Chemistry

Key unit Competence:

To be able to: Compare and contrast the chemical properties of the Group 17 elements and their compounds in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, the learner will be able to:

•Prepare and test halogens.

•Perform experiments to prepare and test chlorine, bromine and iodine.

•Relate the oxidizing power of Group17 elements to their reactivity.

•Relate the acidity strength of oxoacids to the number of oxygen atoms combined with the halogen.

•Compare the reactions of the halogens with cold dilute sodium hydroxide and hot concentrated sodium hydroxide solutions.

•State the uses and hazards of halogens and their compounds.

•Test for the presence of halides ions in aqueous solutions

•State the natural occurrence of halogens•Describe the extraction methods of halogens

•Explain the trends of physical and chemical properties of Group 17 elements down the group

•Describe the trends in strength acidity, volatility and reducing power of halogens hydrides

•Describe the chemical properties of chlorates, iodates, perchlorates and periodates

12.1 Natural occurrence of halogens

Group 17 elements are called halogens. They include: Fluorine (9F), chlorine (17Cl), bromine (35Br), iodine (53I) and astatine (85At). Although astatine is a member of group 17 elements, its chemistry is of little importance in this unit because all of its known isotopes are radioactive. The longest-lived isotope of Astatine (At-210) has a half-life of only 8.3 hours. All these elements have 7 electrons in their outermost shell (ns2np5) which is one electron short of the next noble gas.

Activity 12.1

1. (i).Can halogens occur freely in nature?

(ii).Give explaination to your answer in (i)

2. Which of the halogen compounds is the most abundant in nature?

All halogens occur in sea water as halide ions. Halogens also occur in minerals such as, NaCl (rock salt), CaF2, Ca(PO4)3F and NaAlF6 . Chlorides also occur in the Great Salt Lake and the Dead Sea, and in extensive salt beds that contain NaCl, KCl or MgCl2. Hydrochloric acid is present in the human stomach to help in the digestion process of food.

Bromine compounds occur in the Dead Sea and underground brines. Iodine compounds are found in small quantities in chile saltpeter, underground brines.

The best sources of halogens (with the exception of iodine) are halide salts. It is possible to oxidize the halide ions to produce the diatomic molecules by various methods, depending on the ease of oxidation of the halide ion. Fluoride is the most difficult to oxidize, whereas iodide is the easiest.

Checking up 12.1

1. State 2 locations where chlorine can be found in nature.

2. Write the chemical formulae of the compounds of halogens in nature.

3. Give the name of one lake in Rwanda where salt is abundant in water.

4. Explain the separation method you can use to get the salt crystals from the water.

5. Fluoride ion is the most difficult to oxidise into fluorine, whereas iodide ion is the easiest. Explain why.

12.2. Preparation methods of halogens

Activity 12.212.2 (a). Preparation of Cl2 and I2

1. Dissolve a small quantity of potassium iodide (2g) in a beaker containing 120 ml of water. Keep this solution.

2. Pour 20 ml of a 0.02 mole/litre potassium permanganate solution into a flat-bottomed flask.

3. Pour concentrated hydrochloric acid into potassium permanganate in the flat-bottomed flask (20 ml of a 2 mol/litre HCl).

4. Pour potassium iodide (10 ml of the potassium iodide solution) in the boiling tube.

5. Direct the delivery tube into a boiling tube containing potassium iodide solution as shown in the set-up diagram.

6. Heat the flask until there is a colour change in the boiling tube. Record the observations.

7. Add few drops of starch indicator to the resultant solution in the boiling tube

8. Record the observations.

Repeat procedures steps 1 to 7 but this time use solutions of KBr or NaBr instead of KI in the boiling tube to prepare bromine and chlorine.

Activity 12.2 (b) Preparation of bromine and iodine

1. Put 0.5 gram of MnO2 in a round bottomed flask.

2. Pour concentrated NaBr solution (5 ml of a 0.1 mol/litre) in the round bottomed flask.

3. Pour 5 ml of 1 mol/litre HCl solution in the round bottomed flask mixture.

4. Connect the apparatus to a delivery tube using a rubber stopper.

5. Heat the round bottomed flask mixture.

6. Direct the delivery tube in a solution of AgNO3 in a test tube.

7. Note the observable changes.

Activity 12.2. (c): Electrolysis of concentrated NaCl solution

a) Put 1 g of NaCl in a beaker.

b) Add water and stir using a glass rod until all the salt dissolves.

c) Pour the solution in an electrolyser.

d) Connect the electrolyser to the source of direct current and switch on.

e) Dip a test tube full of water in the NaCl solution in inverted position from above each electrode.

f ) Put 2 a little phenolphthalein indicator in the solution under each test tube.

g) Record the observations that take place for 5 minutes.

