Welcome to Group 17: The Halogens!
Hello future chemist! Get ready to dive into Group 17, one of the most exciting groups in the Periodic Table. These elements—Fluorine (F), Chlorine (Cl), Bromine (Br), and Iodine (I)—are known as the Halogens (meaning 'salt formers').
In these notes, we will explore why these elements are so aggressive, why their compounds behave differently as you go down the group, and how we use chlorine to keep our swimming pools clean! This chapter is essential for understanding trends and redox chemistry. Don't worry if it seems tricky; we’ll break down the patterns step-by-step.
1. The Halogen Elements (\(X_2\)) as Oxidising Agents
1.1 Understanding Reactivity: Halogens as Electron Thieves
Halogens are highly reactive non-metals because they are only one electron away from achieving a stable noble gas configuration. They desperately want to gain an electron.
When a substance gains electrons, it causes another substance to be oxidized. Therefore, the halogens act as powerful oxidising agents.
- Definition: An oxidising agent is an electron acceptor (it is itself reduced).
- The general reaction is: \[ X_2 + 2e^- \longrightarrow 2X^- \]
1.2 Trend in Oxidising Power (Reactivity)
The reactivity of the halogens decreases as you move down Group 17:
\(F_2 \gt Cl_2 \gt Br_2 \gt I_2\)
Why does reactivity decrease?
- Atomic Size: Going down the group, the atoms get larger.
- Shielding: The number of inner electron shells increases, leading to greater electron shielding.
- Nuclear Attraction: The incoming electron needs to be attracted by the positive nucleus. Because the nucleus is further away and more shielded in larger atoms (like Iodine), the attraction is weaker.
Analogy: Think of the halogens as electron magnets. Fluorine is a tiny, powerful magnet right next to the incoming electron. Iodine is a much larger magnet placed behind a thick wall (shielding); it struggles to pull the electron in.
1.3 Displacement Reactions
A more reactive halogen (higher up the group) can displace a less reactive halide ion (lower down the group) from its solution.
This is the classic experimental demonstration of the trend in oxidizing ability.
- Chlorine (\(Cl_2\)) with potassium bromide (\(KBr\)): Chlorine is more reactive than Bromine. \[ Cl_2(aq) + 2Br^-(aq) \longrightarrow 2Cl^-(aq) + Br_2(aq) \] (Chlorine oxidises the bromide ions to bromine).
- Bromine (\(Br_2\)) with potassium iodide (\(KI\)): Bromine is more reactive than Iodine. \[ Br_2(aq) + 2I^-(aq) \longrightarrow 2Br^-(aq) + I_2(aq) \] (Bromine oxidises the iodide ions to iodine).
Quick Review: Halogen Reactivity
Halogen reactivity (as oxidising agents) decreases down the group because the ability to attract an incoming electron decreases due to increased size and shielding.
2. Reactions with Hydrogen: Formation of Hydrogen Halides (HX)
2.1 Direct Reaction of \(X_2\) with Hydrogen (\(H_2\))
Halogens react with hydrogen gas to form the gaseous hydrogen halides: \[ H_2(g) + X_2(g) \longrightarrow 2HX(g) \]
The conditions required for this reaction demonstrate the trend in reactivity:
- Fluorine (\(F_2\)): Reacts explosively even in the dark and at very low temperatures. (Most reactive).
- Chlorine (\(Cl_2\)): Requires UV light or heat to react rapidly (often explosive).
- Bromine (\(Br_2\)): Requires moderate heat (e.g., passing vapors over hot platinum catalyst).
- Iodine (\(I_2\)): Requires continuous heating (high temperature) and the reaction is often incomplete (reversible). (Least reactive).
The trend confirms: Reactivity decreases \(F_2 \gt Cl_2 \gt Br_2 \gt I_2\).
3. Thermal Stability of Hydrogen Halides (HX)
3.1 The Trend in Stability
Once formed, how easy is it to break down the hydrogen halide (HX) molecule into its constituent elements (\(H_2\) and \(X_2\))? This is known as thermal stability.
The thermal stability of hydrogen halides decreases down the group:
\(HF \gt HCl \gt HBr \gt HI\)
3.2 Explanation using Bond Strength
Thermal stability is directly related to the strength of the H–X covalent bond. The weaker the bond, the easier it is to break apart using heat, and thus the less stable the compound is.
- Down the group: The size of the halogen atom (X) increases significantly.
- The overlap between the small H (1s) orbital and the increasingly large X orbital becomes less effective.
- A smaller overlap means the covalent bond is weaker, requiring less energy (less heat) to break it.
The result: HI has the weakest bond and is the easiest to decompose upon heating. HF has the strongest bond and is the most stable.
Test for Thermal Stability: When hydrogen halides are heated, we observe decomposition (especially of HBr and HI) into hydrogen and the halogen.
