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AS Level Chemistry 9701
11. Group 17
Written by: Pranav I
Formatted by: Pranav I
Index
12.1 Physical properties of Group 17 elements
- Referred to as halogens
- Their atoms have seven electrons in their outermost principal quantum shell
- All are non-metals
- Exist as diatomic molecules at room temperature (single covalent bond between the atoms)
- Melting points and boiling points increase going down the group (volatility decreases down the group)
- Weak instantaneous dipole-induced dipole forces hold molecules together
- Become stronger as the number of electrons increases
- Colors of halogens becomes darker down the group
- Fluorine (exists as a gas) → pale yellow
- Chlorine (exists as a gas) → green
- Bromine (exists as a liquid) → brown
- Iodine (exists as a solid) → grey solid and purple vapour
12.2 Reactions of Group 17 elements
- React with metallic elements
- Need to gain 1 more electron to achieve a stable electronic configuration
- Each atom gains one electron from atoms of the metal to form halide ions with a ‘1-‘ charge
Ca(s) + Cl2(g) → Ca2+(Cl–)2 - Halogens are oxidising agents
- Electron acceptors
- They are themselves reduced (oxidation number reduces from 0 to -1)
- Also react with non-metals
- Share a pair of electrons with another non-metal atom in a covalent bond
H2(g) + Cl2(g) → 2HCl(g)
🚨 The halogens get less reactive going down Group 17
- Electronegativity is the ability of a covalently bonded atom to attract the bonding pair of electrons towards itself
- Electronegativity decreases down Group 17
- Fluorine is a stronger oxidizing agent (more reactive) than all the other halogens
- More electronegative
- Lesser shielding of outer electrons due to lesser number of shells → greater nuclear attraction
Evidence of the oxidizing power of the halogens
- Relative reactivity of the halogens can be judged using their displacement reactions with other halide ions in solution
🚨 A more reactive halogen can displace a less reactive halogen from a halide solution of the less reactive halogen
\[
\text{Cl}_2(aq) + 2\text{NaBr}(aq) \rightarrow 2\text{NaCl}(aq) + \text{Br}_2(aq)
\]
- Chlorine atoms are more electronegative than bromine atoms (stronger oxidizing agents)
- The oxidizing power of the halogens (their ability to gain electrons) decreases down the group
- Colours of halogen solutions are difficult to identify positively
- Dissolve well in cyclohexane producing clearly different colours
- Add this organic solvent after mixing the halogen / halide solutions
- Shake the mixture and allow it to settle
- Identify the halogen present based on the colour produced
Reactions of halogens with hydrogen
- Form hydrogen halides when combined with hydrogen gas
- The thermal stability of the hydrogen halides decreases down the group
- Hydrogen iodide decomposes when a red-hot wire is inserted into a sample of it
- 10% of a hydrogen bromide sample decomposes at 430 °C
- Hydrogen chloride and hydrogen fluoride are not decomposed in temperatures up to 1500 °C
- The H-X bond energy decreases down the group
- Larger bond length down the group
- Due to the increase in the size of the halogen atoms
12.3 Reactions of the halide ions
Testing for halide ions
- Dissolve the unknown compound in dilute nitric acid and add silver nitrate solution
| # | Halide Ion | Colour of Silver Halide Precipitate on Addition of Silver Nitrate Solution | Effect on Precipitate of Adding Dilute Ammonia Solution | Effect on Precipitate of Adding Concentrated Ammonia Solution |
|---|---|---|---|---|
| 1 | Chloride | White | Dissolves | Dissolves |
| 2 | Bromide | Cream | Remains insoluble | Dissolves |
| 3 | Iodide | Pale Yellow | Remains insoluble | Remains insoluble |
Reactions of halide ions with concentrated sulfuric acid
- Cl–, Br– and I– will react with concentrated sulfuric acid
- All of these reactions produce one or more toxic gases → must be performed in a fume cupboard
- NaCl(s) + H2SO4(l) → NaHSO4(s) + HCl(g)
- Observation: white fumes
- Since HBr and HI decompose easily, we cannot use the above reaction to prepare samples of these gases
- When hydrogen halides decompose into their elements, the halide ions are oxidized
- Concentrated sulfuric acid is a relatively strong oxidizing agent
- Not strong enough to oxidize HCl
- Will oxidize and decompose HBr and HI
- With sodium bromide
- NaBr(s) + H2SO4(l) → NaHSO4(s) + HBr(g)
- 2HBr(g) + H2SO4(l) → Br2(g) + SO2(g) + 2H2O(l)
- Observation: reddish-brown gas
- With sodium iodide → sulfuric acid is reduced to several sulfur products
- NaI(s) + H2SO4(l) → NaHSO4(s) + HI(g)
- 2HI(g) + H2SO4(l) → I2(g) + SO2(g) + 2H2O(l)
- 6HI(g) + H2SO4(l) → 3I2(g) + S(s) + 4H2O(l)
- Observation: yellow solid (sulfur)
- 8HI(g) + H2SO4(l) → 4I2(g) + H2S(g) + 4H2O(l)
- Observation: strong smell of bad eggs (hydrogen sulfide)
- Observation: purple vapor (iodine)
- The Br– and I– ions act as reducing agents (electron donors)
- Reduce sulfuric acid
- Are themselves oxidized
- The halide ions are increasingly effective as reducing agents down the group
12.4 Disproportionation reactions
- When Cl2 reacts with dilute alkali, some Cl atoms are reduced and some are oxidized in the same reaction (disproportionation reaction)
- The actual reaction depends on the conditions (temperature)
Chlorine in cold alkali (15 °C)
\[
\text{Cl}_2(aq) + 2\text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{NaClO}(aq) + \text{H}_2\text{O}(l)
\]
- The oxidation number of chlorine increases from 0 to +1
- Sodium chlorate(I) is formed
Chloride in hot alkali (70 °C)
\[
3\text{Cl}_2(aq) + 6\text{NaOH}(aq) \rightarrow 5\text{NaCl}(aq) + \text{NaClO}_3(aq) + 3\text{H}_2\text{O}(l)
\]
- The oxidation number of chlorine increases from 0 to +5
- Sodium chlorate(V) is formed
Chlorination of water
- Adding a small amount of chlorine to a water supply will kill bacteria and make it safe for drinking
- Chlorine undergoes a disproportionation reaction in water
- HClO is called chloric(I) acid
- Sterilizes water by killing bacteria
- Some of it dissociates to form ClO–(aq) → also a sterilizing agent