Apparatus set-up: Electrolysis of concentrated NaCl solution

12.2.1. Chlorine

Most commercial chlorine is obtained by electrolysis of chloride ions in aqueous solutions of sodium chloride or molten NaCl.

The reactions that take place are shown by the following chemical equations:

At the cathode, reduction reaction that takes place depends on the state of NaCl:If molten NaCl is used,

Sodium can be collected in metallic form or dissolved in water to produce NaOH solution:

If salt water or brine is used: hydrogen will form at the cathode, not sodium:

It is feasible to prepare chlorine, bromine and iodine in the laboratory by the chemical oxidation of the halide ion in acid solution with strong oxidizing agents such as manganese dioxide (MnO2) or sodium dichromate (Na2Cr2O7).

The reaction with manganese dioxide is represented by the equation given below:

Where X= chlorine, bromine or iodine. The halide used here may be NaX or KX

12.2.2. Fluorine

The common method for preparing fluorine is electrolysis. The most common electrolysis procedure is to use a molten mixture of potassium hydrogen fluoride, KHF2, and anhydrous hydrogen fluoride. During the process, fluoride ion, F- is oxidized to F2, whereas hydrogen ion, H+, is reduced to H2 .

Electrolysis causes HF to decompose, forming fluorine gas at the anode and hydrogen at the cathode. The two gases are separated to prevent their explosive recombination to reform hydrogen fluoride, HF.

12.2.3. Bromine

The industrial production of bromine involves the oxidation of bromide ion by chlorine:

Chlorine is a stronger oxidizing agent than bromine.

12.2.4. Iodine

Iodine can be obtained by oxidation of iodide ions by chlorine gas or another halogen that is higher than iodine in the group.

The industrial production of iodine can be done using the reduction of sodium iodate, NaIO3. The reaction is carried out using sodium hydrogen sulfite, NaHSO3, as reducing agent:

Checking up 12.2

1. Describe the preparation of chlorine in the laboratory.

2. How is electrolysis used to produce chlorine industrially?

3. Explain the method used to produce bromine.

4. Explain the hazards that may be encountered during the production of fluorine from its ores.

5. Say if the following statement is correct or wrong and justify your answer: “Bromine can be produced by oxidation of bromide by chlorine and vice-versa”.

12.3. Trends of physical and chemical properties

Activity 12.3: Testing a few properties of halogens

1. Dissolve about 0.5g of sodium chloride crystals in a test tube containing 20cm3of water. Keep this solution.

2. Pour 4 ml of the NaCl solution in 4 ml of a solution of conc H2SO4 in a test tube.3. Heat the above mixture in (NaCl and H2SO4 ) on a Bunsen burner flame.

4. Test the gas evolved with blue and red litmus papers.

5. Repeat the above procedures but using NaBr instead of NaCl.

6. Repeat the above procedures but these times use KI instead of NaCl.

7. Test the presence of any evolved gas with litmus papers.

8. Test the resultant solution of the mixture of KI and concentrated H2SO4 with a solution of starch indicator.

7. Record the observations

8. Interpret the results of the reactions in the experiment and draw a conclusion of the reactions.

Observations: The colourless solution of KI(aq) gives red or reddish mixture when chlorine gas is bubbled in it.

The red mixture turns the colourless starch indicator solution to dark–blue. On standing for a while, a black or deep violet residue settles.

Interpretations

Chlorine is liberated by the reaction between 2M hydrochloric acid and potassium permanganate solution. Chlorine displaces iodine from potassium iodide solution, which dissolves in water to give a dark-red solution, and turns starch indicator dark-blue. The greyish-black residue is due to the formation of Iodine solid.

12.3.1. Physical properties of group 17 elements down the group

a. Appearance

At room temperature and pressure, fluorine and chlorine are gases, bromine is a liquid and iodine is a solid. The table below shows the physical properties of halogens.