- HCl is very stable, decomposing only minimally at very high temperatures.
- HBr decomposes noticeably at moderate heat.
- HI decomposes readily upon gentle heating.
Key Takeaway for Stability
Stability (\(HF \longrightarrow HI\)) decreases because the H–X bond strength decreases due to increasing atomic size and reduced orbital overlap.
4. The Halide Ions (\(X^-\)) as Reducing Agents
Now let's look at the negative ions: Chloride (\(Cl^-\)), Bromide (\(Br^-\)), and Iodide (\(I^-\)). These ions are ready to lose an electron and return to the neutral halogen atom (\(X_2\)).
When a species loses electrons, it causes another substance to be reduced. Thus, halide ions act as reducing agents.
4.1 Trend in Reducing Power
The reducing power of the halide ions increases as you move down Group 17:
\(I^- \gt Br^- \gt Cl^- \gt F^-\)
Why does reducing power increase?
Reducing power depends on how easily the ion can give away its outermost electron.
- Ionic Size: Going down the group, the ionic radius increases (e.g., \(I^-\) is much larger than \(Cl^-\)).
- Nuclear Attraction: The outer electron in a large ion (like \(I^-\)) is further from the nucleus and experiences greater shielding.
- This outer electron is therefore held less strongly and is easier to remove (oxidize) by an external reagent.
Mnemonic: In the battle of $X^-$ vs. the nucleus, the electron is easier to escape from the Big Guy (Iodide) because the nucleus can’t hold onto it as tightly.
4.2 Reaction with Concentrated Sulfuric Acid (\(H_2SO_4\))
This reaction is crucial as it demonstrates the differing reducing powers of the halide ions. Concentrated sulfuric acid is itself a moderately powerful oxidising agent.
Step 1: Acid-Base Reaction (Occurs for all Halides)
All solid halide salts react with conc. \(H_2SO_4\) to produce the corresponding volatile hydrogen halide gas (HX). This is an acid-base reaction, not a redox reaction. \[ X^- + H_2SO_4 \longrightarrow HX + HSO_4^- \]
Step 2: Redox Reaction (The Test of Reducing Power)
After the initial acid-base reaction, the hydrogen halide (HX) produced must be strong enough to reduce the \(H_2SO_4\).
A. Chloride Ions (\(Cl^-\)) – The Weakling
- Chloride ions (\(Cl^-\)) are weak reducing agents. They are not strong enough to reduce conc. \(H_2SO_4\).
- Observation: Only dense white fumes of HCl gas are produced. No further redox reaction occurs.
\[ NaCl(s) + H_2SO_4(l) \longrightarrow HCl(g) + NaHSO_4(s) \]
B. Bromide Ions (\(Br^-\)) – The Moderate Reducer
- Bromide ions (\(Br^-\)) are moderately strong reducing agents. They reduce conc. \(H_2SO_4\) to \(SO_2\).
- Observation: Dense white fumes (HBr) plus brown fumes (Bromine, \(Br_2\)) and colourless, pungent \(SO_2\) gas.
- Redox Equation: \[ 2Br^- + 2H_2SO_4 \longrightarrow Br_2 + SO_2 + SO_4^{2-} + 2H_2O \] (Oxidation number of Br changes from -1 to 0; S changes from +6 to +4).
C. Iodide Ions (\(I^-\)) – The Strong Reducer
- Iodide ions (\(I^-\)) are strong reducing agents. They reduce conc. \(H_2SO_4\) multiple times, producing \(SO_2\), sulfur (S), and even hydrogen sulfide (\(H_2S\)).
- Observation: Dense white fumes (HI) plus purple vapour (Iodine, \(I_2\)) which solidifies to black, and often yellow solid (Sulfur) and a smell of rotten eggs (\(H_2S\)).
- Iodide Reducing \(H_2SO_4\) to \(SO_2\): \[ 2I^- + 2H_2SO_4 \longrightarrow I_2 + SO_2 + SO_4^{2-} + 2H_2O \]
- Iodide Reducing \(H_2SO_4\) to \(H_2S\) (strongest reduction): \[ 8I^- + 5H_2SO_4 \longrightarrow 4I_2 + H_2S + 4SO_4^{2-} + 4H_2O \] (Oxidation number of S changes from +6 to -2).
Struggling with Redox? Remember, the stronger the reducing agent (\(I^-\)), the further it can force the oxidation number of the oxidizing agent (\(S\) in \(H_2SO_4\)) down the scale (from +6 to +4, 0, or -2).
5. Qualitative Analysis of Halide Ions (Silver Nitrate Test)
We can identify chloride, bromide, and iodide ions in solution using aqueous silver nitrate, \(AgNO_3(aq)\), followed by aqueous ammonia, \(NH_3(aq)\).
