NON-METALS : HALOGEN
Halogens:
Group 17(VIIA) of the periodic table consists of five elements, i.e., fluorine, chlorine, bromine, iodine and astatine. They are all non-metallic elements and are collectively named as halogens. The name halogen (Greek, halos = sea salts, genes = born) meaning sea salt formers was given to them by Schweiger in 1811 because the salts (chlorides, bromides and iodides) of the first three elements occur in sea water. Of all the halogens, fluorine is the most reactive and hence is also called super halogen, Astatine, however, is radioactive and hence occurs in nature only in traces. Therefore, not much is known about its chemistry.
Why diatomic? The electronic configurations of halogens show that each of them has one electron less than the nearest inert gas. As a result, they are very reactive elements. They readily accept one electron to form uninegative ions or share their single unpaired electron of p-orbitals with other atoms to form covalent bonds. Thus, all halogens exist as diatomic molecules.
General characteristics of halogens:
Some important properties of halogens are discussed below:
i) Electronic configuration and oxidation states:
| Fluorine | 1s2 2s2 2p5 |
| Chlorine | [Ne]3s2 3p5 |
| Bromine | [Ar]3d10 4s2 4p5 |
| Iodine | [Kr]4d10 5s2 5p5 |
| Astatine | [Xe]4f14 5d10 6s2 6p5 |
The melting and boiling points increase down the group because of the Vander Waals forces. The size of the molecules increases down the group. This increase in size means an increase in the strength of the van der Waals forces.
F<Cl<Br<I<At
(Group| Fluorine | -220 | -188 |
| Chlorine | -101 | -35 |
| Bromine | -7.2 | 58.8 |
| Iodine | 114 | 184 |
| Astatine | 302 | 337 |
iii) Atomic size/radius:
The size of the nucleus increases down a group (F < Cl < Br < I < At) because the numbers of protons and neutrons increase. In addition, more energy levels are added with each period. This results in a larger orbital, and therefore a longer atomic radius.
| Fluorine | 71 | 133 |
| Chlorine | 99 | 181 |
| Bromine | 114 | 196 |
| Iodine | 133 | 220 |
| Astatine | 150 |
iv) Ionization Energy
If the outer valence electrons are not near the nucleus, it does not take as much energy to remove them. Therefore, the energy required to pull off the outermost electron is not as high for the elements at the bottom of the group since there are more energy levels. Also, the high ionization energy makes the element appear non-metallic. Iodine and astatine display metallic properties, so ionization energy decreases down the group (At < I < Br < Cl < F).
| Fluorine | 1681 |
| Chlorine | 1251 |
| Bromine | 1140 |
| Iodine | 1008 |
| Astatine | 890±40 |
The number of valence electrons in an atom increases down the group due to the increase in energy levels at progressively lower levels. The electrons are progressively further from the nucleus; therefore, the nucleus and the electrons are not as attracted to each other. An increase in shielding is observed. Electronegativity therefore decreases down the group (At < I < Br < Cl < F).
| Fluorine | 4.0 |
| Chlorine | 3.0 |
| Bromine | 2.8 |
| Iodine | 2.5 |
| Astatine | 2.2 |
vi) Electron Affinity
Since the atomic size increases down the group, electron affinity generally decreases (At < I < Br < F < Cl). An electron will not be as attracted to the nucleus, resulting in a low electron affinity. However, fluorine has a lower electron affinity than chlorine. This can be explained by the small size of fluorine, compared to chlorine.
| Fluorine | -328.0 |
| Chlorine | -349.0 |
| Bromine | -324.6 |
| Iodine | -295.2 |
| Astatine | -270.1 |
vii) Reactivity
The reactivities of the halogens decrease down the group ( At < I < Br < Cl < F). This is due to the fact that atomic radius increases in size with an increase of electronic energy levels. This lessens the attraction for valence electrons of other atoms, decreasing reactivity. This decrease also occurs because electronegativity decreases down a group; therefore, there is less electron "pulling." In addition, there is a decrease in oxidizing ability down the group.