Table12.1: Physical properties of halogens

As shown in the table above, the boiling and melting points of halogens increase as you move down the group due to increase of the atomic mass which increases the strength of van der Waals forces.

a. Solubility in water

•Fluorine and chlorine dissolve in cold water.

•Bromine and iodine are sparingly soluble in polar solvents like water and they form coloured solutions. However they are soluble in organic solvents

like carbon tetrachloride, carbon disulphide and hydrocarbons.

•Halogen-halogen bond dissociation energy decreases from chlorine downwards as shown in the following trend:

The reason why the trend is in this order is that in a smaller molecule such

as Cl-Cl, the bond length is shorter so the internuclear attraction becomes

bigger and therefore more energy is required to break the bond Cl-Cl.

•Fluorine, F2 presents an exception because its energy of dissociation F-F is far smaller than that of Cl-Cl (opposite to the general trend from Cl2 to I2) because there is a repulsion between the lone pairs of electrons of fluorine atoms in the molecule due to the small size of fluorine.

12.3.2. Trend in chemical properties of group 17 elements down the group

Due to their valence electronic structure of ns2 np5 halogens gain easily one electron to complete the octet structure, that is why they exhibit an oxidation state (-1) in most of their compounds. They are generally considered as good oxidizing agents.

All the halogens are highly reactive; they are the most reactive group of non-metals in the periodic table. They react with metals and non-metals to form halides. But in Group 17 elements (halogens), the reactivity of halogens decreases down the group in the order: F2> Cl2>Br2 >I2. Hence each halogen displaces those below it from their salts.

The ease of capturing an electron indicates the strength of oxidising nature of halogens. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution and even in the solid phase. In general, a halogen oxidises halide ions of higher atomic number. The halogens decrease in strength of oxidising power as we descend the group.

Examples:

In addition to the normal oxidation state -1, chlorine, bromine and iodine exhibit +1, +3, +5 and +7 oxidation states. The higher oxidation states is achieved when the halogen is in combination with a small and highly electronegative atom such as fluorine and oxygen. Example: Interhalogens, oxides and oxoacids. The oxidation states of+1, +5, +7 occur in the oxides and oxoacids of chlorine, bromine and iodine

Compounds where halogens have positive oxidation states are generally very reactive, because they tend to be reduced to the oxidation state -1, the most stable oxidation state of halogens.

For example, perchlorate, ClO4-, is a powerfull oxidizing agent due the the presence of Cl(VII) species that is eager to capture electrons and get the oxidation state -1.

Being the most electronegative element, fluorine exhibits only –1 oxidation state.

a)Reactions of halogens with water

Fluorine exhibits anomalous behaviour in many properties. The anomalous behaviour of fluorine is caused by its small size and very high electronegativity.

Fluorine reacts vigorously with water to produce oxygen O2 and hydrofluoric acid HF.

Chlorine dissolves slightly in water at room temperature (25oC). At 800C and above, Cl2 gas is insoluble in water. Chlorine reacts with water to produce hydrochloric acid, HCl and hypochlorous acid, HClO:

Hypochlorous acid is a bleaching agent.Bromine dissolves in water; it reacts with water slowly to form hydrogen bromide HBr and hypobromous acid, HBrO:

Solubility of iodine in water is very low. Iodine does not react with water. Table 12.2 compares the oxidizing power of halogens when a given halogen is mixed

with an aqueous solution of another halide.

Checking up 12.3

1 State different hydrogen halides that can be formed.

2. What is the difference between hydrogen halides and hydrohalic acids?

3. Describe the anomalous behaviour of fluorine in the group (explain at least 3 differences from the rest of the halogens).

4. Write the reaction equation that takes place between:

a) Solid chloride with concentrated sulphuric acid. b)Bromine in hot concentrated sodium hydroxide solution.

5. Briefly explain the trend in volatility of hydrogen halides as you move down the group.

6. Explain why the hydrogen halides acidity increases in the order: HF < HCl < HBr <HI

7. Explain the trend in solubility of halogens in water as you move down the group.

8. You have two test tubes; one contains water, the other contains a solution of chlorine in water, and you are asked to identify them. What test are you going to do in order to identify the two test tubes?

Table 12.2: Comparison of oxidising power of halogens

b) Reaction of halogens with hydroxide:

• Cl2 reacts with excess cold dilute OH- and the reaction equation is:

With hot concentrated OH-, chlorate ClO3- , is formed:

The same reactions occur for Br2 and I2.

