Chemistry | Redox Chemistry and Groups 1, 2 and 7

🌟 Introduction: Welcome to Redox and Group Chemistry! 🌟

Hello future Chemist! This chapter is incredibly important because it brings together fundamental concepts – electron movement (Redox) – and applies them to specific families of the Periodic Table: Groups 1, 2, and 7.

Don't worry if Redox seems tricky at first; we will break down the process of tracking electrons using simple rules. By the end of these notes, you’ll understand why calcium is more reactive than magnesium, and why chlorine is a stronger oxidising agent than iodine!

Key Goal: Mastering how electron transfer dictates chemical reactions and predicting the behaviour of key elements.

⚛️ SECTION 1: The Core of Redox Chemistry ⚛️

The term Redox is short for Reduction and Oxidation. A redox reaction is simply a reaction where electrons are transferred between species.

1.1 Defining Oxidation and Reduction

We define oxidation and reduction in three ways, but the electron definition is the most common at this level:

• Oxidation: Losing electrons; an increase in oxidation state.
• Reduction: Gaining electrons; a decrease in oxidation state.

🔥 Memory Aid: OIL RIG 🔥

This is the classic mnemonic you must know:
Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

1.2 Understanding Oxidation Numbers (States)

The oxidation number is a hypothetical charge an atom would have if the compound were purely ionic. It is essential for determining if a reaction is redox.

Rules for Assigning Oxidation Numbers:

1. An atom in its elemental state always has an oxidation number of zero (e.g., \( Na \), \( O_2 \), \( Cl_2 \)).
2. In a neutral compound, the sum of all oxidation numbers is zero.
3. In an ion, the sum of all oxidation numbers equals the charge of the ion (e.g., \( SO_4^{2-} \) totals -2).
4. Group 1 metals are always +1. Group 2 metals are always +2.
5. Hydrogen is usually +1 (except in metal hydrides, where it is -1).
6. Oxygen is usually -2 (except in peroxides, e.g., \( H_2O_2 \), where it is -1).
7. Fluorine is always -1. Other halogens are usually -1.

Example: Find the oxidation state of Sulfur (S) in sulfuric acid, \( H_2SO_4 \).
\( (2 \times H) + (1 \times S) + (4 \times O) = 0 \)
\( (2 \times +1) + S + (4 \times -2) = 0 \)
\( +2 + S - 8 = 0 \)
\( S = +6 \)
So, Sulfur's oxidation number is +6.

1.3 Oxidising and Reducing Agents

When one species is oxidized, it causes the other species to be reduced. We name the agents based on what they make the other molecule do:

• A Reducing Agent: It gets oxidized itself, but causes the reduction of another species. (It provides electrons).
• An Oxidising Agent: It gets reduced itself, but causes the oxidation of another species. (It accepts electrons).

💡 Analogy 💡

Think of a bully (the Oxidising Agent). The bully steals the electron (gets reduced), but by stealing it, they have forced the victim (the Reducing Agent) to lose the electron (get oxidised).

Key Takeaway for Section 1

Redox reactions involve electron transfer. We track this transfer using oxidation numbers. Reduction is gaining electrons (O.N. decreases); Oxidation is losing electrons (O.N. increases).

🛠️ SECTION 2: Group 2 – The Alkaline Earth Metals 🛠️

Group 2 elements (Be, Mg, Ca, Sr, Ba) are metals that typically form ions with a +2 charge. Their chemistry is dominated by their willingness to lose two outer electrons.

2.1 Trends in Reactivity

When moving down Group 2:

1. Atomic Radius Increases: We add new electron shells.
2. Nuclear Attraction Decreases: The outer electrons are further from the nucleus.
3. Shielding Increases: More inner shells block the nuclear charge.
4. First and Second Ionisation Energies Decrease: Less energy is required to remove the outer two electrons.

Conclusion on Reactivity: Because it takes less energy to remove the electrons as you move down the group, the elements become more powerful reducing agents (they are more easily oxidized) and are therefore more reactive.

2.2 Reactions with Water

Group 2 metals react with water to form the metal hydroxide and hydrogen gas. \[ M(s) + 2H_2O(l) \rightarrow M(OH)_2(aq) + H_2(g) \]

• Magnesium (Mg): Reacts very slowly with cold water. It reacts much faster with steam to form magnesium oxide, \( MgO \). \[ Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g) \]
• Calcium (Ca), Strontium (Sr), Barium (Ba): React readily with cold water. Reactivity increases down the group.

2.3 Solubility Trends (A Quick Note)

The solubility of Group 2 compounds shows opposite trends for hydroxides and sulfates:

• Hydroxides (e.g., \( Ca(OH)_2 \)): Solubility increases down the group.
• Sulfates (e.g., \( CaSO_4 \)): Solubility decreases down the group. (Barium sulfate, \( BaSO_4 \), is extremely insoluble and used in medical imaging.)

Key Takeaway for Section 2

Group 2 metals are reducing agents. Their reactivity increases down the group because it becomes easier to lose electrons due to increased atomic size and shielding.