States of halogens at room temperature:| Solid | Iodine | Violet |
| Astatine | Black/Metallic [Assumed] | |
| Liquid | Bromine | Reddish-Brown |
| Gas | Fluorine | Pale Yellow-Brown |
| Chlorine | Pale Green |
viii) Appearance/Colors
The halogens' colors are results of the absorption of visible light by the molecules, which causes electronic excitation. Fluorine absorbs violet light, and therefore appears light yellow. Iodine, on the other hand, absorbs yellow light and appears violet (yellow and violet are complementary colors, which can be determined using a color wheel. The colors of the halogens grow darker down the group:
- Fluorine → pale yellow/brown
- Chlorine → pale green
- Bromine → red-brown
- www.crscientific.com/brominecell4.jpg
- Iodine → violet
- genchem.chem.wisc.edu/lab/PTL...ments/I/I.jpeg
- Astatine* → black/metallic
- www4.msu.ac.th/satit/studentP...t/astatine.jpg
In closed containers, liquid bromine and solid iodine are in equilibrium with their vapors, which can often be seen as colored gases. Although the color for astatine is unknown, it is assumed that astatine must be darker than iodine's violet (i.e. black) based on the preceding trend.
ix) Oxidation States of Halogens
As a general rule, halogens usually have an oxidation state of -1. However, if the halogen is bonded to oxygen or to another halogen, it can adopt different states: the -2 rule for oxygen takes precedence over this rule; in the case of two different halogens bonded together, the more electronegative atom takes precedence and adopts the -1 oxidation state.
One third exception to the rule is this: if a halogen exists in its elemental form (X2), its oxidation state is zero.
| Fluorine | (always) -1* |
| Chlorine | -1, +1, +3, +5, +7 |
| Bromine | -1, +1, +3, +4, +5 |
| Iodine | -1, +1,+5, +7 |
| Astatine | -1, +1, +3, +5, +7 |
Comparative Study on Preparation
In principle all halogens can be prepared by the oxidation of their corresponding halides either electrolytically or chemically by means of oxidizing agents stronger than halogens.
2X → X2
The decreasing ease of oxidation of halides is in the order:
I- > Br- > Cl- > F-
Thus I- is easy to oxidise while F- is difficult to oxidise. Since fluorine is the strongest oxidizing agent, therefore it cannot be prepared by chemical oxidation of fluorine ions. Thus, fluorine can be prepared only by electrolytic method.
(a) Preparation of Chlorine
1. From concentrated hydrochloric acid:
Any strong oxidising agent like MnO2, K2Cr2O7 oxidises concentrated hydrochloric acid to chlorine. The reactions of some oxidising agents with concentrated hydrochloric acid are:
2KMnO4 + 16HCl → 2KCI+ 2MnCl₂ + 5Cl2 + 8H₂O
HNO3 +3HCl → NOCI + CI2 +2H₂O
CaOCl2 + 2HCl → CaCl2 + Cl2 + H₂O
Instead of hydrochloric acid, sodium chloride and concentrated sulphuric acid may be used with the oxidising agent. The sulphuric acid first reacts with the salt to form hydrogen chloride, which is then oxidised by the oxidising agent.
2NaCl + 2H2SO4 + MnO2 → MnSO4 + Na2SO4 + Cl₂ + 2H₂O
2. Laboratory preparation:
i) With application of heat: In laboratary chlorine gas is prepared by heating manganese dioxide with concentrated hydrochloric acid.
MnO2 + 2HCI → MnCl₂ + H₂O + [O]
2HCl + O → H2O + Cl2
MnO2 + Conc. 4HCI → MnCl2 + 2H2O + Cl₂ ↑
Similarly Br2 and I2 , can also prepared by heating HBr and HI respectively with MnO, But HBr and HI are unstable compound can't used as laboratory reagent.
Instead of hydrochloric acid, a mixture of sodium chloride and concentrated sulphuric acid is heated with manganese dioxide.