• F2, unlike other halogens, does not form oxosalts with alkalis and it does not form oxoacids as shown in the following reactions:

c) Reaction with metals:

Halogens react with metals to form salts, metal halides. Example: bromine reacts with magnesium to form magnesium bromide

The preparation of certain anhydrous metal halides is carried out using this method of reacting a halogen and a metal. Examples: MgCl2 (white solid), AlCl3 (white solid), FeCl3 (brown solid), FeBr3 (dark red solid) etc... are halides that can be prepared in this way.

Generally metal halides are ionic, the most ionic being the ones where the metal has low oxidation states (+1, +2).

For the same metal in the same oxidation state, the ionic character of the metal halides decreases down the group:

MF >MCl > MBr > MI; whereby M is a monovalent metal.

If a metal exhibits more than one oxidation state, the halide in a higher oxidation state will be more covalent than the one with a lower oxidation state.

Example: SnCl4 and PbCl4 are more covalent than SnCl2 and PbCl2 respectively.

d) Reaction with non-metals

Halogens react with hydrogen gas to form hydrogen halides but the affinity for hydrogen decreases from fluorine to iodine.Hydrogen halides dissolve in water to form hydrohalic acids. The acidic strength of these acids varies in the order:

HF < HCl < HBr < HI

The stability of these halides decreases down the group due to the decrease in (H–X) bond strength order:

Halogens react with oxygen to form many oxides such as, Cl2O, Cl2O6, etc.. In those oxides, except for fluorine, halogens have positive oxidation states due to the high electronegativity of oxygen; this makes them unstable and very reactive. Fluorine forms compounds with oxygen where it has a negative charge; they are rather oxygen fluoride; e.g. OF2With phosphorus, all halogens react with phosphorus to form phosphorus (III) halides of the form PX3. The reaction equation of phosphorus with bromine is represented in a general equation as:

In excess chlorine or bromine, phosphorus reacts to form phosphorus (V) chloride or bromide:

The reaction between phosphorus (III) chloride and chlorine to form phosphorus (V) chloride is reversible:

Halogens combine with other halogens to form various compounds known asinterhalogensof the types XY, XY3, XY5 and XY7, whereby X is a larger size halogen and Y is smaller size halogen.These compounds are easiest to form when Y is fluorine.

d) Bleaching property

Bleaching is the process of removing stains or colours in fabrics, especially by the use of chemical agents such as halogens, chlorine and bromine. The bleaching action of chlorine is an oxidizing action of hypochlorous acid, HOCl:

The hypochlorous acid oxidizes the dye material to form a colourless substance:

Chlorine bleaches ordinary inks which are solutions of dyes.

Bromine has a similar action but with a much less vigorous action.

Iodine has no reaction with water and therefore it is not a bleaching agent.

12.4. Preparation of Hydrogen halides

Activity 12.4.(a)

Laboratory preparation of Chlorine

Reactants:Sodium Chloride and sulfuric acid

procedure:

1.Put 50g of NaCl in round bottomed flask

2.Pour conc sulfuric acid throuth the filter fannel and heat

3.The liberated gas is pased throuth the concentrated sulfuric acid

4.Collect the gas by the Downward delivery

Note:The lower end of the thistle funnel must be dippen in acid,or you can use the funnel with Syphon.

Halogens form hydrogen halides (HF(g), HCl(g), HBr(g) and HI(g), these hydrogen halides form a series of acids when they are dissolved in water. When in aqueous solution, the hydrogen halides are known as hydrohalic acids. Hydrohalic acids are strong acids, except HF(aq) that is a weak acid.

Hydrogen halides are formed by reacting a solid halide (such as NaBr) with concentrated phosphoric (V) acid (H3PO4):

The hydrogen halides fume on contact with air and with moisture, they react to form acids:

HF and HCl can be prepared by using concentrated sulphuric acid, H2SO4, a less expensive method:

Similar equation for NaF.

Concentrated sulphuric acid is an oxidising agent and can’t be used to prepare HBr or HI. It is observed that as HBr or HI are formed, they are oxidised by sulphuric acid:

HBr and HI are prepared by hydrolysis of phosphorus trihalides.