🔬 SECTION 3: Group 7 – The Halogens 🔬

Group 7 elements (F, Cl, Br, I, At) are known as the Halogens (meaning 'salt formers'). They exist as diatomic molecules (e.g., \( Cl_2 \)) and are powerful oxidising agents, as they desperately want to gain one electron to achieve a full outer shell.

3.1 Physical Trends

As you move down Group 7:

• Boiling Point and Melting Point Increase: The atoms get larger, meaning the instantaneous dipoles get larger, increasing the strength of the van der Waals forces between molecules.
• Colour Deepens: \( Cl_2 \) (green-yellow gas), \( Br_2 \) (red-brown liquid), \( I_2 \) (grey solid/purple vapour).

3.2 Trends in Reactivity (Oxidising Power)

Halogens react by gaining one electron to form a halide ion, \( X^- \). \[ X_2 + 2e^- \rightarrow 2X^- \]

When moving down Group 7:

1. Atomic Radius Increases: The electron shell count increases.
2. Nuclear Attraction Decreases: The nucleus's pull on incoming electrons weakens.
3. Shielding Increases: More inner shells block the nuclear charge.

Conclusion on Reactivity: It becomes harder to attract and gain an external electron as you go down the group. Therefore, the halogens become less powerful oxidising agents, and their reactivity decreases down the group.

3.3 Halogen Displacement Reactions

A more reactive halogen (higher up the group) can oxidise the halide ion of a less reactive halogen (lower down the group).

Example: Chlorine gas (\( Cl_2 \)) is added to a solution containing bromide ions (\( Br^- \)).
Chlorine is more reactive than bromine, so it displaces the bromide ions: \[ Cl_2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br_2(aq) \]

• Observation: The colourless solution turns brown/orange (due to the formation of liquid bromine, \( Br_2 \)).
• Role of Chlorine: Oxidising Agent (it is reduced from O.N. 0 to -1).
• Role of Bromide: Reducing Agent (it is oxidized from O.N. -1 to 0).

3.4 Testing for Halide Ions (The Silver Nitrate Test)

This standard test confirms the presence of halide ions (\( Cl^-\), \( Br^-\), \( I^- \)) in solution. We add dilute nitric acid (to remove carbonate/hydroxide impurities) followed by aqueous silver nitrate, \( AgNO_3 \). \[ Ag^+(aq) + X^-(aq) \rightarrow AgX(s) \]

The colour of the precipitate identifies the halide:

• Chloride (\( Cl^- \)): White precipitate (\( AgCl \))
• Bromide (\( Br^- \)): Cream precipitate (\( AgBr \))
• Iodide (\( I^- \)): Yellow precipitate (\( AgI \))

Confirming the Halide using Ammonia Solution

We add aqueous ammonia to test the precipitate's solubility:

1. \( AgCl \) (White): Dissolves in dilute ammonia.
2. \( AgBr \) (Cream): Dissolves only in concentrated ammonia.
3. \( AgI \) (Yellow): Insoluble in both dilute and concentrated ammonia.

3.5 Disproportionation Reactions of Chlorine

A disproportionation reaction is a special type of redox reaction where the same element is simultaneously oxidised and reduced. This is often tested using chlorine.

A) Chlorine reacting with Water (Used for Sterilisation)

Chlorine is reduced (to \( Cl^- \)) and oxidized (to hypochlorous acid, \( HClO \)) \[ Cl_2(aq) + H_2O(l) \rightleftharpoons HClO(aq) + HCl(aq) \] In this reaction: Cl goes from O.N. 0 to -1 (in HCl) and O.N. 0 to +1 (in HClO).
The hypochlorous acid, \( HClO \), is the active ingredient used to kill bacteria.

B) Chlorine reacting with Cold, Dilute Aqueous Sodium Hydroxide (Alkali)

This reaction is used commercially to produce bleach (sodium chlorate(I)). \[ Cl_2(aq) + 2NaOH(aq) \rightarrow NaClO(aq) + NaCl(aq) + H_2O(l) \] The chlorine again disproportionates:

• Reduced to \( NaCl \) (O.N. -1)
• Oxidised to \( NaClO \) (Sodium Chlorate(I)) (O.N. +1)

Key Takeaway for Section 3

Group 7 halogens are strong oxidising agents, with reactivity decreasing down the group. They are identifiable by the colour and solubility of their silver precipitates, and chlorine can undergo disproportionation where it simultaneously gains and loses electrons.

🧠 Final Summary and Checklist 🧠

You have now mastered the fundamentals of Redox and applied these principles to two critical groups! Use this checklist to review:

• Can I assign oxidation numbers correctly?
• Can I define oxidation/reduction using electrons and oxidation numbers (OIL RIG)?
• Can I explain why Group 2 reactivity increases down the group (easier to lose electrons)?
• Can I explain why Group 7 reactivity decreases down the group (harder to gain electrons)?
• Can I predict the outcome of a halogen displacement reaction?
• Do I know the colours and ammonia solubility of the silver halide precipitates?
• Can I write the equation for the disproportionation of chlorine in alkali?

Well done! Keep practicing those oxidation number calculations—they are the key to unlocking all redox problems.