NaCl + H₂SO4 → NaHSO4+HCI] x2
MnO2 + H2SO4 → MnSO4 + H₂O + O
2HCl + O → H2O + Cl2
2NaCl + Conc. 3H2SO4 + MnO2 → 2NaHSO4 + MnSO4 + 2H2O + Cl₂
ii) Without application of heat: In laboratory chlorine gas can be prepared by dropping concentrated hydrochloric acid on potassium permanganate.
2KMnO4 + 6HCI → 2KCl + 2MnCl2 + 3H2O + 5(O)
2HCl + (O) → H2O + Cl2]*5
2KMnO4 + Conc. 16HCI → 2KCL + MnCl2 + 8H2O + 5Cl2
(b) Preparation of Bromine
Laboratory Preparation
Bromine is prepared in the laboratory by heating sodium bromide (or Potassium bromide) with conc. H2SO4 and manganese dioxide. The following reaction occurs and brown vapours of bromine are condensed to liquid bromine and collected.
2NaBr + MnO2 + Conc. 3H2SO4 → 2NaHSO4 + MnSO4 + Br2 + 2H₂O
Since bromine is poisonous, it is prepared in a fume cup-board.
(c) Preparation of lodine
Laboratory Preparation
lodine is prepared in the laboratory by heating a mixture of potassium iodide, MnO, and concentrated sulphuric acid
2KI + 3H₂SO4 + MnO₂ → 2KHSO4 + MnSO4 +2H₂O +I2
It may also be obtained by heating sodium iodide, MnO2 and conc. H2SO4
2Nal + 3H₂SO4 + MnO2 → 2NaHSO4+ MnSO4 + 2H₂O + I2
Chemical properties of halogens
All the halogens are chemically very reactive. Fluorine, however, is the most reactive of all the halogens. The high reactivity of halogens is due to their
(i) low bond dissociation energy and
(ii) high electron affinity.
As we move down the group, both the bond dissociation energy and electron affinity decreases and thus the reactivity of halogens also decreases in the same order, i.e.
F2 > Cl2 > Br2 > l2
Some important chemical reactions of halogens are discussed below:
a. Reaction with hydrogen
All the halogens react with hydrogen to form halogen acids (HX). The reactivity towards hydrogen, however, decreases from fluorine to iodine.
H2 + X2 → 2HX (Where X = F Cl, Br or I)
For eg.
H2 + F2 → 2HF
H2 + Cl2 → 2HCl
b. Reaction with metals
Halogens combine with most of the metals to form their corresponding halides.
2Na + F2 → 2 NaF
2Na + Cl2 → 2NaCl etc.
Note that the metals with variable valency give higher oxidation state halide e.g. especially with F2, Cl2 and Br2.
Cu + X2 → CuX2 (X=Cl, Br, F₂)
But 2Cu + I2 → 2CuI
c. Reaction with Non metals
All halogens combine directly with a number of non metals to form their corresponding halides.
Xe + 3F2 → XeF6
P4 + F2 → PF3 OR PF5 and likewise with other halogens.
d. Reaction with alkalies
(i) With Fluorine: Fluorine reacts with cold and dilute alkali solution to give oxygen difluoride (OF2).
2F2 + 2NaOH (dil.) → 2NaF + H₂O + OF₂
However, with hot and concentrated alkali, fluorine gives oxygen.
2F2 + 4NaOH (conc.) → 4NaF + 2H₂O + O2
(ii) With Chlorine: Chlorine reacts with cold and dilute alkali solution to give chloride and hypochlorite.
2NaOH(dil) + Cl₂ → NaCl + NaOCl + H₂O
However, with hot and concentrated alkali, chlorine gives chloride and chlorate.
(iii) With Bromine: Bromine reacts with cold and dilute alkali solution to give bromide and hypobromite.
2NaOH(dil.) + Br₂ → NaBr + NaOBr + H₂O
However, with hot and concentrated alkali, bromine gives bromides and bromates.