Checking up 12.4

1. Explain why fluorine is the most oxidising element in group 17 and of all chem-ical elements?

2. Explain why the acidity of group 17 hydrohalides increases down the group.

3. Explain why HF does not follow the general trend of volatility but instead it has a higher boiling point than HCl.

4. Describe the natural state and appearance of halogen elements at room tem-perature and pressure.

5.You have two unknown sample solutions A and B. You are told that one is a solution of NaCl(aq), the other is a solution of NaI(aq). You are asked to identify them. When you add a solution of Fe3+(aq) which has a yellow colour, to both sample solutions, you get the following results: (a) Adding Fe3+(aq) to A gives a yellow solution; (b) Adding Fe3+(aq) to B gives a solution with a complicated mixture of colours between green and violet. Question: Which solution is in A, which solution is in B? Justify.

12.5. Trends in strength of acidity, volatility and reducing power of hydrogen halides

12.5.1. Acid strength

The acid strength is a measure of how an acid dissociates in water into its ions. Strong acids dissociate completely into their ions, whereas a weak acid dissociates partially into its ions.

Examples:

•Hydrochloric acid is a strong acid:

•Hydrofluoric acid is a weak acid:

The acid strength of hydrogen halides (HX) increases down the group from HF to HI:

Because of the very electronegativity and small size of F, it forms very strong H-F bond. In water, it is slightly dissociated to give few H+ ions in solution; therefore it is a weak acid.

This is due to the strong bond H-F; in other words, the bond strength HX decreases down the group.

12.5.2. Volatility

Volatility is the state of having a low boiling point to evaporate easily. The boiling points of hydrogen halides generally increase down the group. This is because the van der Waals’ forces increases with size (the forces increase with increase of the surface of contact between molecules).

The high electronegativity of fluorine atom is the root cause of the strong hydrogen bonds. It is liquid at room temperature while other hydrogen halides are gases.

The trend in volatility:

HF is an exception to the general trend (it has a high boiling point of 19.9 0C). It is almost liquid at room temperature while the rest are gases because there are strong hydrogen bonds between HF molecules. For the other Hydrogen halides, the boiling point increases with increasing molecular mass.

12.5.3. Reducing power

The reducing power of hydrogen halides is dependent on the ability of their anions (halide ions) to liberate electrons.

It becomes easier to oxidise the hydrogen halides as we go down the group because the electron on the halide ion becomes less attracted to the nucleus as we move down the group.

Hence the order of reducing power is: I-> Br- > Cl-> F-

Given that halogens are the best oxidizing agents since they tend to capture electrons from other elements to achieve the octet electronic configuration, the halide ions are not expected to be good reducing agents.

However, an iodide ion, due to its big size, can easily liberate an electron and therefore act as a mild reducing agent.

Hence HI is the strongest reducing agent of all hydrogen halides.

HI can reduce Cl2, Br2, O2, Fe3+ salts, etc... to produce molecular iodine, I2:

That is why when a colorless aqueous solution of HI is exposed to air, it turns brown because of the presence of I2.

Concentrated H2SO4 oxidises Br- and I- to form Br2 and I2 elements respectively.

Concentrated H2SO4 cannot oxidise F- and Cl- ions.

12.5.4. Tests for halide ions in aqueous solution

Test of substance X with an unknown anion

Activity 12.5

Identification of ions:

i) You are provided with a solution of X substance.

ii) Put 1 ml of X solution in each of the 4 test tubes.

iii) Add in each test tube the reagent solutions as indicated in the table be-low.

iv) Note down the observations for interpretation later in each test.

Halide ions are detected by precipitation reactions using silver nitrate solution AgNO3 or lead (II) nitrate Pb(NO3)2.The ionic equations below illustrate the general equation reactions:

i) Testing using AgNO3 and ammonia solution, NH3

The first step is to dissolve 1 ml sample solution to be tested in about 2 ml of dilute nitric acid.The table given below shows the observable changes when we add the silver nitrate solution, AgNO3(aq), followed by ammonia solution, NH3(aq)

The addition of NH3 ammonia solution is called ‘confirmatory test”, i.e. to confirm the first one.

Table 12.5: Tests of halides

If the test uses Pb2+(aq) and Cl- ions, heating the precipitate, it will dissolve; PbCl2(s) is insoluble in cold water, but soluble in hot water.

g) Lead nitrate Pb(NO3)2 solution

Chloride, Cl- ion: a white precipitate solution is formed in cold water, the precipitate dissolves when it is hot water

Checking up 12.5

1. You are given 1ml of solution Y in which you add 1 ml of AgNO3 solution and there is formation of a white precipitate that dissolves on adding 1 ml of ammonia solution. Deduce the anion present in solution Y.

