to give aluminium
  chloride solution and hydrogen. Explain why the reaction of dilute hydrochloric
  acid with aluminium foil is fairly slow whereas the reaction with the
  same mass of aluminium powder can be extremely vigorous.
  
  12. For the reaction A B, the concentration of a reactant changes from 0.03
  M to 0.02 M in 25 minutes. Calculate the average rate of reaction using units
  of time both in minutes and seconds.
  
  13. The decomposition of dinitrogen pentoxide (N2O5) in tetrachloromethane
  (CCl4) at 318K has been studied by monitoring the concentration of dinitrogen
  pentoxide in the solution. Initially, the concentration of dinitrogen
  pentoxide is 2.33mol/L and after 184 minutes. It is reduced to 2.08mol/L. The
  reaction takes place according to the equation:
  Calculate the average rate of this reaction in terms of hours, minutes and
  seconds. What is the rate of production of nitrogen dioxide during this period?
  
  14. Calculate the activation energy for a reaction whose rate constants are
  2.15x10-1dm3mol-1s-1 and 7.25x10-3dm3mol-1s-1 at 991K and 500K respectively.
  (R = 8.3145 JK-1mol-1).
  
  15. How much faster will the reaction be at 781K than 550K if the activation energy
  of the reaction is 190 kJ/mol? (R = 8.3145 JK-1mol-1)
- Select activity S6 CHEMISTRY QUESTIONS
  
  S6 CHEMISTRY QUESTIONS File
  
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  Mark as done
  
  Uploaded 28/11/22, 10:05

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