3Br₂ + 6NaOH (conc.) → 5NaBr+ 3H2O+NaBrO3
(iv) With Iodine: lodine reacts with cold and dilute alkali solution to give hypo iodous acid
2NaOH(dil) + I2 → Nal + HOI
However, with hot and concentrated alkali Iodine gives lodide and lodate.
3I2 + 6NaOH (conc.) → 5Nal + 3H2O + NalO3
e. Reactions with water
Flourine is highly soluble in water. It reacts vigorously with water to give HF, O2 and O3.
2H2O + 2F2 → 4HF + O2
3H2O + 3F2 → 6HF + O3
Chlorine and bromine are fairly soluble in water to give chlorine water and bromine water respectively. A part of dissolve chlorine or bromine reacts with water less vigorously forming a mixture of hydrohalic acid (HX) and hypohalous acid (HOX):
X2 + H2O → HX + HOX (X = Cl or Br)
Example:
Cl₂+ H₂O → HCI + HOCI (Hypochlorous acid)
Br₂+H₂O → HBr + HOBr (Hypohromous acid)
Hypohalous acid being unstable when exposed to sunlight decompose to give oxygen.
HOX → O + HX
O + O → O2
Example:
2HOCI → O₂+ 2HCI
2HOBr → O₂+2HBr
lodine has the least affinity for water and is only slightly soluble in it. However, it dissolves in 10% aqueous solution of KI due to the formation a complex ion, i.e. l3.
I2 + KI (rightleftharpoons) KI3 (potassium triodide)
f. Oxidising agent
Since all the halogens have a strong tendency to accept electrons, they are strong oxidising agents. Their oxidising power decreases from F2 to I2 Thus, fluorine is the strongest oxidising agent of all the halogens. As such it oxidises all other halide ions to the corresponding halogens:
F2 +2X- → 2F- + X2 (X = Cl, Br or l)
Similarly, Cl2 will oxidise Br- ions to Br2 and I- ions to l2
Cl + 2X → 2Cl- + X2(X = Brorl)
While Br2 will oxidise only I- ions to l2
Br2 + 2I- → 2Br- + I2
Chlorine, bromine and lodine (X) oxidise with different compounds to form corresponding compounds.
g. Bleaching properties of halogens
F2 being the strongest oxidising agent destroy the substances to be bleached in the presence of moisture. Thus it can not be used as bleaching agent. I2 however being weakest oxidising agent has no bleaching property. While Br2 is mild bleaching agent. Cl2 gas, the colour is discharged due to its oxidation to colourless flower. In moist, Cl2 produce HCI + HOCI.
Thus Cl2 is used as bleaching agent. When a moist coloured flower is placed in a jar containing nascent oxygen.
Cl2 + H2O → HCl + HOCl
HCIO → HCI+ [O]
Coloured flower + [O] → colourless
Thus Cl2 bleaches colour by oxidation and bleached matter does not regain its original colour. Hence bleaching by Cl2 is permanent.
It is noted that Cl2 can not bleach printed matter and pencil marked due to inertness of carbon. Also dry Cl2 does not bleach.
h. Reaction with ammonia
Halogens give different product with ammonia
(i) With fluorine: Fluorine reacts with ammonia violently to produce nitrogen.
8NH3 + 3F2 -> 6NH4F + N2
(ii) With Chlorine: Chlorine reacts with excess of ammonia to produce nitrogen.
8NH3 +3Cl2 → 6 NH4Cl + N2
But excess of chlorine gives NCl3
(iii) With bromine: Bromine reacts with excess of ammonia to produce N2
8NH3 + 3Br2 → 6NH4Br + Br2
But excess of bromine gives nitrogen tribromine hexa ammoniate.
(iv) With iodine: lodine reacts with liquid ammonia to form a brown ppt. of nitrogen tri-iodide ammonate.