2. Write a balanced equation for the reaction of: a) Decomposition of KClO4b) Oxidation of Cl- by ClO3-in acid medium.

3. Explain how ClO3- is prepared in the laboratory from chlorine.

4. Explain why fluorine forms only 1 oxoacid (HFO).

12.6. Chemical properties of chlorates, iodates, perchlorates and periodates

Activity 12.6 (a)

1. Put 0.5g of solid KClO3 in a pyrex test tube.

2. Using a test tube holder, heat the KClO3 strongly while placing a glowing splint on the test tube outlet.

3. Note the observation as heating takes place.

4. Dissolve a small portion of the solid residue in water.

5. Put a little solid residue in 1 ml of aqueous solution of AgNO3.

6. Record the observed changes.

7. Deduce the chemical changes that took place. a)While heating b) On adding the residue to AgNO3 solution.8. Write the equations for the reaction that took place in 7 (a) and 7(b).

Activity 12.6 (b)

1. Dissolve 0.5g of KIO3 solid in water and make 20 ml of solution.

2. Add 20 ml of a 1mol/litre solution of HCl acid to the solution of KIO3.

3. Pour 2 ml of the above mixture in a test tube.

4. Add 1 ml of a solution of KI to the solution containing KIO3 and HCl acid.

5. Note down the changes observed.

6. Identify the reagent compound that is oxidizing and the one that is reducing in the above reactions.

Due to the high electronegativity and small size, fluorine forms only one oxoacid, HOF known as fluoric (I) acid or hypofluorous acid. Other halogens form many oxoacids. Most of them cannot be isolated in pure state. They are stable only in aqueous solutions but they form stable salts.

Halogens normally form 4 series of oxoacids, namely: Hypohalous acids (HXO), halous acids (HXO2 ), halic acids (HXO3) and perhalic acids (HXO4).

The difference betweem oxoacids is the oxidation state of the halogen in each oxoacids or halates.

Table 12.3: Oxohalic acids

a) Acid character

Table 12.3 shows that the acid strength increases with increasing oxidation state of the halogen.

Example: Among of the four oxoacids of chlorine, the acidity increases as follows:

HClO < HClO2 < HClO3 < HClO4

In other words the acidity increases with the number of oxygen atoms forming bonds with the halogen atom.

Example: Potassium chlorate decomposes at 400oC to give potassium perchlorate and potassium chloride.

At higher temperatures, the perchlorate decomposes to give the chloride and oxygen.

Potassium bromate and iodate give similar products on strong heating. No perbromate and periodate is formed during the heating. The halates are also strong oxidizing agents as their corresponding acids behave. Potassium chlorate (V) oxidizes hot concentrated hydrochloric acid to an explosive bright yellow gaseous mixture

Checking up 12.6

1. Write a balanced equation for the reaction of:

a) Decomposition of KClO4

b) Oxidation of Cl- by ClO3-in acid medium.

2. Explain why fluorine forms only 1 oxoacid (HFO).

12.7. Uses of halogens and their compounds

Activity 12.7

1. When you want to eat food, salt is dissolved in it. Indicate the chemical composition of table salt and its natural occurrence.

2. Chlorine compounds are used in the treatment of water. Explain how chlorine re-acts to be a good disinfectant in water treatment.

3. a)Write the observations of the phenomenon that takes place when electrolysis of a concentrated solution of chlorine is carried out in the laboratory.

b) Deduce the product of reaction that is formed at the anode.

12.7.1. Uses of Halogens

Halogens and their compounds have many applications and uses in different domains:

•Fluorine is used in production of synthetic fibres. It is also used as an ingredient in toothpastes as well as in the manufacture of HF.

•NaCl is the main food seasoning table and kitchen salt; it is also used in many industrial processes such as: NaOH production, soap manufacturing, etc...

•Chlorine is used to synthesise products for bleaching clothes, papers, etc. it is used as an antiseptic and as fungicide.

•Chlorine is used in water treatment as well as in the manufacture of plastics such as Polychloroethene (polyvinyl chloride: PVC).