5.1 Reaction with Silver Ions (\(Ag^+\))
Silver ions react with halide ions to produce precipitates of insoluble silver halides. \[ Ag^+(aq) + X^-(aq) \longrightarrow AgX(s) \]
-
Chloride (\(Cl^-\)): Forms Silver Chloride (AgCl).
Observation: White precipitate. -
Bromide (\(Br^-\)): Forms Silver Bromide (AgBr).
Observation: Cream/Off-white precipitate. -
Iodide (\(I^-\)): Forms Silver Iodide (AgI).
Observation: Pale Yellow precipitate.
Memory Aid: Down the Group: White, Cream, Yellow (W.C.Y.).
5.2 Testing Solubility in Aqueous Ammonia (\(NH_3(aq)\))
This step is vital because the white and cream colours are often hard to distinguish. Silver halide precipitates can sometimes dissolve in ammonia because the silver ion forms a soluble complex ion, \([Ag(NH_3)_2]^+\) (though you don't need to know the formula).
-
AgCl (White): Soluble in dilute aqueous ammonia.
The white precipitate dissolves, forming a clear solution. -
AgBr (Cream): Partially soluble (or only soluble in concentrated aqueous ammonia).
The cream precipitate remains undissolved in dilute \(NH_3\) but dissolves in concentrated \(NH_3\). -
AgI (Yellow): Insoluble in both dilute and concentrated aqueous ammonia.
The yellow precipitate remains.
Quick Review Box: Silver Nitrate Test
| Halide Ion | Precipitate Colour (\(AgNO_3\)) | Solubility in \(NH_3(aq)\) |
|---|---|---|
| \(Cl^-\) | White | Soluble (Dilute) |
| \(Br^-\) | Cream | Partially Soluble (Conc. only) |
| \(I^-\) | Pale Yellow | Insoluble |
6. The Special Case: Reactions of Chlorine (\(Cl_2\))
Chlorine exhibits a fascinating type of reaction called disproportionation, where the same element is simultaneously oxidised and reduced.
Chlorine is unusual because it can exist in several oxidation states, including 0 (\(Cl_2\)), -1 (\(Cl^-\)), +1 (in HOCl), and higher states up to +7.
6.1 Disproportionation with Aqueous Sodium Hydroxide (\(NaOH\))
A. Cold, Dilute Aqueous NaOH
Chlorine gas is bubbled into cold, dilute sodium hydroxide (alkali). \[ Cl_2(aq) + 2NaOH(aq) \longrightarrow NaCl(aq) + NaOCl(aq) + H_2O(l) \]
The products are sodium chloride (oxidation state -1) and sodium chlorate(I) (NaOCl, oxidation state +1). NaOCl is the active ingredient in household bleach.
-
Oxidation Changes:
- \(Cl_2 \longrightarrow Cl^-\): Reduction (0 to -1)
- \(Cl_2 \longrightarrow ClO^-\): Oxidation (0 to +1)
- Application: This reaction produces bleaching agents.
B. Hot, Concentrated Aqueous NaOH
If the reaction is carried out using hot, concentrated alkali, the disproportionation goes further. \[ 3Cl_2(aq) + 6NaOH(aq) \longrightarrow 5NaCl(aq) + NaClO_3(aq) + 3H_2O(l) \]
The products are sodium chloride (oxidation state -1) and sodium chlorate(V) (\(NaClO_3\), oxidation state +5).
-
Oxidation Changes:
- \(Cl_2 \longrightarrow Cl^-\): Reduction (0 to -1)
- \(Cl_2 \longrightarrow ClO_3^-\): Oxidation (0 to +5)
6.2 Chlorine in Water Purification
Chlorine is added to drinking water supplies and swimming pools to kill harmful bacteria and microorganisms. When chlorine dissolves in water, it undergoes a similar disproportionation reaction: \[ Cl_2(aq) + H_2O(l) \rightleftharpoons HCl(aq) + HOCl(aq) \]
- This equilibrium produces two species containing chlorine: Hydrochloric acid (HCl) and Hypochlorous acid (HOCl).
- HOCl is a weak acid and further dissociates: \[ HOCl(aq) \rightleftharpoons H^+(aq) + ClO^-(aq) \]
The Active Species:
The species responsible for killing bacteria are HOCl (hypochlorous acid) and the \(ClO^-\) ion (chlorate(I) or hypochlorite ion).
These two active species are powerful oxidising agents that damage the cellular structures of microorganisms, effectively sterilising the water.
Did you know? The effectiveness of chlorine in water purification is highly dependent on the pH. At very low pH, the amount of HOCl (the most effective species) decreases. If the pH is too high (alkaline), the concentration of \(ClO^-\) increases, but this ion is less powerful than HOCl, meaning more chlorine is needed to achieve the same disinfection. This is why pool pH is carefully monitored.