5NH3 + 3I2 → 3NH4I + NI3.NH3
Test for Halide lon
Test for chloride ion (CI)
In order to test the presence of CI ion in a given solution, about 1 ml of solution is taken in clean and dry test tube and is acidified with 1-2 drops of conc. HNO3 Then the acidified solution is then heated and then cooled under tap water and 1-2 drops of AgNO3 solution is added to it. A white ppt. will be seen which indicates the presence of Cl ion in the given solution.
Cl- + AgNO3 → AgCI + NO3- (aq.)
For further confirmation, the white ppt. is dissolved in NH,OH solution and again reappears by adding few drops of dil. HNO3 solution.
AgCI + 2NH4OH → [Ag(NH)]+ + 2H2O + CI-
[Ag(NH3)2]+ + CI- + 2HNO3 → AgCl + 2NH4+ + 2NO3-
If ppt, is dissolved soon and also reappears soon it confirms the presence of CI-.
Test for bromide ion (Br)
In order to test the presence of Br- ion in a given solution, about 1 ml of solution is taken in clean and dry test tube and is acidified with 1-2 drops of conc. HNO3 Then the acidified solution is then heated and then cooled under tap water and 1-2 drops of AgNO3 solution is added to it. A pale yellow ppt. will be seen. Which indicates the presence of Br ion in the given solution.
Br- (aq) + AgNO3 (aq) → AgBr + NO3- (aq)
For further confirmation, the pale yellow ppt. is seen by adding few drops of dil. HNO3 solution.
Test for iodide ion (I-)
In order to test the presence of Iodide ion in a given solution, about 1 ml of solution is taken in clean and dry test tube and is acidified with 1-2 drops of conc. HNO3. Then the acidified solution is then heated and then cooled under tap water and 1-2 drops of AgNO3 solution is added to it. A yellow ppt. will seen which indicates the presence of ion in the given solution.
I- (aq) + AgNO3 (aq) → Agl + NO3- (aq)
Yellow ppt. does not dissolve in NH4OH and dil. HNO3
Uses of Halogen (Cl2, Br2, I2)
A. Uses of Chlorine
(i) Chlorine is used in sterilization of drinking water.
(ii) Large quantities of chlorine are used for bleaching paper, pulp and textiles.
(iii) It is used in the manufacture of inorganic chemicals such as hydrochloric acid, sodium hypochlorite (NaOCI), belaching powder (CaOCl), phosphorus trochloride (PCI). phosphorus pentachloride etc.
(iv) It is used in the manufacture of vinyl chloride which is the starting material for plastic polyvinyl chloride (PVC)
(v) It is used in the manufacture of insecticides like D.D.T., germicides, dyes and drugs.
(vi) It is used in the manufacture of chlorinated organic solvents like CHCI, CCI CCI CCI, (westron), CHCI-CCI, (westrosol) which are used in dry-cleaning and degreasing machinery.
(vii) It is used in the manufacture of chlorates which are used in flash light powders and explosives.
B. Uses of Bromine
(i) The main use of bromine is in the manufacture of ethylene bromide (BrCH-CH,Br) which is used as an additive to leaded petrol.
(ii) Bromine is also used to prepare AgBr which is extensively used in photography.
(iii) In the laboratory, bromine and brominated water are used in the detection of unsaturation (double or triple bonds) in an organic compound.
(iv) It is used in the manufacture of dyes, drugs, etc.
(v) It is used in the manufacture of benzyl bromide which is an effective tear gas.
Uses of Iodine
(i) Iodine has powerful germicidal properties and a 2% solution of iodine in alcohol called tincture of iodine is used as an antiseptic.
(ii) lodine is also used for the preparation of iodoform and potassium iodide.
(iii) Deficiency of iodide ion in the diet of human beings leads to goitre (enlargement of thyroid gland). To check this, traces of iodide (0.5 g per kg of NaCl) in form of Nal or Kl is added to the salt and this type of common salt is called iodised salt.
Comparative study of preparation of haloacids
A. Hydrogen Chloride (HCI)
Laboratory Preparation
In the laboratory, hydrogen chloride gas is prepared by heating common salt with concentrated sulphuric acid at low temperature.