•Chlorine is used to make DDT (Dichlorodiphenyltrichloroethane: a banned chemical), and other chlorinated aromatic compounds used as pesticides.

•A Chlorine compound, HCl is important in in the human stomach for digestion.

•Halogen elements are used to manufacture polymers.

•Bromine is used in photographic industry (film manufacture).

•Iodine is used in food in form of iodised salt and in drugs to fight against goitre and to kill bacteria in wounds, etc.

12.7.2. Hazards caused by group 17 elements

Bromine effects

•On heating, toxic fumes are formed.

•Reacts violently (explosively) with many compounds.

•Attacks plastics, rubber and coatings.

Chlorine effects

•It reacts violently with many compounds like ammonia and may cause fire and explosion.

•It attacks many metals in the presence of water.

•It attacks plastics, rubber and coatings.

Chlorine oxide effects

•It may explosively decompose when it encounters shock and friction then it may explode on heating.

•It reacts violently with mercury, phosphorus, sulphur, etc causing fire and explosion hazard.

Fluorine effects

•It reacts violently with water to produce toxic and corrosive vapours: ozone and hydrogen fluoride.

•It reacts violently with ammonia, metals, oxidants, etc, to cause fire and explosion.

Hydrogen bromide effects

•It reacts violently with strong oxidants and many organic compounds to cause fire and explosion.

•It attacks many metals forming flammable hydrogen gas.

Hydrogen fluoride effects

•It reacts violently with many compounds causing fire and explosion.

•On contact with air, it emits corrosive fumes which are heavier than air.

•It attacks glass and other silicon-containing compounds.

SF6 effects

•The substance decomposes in a fire to produce toxic fumes of sulphur oxides and hydrogen fluoride •When it is heated, there is formation of toxic fumes.

CFCs: Chlorofluorocarbons

Chlorofluorocarbons are hydrocarbons (alkanes) where some or all of the hydrogen atoms have been replaced by chlorine and fluorine atoms. Most of those compounds are stable and unreactive at high temperature. That is why they were used as aerosol propellants, refrigerants, and solvents.

One example of CFC is Freon 12 (CCl2F2) that was used in refrigerators.Because of their chemical inertness, CFCs can diffuse unchanged into the upper atmosphere up to the ozone layer (10-15 km). There, photochemical reaction cause them to break down into radicals, .Cl. Radicals, being very active chemical species, react with ozone to form ordinary oxygen, hence destruction of ozone layer. For this reason, their use has been discourage or banned all over the World.

Checking up 12.7

1. Indicate at least 3 uses of chlorine compounds in daily life.

2. State at least 2 uses of bromine and its compounds in daily life.

3. Why do we use iodised salt as table and cooking salt?

4. Discuss the importance of halogen compounds in biological area.

5. Give at least 2 negative effects of halogens compounds

12.8. End of unit assessment

1. Chlorine is used in the preparation of bleaching agents and iodine.

(a) Write the reaction of Chlorine with water and mention the bleaching group

(b) Bromine is extracted from sea water by oxidising bromideions with chlo-rine gas.

(i) Write the ionic equation for this reaction.

(ii) Explain why chlorine is strong oxidising agent compared to bromine

2. Aqueous silver nitrate can be used as a test for halide ions. A student decided to carry out this test on a solution of magnesium chloride. The bottle of mag-nesium chloride that the students used showed the formula MgCl2.6H2O. The student dissolved a small amount of MgCl2.6H2O in water and added aqueous silver nitrate solution.

a)What would the student expect to see after adding aqueous silver nitrate?

b) Write an ionic equation for this reaction. Include state symbols.

c) Using aqueous silver nitrate, it is sometimes difficult to distinguish between chloride, bromide and iodide ions.What test are you going to use to distinguish between them? Explain

d) When carrying out halide test with aqueous silver nitrate, it is important that distilled water is used for all solutions, rather than tap water.

Suggest why it is done this way.

1. Chlorine gas is bubbled through an aqueous solution of bromide ions and alsothrough an aqueous solution of iodide ions. What will be observed? Explain the observation by chemical equations

2. A student carried out experiments using chlorine gas (Cl2) with sodium hydrox-ide solution and the equation for this reaction is shown below:

The student repeated the procedure in but with hot concentrated sodium hydroxide. A different reaction took place in which sodium chlorate (V) was formed instead of NaClO.

Write the equation for the reaction to produce sodium chlorate (V).