NaCl + H₂SO4 → NaHSO4 + HCI At low temperature
2NaCl + H₂SO4 → Na2SO4 +HCI
Drying of HCI gas: HCI gas is dried by passing through concentrated sulphuric acid and is collected in a gas jar, by the upward displacement of air because it is heavier than air. As it is very soluble in water, it cannot be collected over water.
Preparation of hydrochloric acid
Hydrogen chloride gas is highly soluble in water. To prepare hydrochloric acid (HCI solution) care should be taken. If the delivery tube is directly dipped into water, water would suck back into generating flask and cause an explosion with hot sulphuric acid and damaging the apparatus and harming people around.
To prevent this sucking back of water, anti-suction device is made by joining inverted funnel to delivery tube and the rim of funnel just touch the surface of water. Hydrogen chloride gas is dissolution of HCI gas in water absorbed by water and the solution becomes stronger and stronger till a saturated solution called concentrated hydrochloric acid is obtained.
Uses of Hydrochloric acid
Next to sulphuric acid, hydrochloric acid is the most widely used acid in industry.
(i) It is used in making pickling baths for removing the oxide coating from iron before plating or galvanising.
(ii) Large quantities of HCI are used in the hydrolysis of starch to glucose in the manufacturing of corn syrup.
(iii) It is also used in the of textiles, dyes and chemicals.
(iv) Hydrochloric acid is important in the stomach, where it aids in digestion.
B. Hydrogen Bromide (HBr)
Laboratory Preparation: In the laboratory hydrogen bromide is prepared by the action of bromine on moist red-phosphorus.
P4 + 6Br2 + 12H2O → 4H3PO3 + 12HBr
HBr can't prepare in the same way as HCI by heating bromide with concentrated H2SO4. It is because HBr is strong reducing agent and concentrated H2SO4 is reduced to SO2.
NaBr + H2SO4 → Na2SO4 + 2HBr
2HBr + H2SO4 → SO2 + 2H2O + Br2
Drying of HBr gas : HBr gas is dried by passing through anhydrous CaCl2 or P2O5 to remove moisture. Concentrated sulphuric acid cannot be used as it oxidises HBr to bromine. Similarly lime also cannot be used for the same reason.
H₂SO + 2HBr → 2H2O + SO2 + Br2
CaO+2HBr → CaBr+H₂O
Preparation of Hydrobromic Acid
Hydrobromic acid can be prepared by dissolving HBr gas in water using antisuction device by as in hydrochloric acid.
Uses of Hydrogen Bromide
The properties of hydrogen bromide are like those of hydrogen chloride. Therefore hydrogen bromide has not any extensive use because of its high cost. Any how, it may be used
(i) in the preparation of bromides.
(ii) in organic chemistry as a brominating agent.
(iii) as a reducing agent.
(iv) as a laboratory reagent.
C. Hydrogen Iodide
Laboratory Preparation:
In the laboratory, hydrogen iodide is prepared by dropping water on a mixture of moist red phosphorus and iodine.
P4 + 6I2 → 4H3PO3 + 12HI
But HI can't prepare in the same way as HCI by heating iodide with concentrated H2SO4 It is because HI is strong reducing agent and concentrated H2SO4 is reduced to SO4.
Nal + H2SO4 → Na2SO4 + 2HI
2HI + H2SO4 → SO2 + H2O + I2
As in the case of HBr, hydrogen iodide is also dried by passing over anhydrous calcium chloride. Because of its still greater reducing power than HBr, sulphuric acid and lime cannot be used as dehydrating agents.
H2SO4 + 2HI → 2H2O + SO2 + I2
CaO + 2HI → Cal2 + H2O
Preparation of Hydroiodoic Acid
Hydroiodoic Acid can be prepared by dissolving HI gas in water using antisuction device by as in hydrochloric acid.
Uses of Hydriodic Acid
Because of its instability, hydrogen iodide does not find much application. It may be used as a reducing agent and also for the preparation of iodides and iodine.
Properties of Hydrogen Halides
a. Physical properties
1. Hydrogen halides are highly soluble in water.
2. They are colorless gas with pungent smell.
3. They form constant boiling mixture with water.
4. They can be liquefied into colorless liquid and solidified into a white crystalline solid.
5. The melting point and boiling point increase with increase in molecular mass.
b. Chemical Properties
Combustibility: Hydrogen halides are neither combustible nor supporter of combustion.
Thermal stability: The stability of hydrogen halides decreases from HF to HI. Due to decrease in bond dissociation energies from HF to HI. It means HF is most stable and HI is least stable.
HF is most stable because of high electro-negativity of fluorine and HI is least stable because of least electro-negativity of iodine.
Stability decreases as in the order : HF > HCl > HBr > HI
Acidic Nature: Hydrogen halides are covalent molecule in dry state and don't turn blue litmus paper red. However, in aqueous solution they show acidic character. Hydrogen halides ionize to give proton as:
HX + H2O → H3O+ + X-
The acidic strength of hydrogen halides increases in the order
Acidic strength increases : HF < HCl < HBr < HI
The bond dissociation energies decrease from HF to Hl. So HI dissociate most easily and HF bond dissociate with maximum difficulty. Thus HI is strongest and HF is weakest among hydrohalic acid. Hydrohalic acid reacts with alkalis and metals.
HX + NaOH → NaX + H₂O
Zn + HX → ZnX2 + H2
Na + HX → NaX + H2
Reducing properties: The reducing property of Hydrogen halides increases from HF to HI. It is because of decrease in bond dissociation energy from HF to HI
HF < HCl < HBr < HI
(i) Reducing properties of HF HF is not reducing agent at all.
(ii) Reducing properties of HCI. HCI is weak reducing agent and reduces only strong oxidizing agent like MnO, KMnO HO, K Cr₂O, etc.
Note : HCI can't reduce CuSO, FeCl, and Conc. H₂SO:
iii) Reducing properties of HBr and HI: HBr and HI are strong reducing agent and easily oxidized to Bry and I, respectively. They can reduced acidified KMnO4, K2Cr2O7 solution, iodic acid to iodine, CuSO4 and FeCl3 solution, Nitric acid to nitrogen dioxide, Sulphuric acid to Sulphur dioxide.
Action with ammonia
Hydrogen halides give white fumes of ammonium halide with ammonia.
NH3 + HX → NH4X ( X = Cl, Br, I)
Action with halogens
Chlorine is liberated from HCl by fluorine
2HCl + F2 → 2HF + Cl2
Bromine is liberated from HBr by fluorine and Chlorine
2HBr + F2 → 2HF + Br2
2HBr + Cl2 → 2HCl + Br2
Iodine is liberated from HI by F2, Cl2 and Br2.
2HI + F2 → 2HF + I2
2HI + Cl2 → 2HCl + I2
2HI + Br2 → 2HBr + I2
Action with Silver Nitrate solution:
i) Hydrofluoric acid (HF) gives water soluble compound AgF, with AgNO3
HF + AgNO3 → AgF + HNO3
ii) Hydrochloric acid (HCl) gives white precipitate with AgNO3 solution which is insoluble in HNO3 but soluble in ammonia solution.
HCl + AgNO3 → AgCl + HNO3
iii) HBr gives a pale yellow precipitate with silver nitrate solution which is sparingly soluble in ammonia but insoluble in dilute HNO3
AgNO3 + HBr → AgBr + HNO3
iv) HI gives a yellow precipitate with silver nitrate which is insoluble both ammonia nitric acid.
AgNO3 + HI → AgI + HNO3
Action with lead acetate:
HCl and HBr give a white precipitate which are soluble in hot water.
2HX + (CH3COO)2Pb → PbX + 2CH3COOH
But HI gives yellow precipitate of lead iodide which soluble in boiling water.
(CH3COO)2Pb + HI → CH3COO + PbI2