Redox Chemistry and Groups 1, 2 and 7

Welcome to the World of Redox and Groups 1, 2, and 7!

In this chapter, we are going to explore some of the most reactive and exciting elements on the Periodic Table. We will start by learning a new way to "keep books" on electrons through Redox Chemistry. Then, we will look at the personalities of the s-block metals (Groups 1 and 2) and the Halogens (Group 7). By the end of these notes, you’ll understand why certain metals explode in water and why chlorine is so good at cleaning swimming pools. Don't worry if it seems like a lot to remember—we’ll break it down step-by-step!

Section 8A: Redox Chemistry

Redox is short for Reduction-Oxidation. In chemistry, many reactions involve the movement of electrons from one place to another. Think of it like a game of catch: one atom throws an electron, and another one catches it.

1. Oxidation Numbers (The "Charge Score")

An oxidation number is a number assigned to an atom to show how many electrons it has lost or gained. It’s like a "score" for an atom's electron status.

The Golden Rules for Assigning Numbers:
- 1. Elements on their own (like \(Na\), \(Cl_2\), or \(O_2\)) are always 0.
- 2. Simple ions have the same number as their charge (e.g., \(Na^+\) is +1, \(Mg^{2+}\) is +2).
- 3. Fluorine is always -1 in compounds.
- 4. Hydrogen is usually +1 (except in metal hydrides like \(NaH\), where it is -1).
- 5. Oxygen is usually -2 (except in peroxides like \(H_2O_2\), where it is -1).
- 6. The sum of all numbers in a neutral compound must be 0.

Example: What is the oxidation number of Sulfur in \(H_2SO_4\)?
Hydrogen is \(+1 \times 2 = +2\). Oxygen is \(-2 \times 4 = -8\). To make the total zero, Sulfur must be +6 because \((+2) + (+6) + (-8) = 0\).

2. Defining Redox (OIL RIG)

To remember what happens to electrons, use the most famous mnemonic in chemistry:
OIL RIG
Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

Oxidising Agents: These are "electron thieves." They take electrons from others, so they get reduced themselves.
Reducing Agents: These are "electron givers." They give electrons away, so they get oxidised themselves.

3. Disproportionation

This is a fancy word for a specific situation where the same element in a single species is simultaneously oxidised and reduced. It’s like one atom of an element splitting its personality—half goes up in oxidation number, and half goes down.

Quick Review Box:
- Oxidation: Oxidation number increases.
- Reduction: Oxidation number decreases.
- Metals: Usually form positive ions (lose electrons/oxidised).
- Non-metals: Usually form negative ions (gain electrons/reduced).

Key Takeaway: Redox is all about electron transfer. Use oxidation numbers to track where the electrons are going. If the number goes up, it's oxidation!

Section 8B: The elements of Groups 1 and 2

Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals) are the "reactive extroverts" of the metal world. They really want to lose their outer electrons to become stable.

1. Trends in Reactivity and Ionisation Energy

As you go down Groups 1 and 2:
- First Ionisation Energy decreases: The outer electron gets further from the nucleus and is "shielded" by more inner shells, so it’s easier to pull away.
- Reactivity increases: Because it’s easier to lose that outer electron, the metals become more reactive as you go down the group. (Potassium is more reactive than Sodium!).

2. Reactions with Water, Oxygen, and Chlorine

- With Water: Group 1 metals react vigorously to form a metal hydroxide and hydrogen gas. (Example: \(2Na + 2H_2O \rightarrow 2NaOH + H_2\)). Group 2 metals are similar but generally slower (Magnesium reacts very slowly with cold water but quickly with steam).
- With Oxygen: They burn to form oxides. (Example: \(2Mg + O_2 \rightarrow 2MgO\)).
- With Chlorine: They react to form white chloride salts. (Example: \(Ca + Cl_2 \rightarrow CaCl_2\)).

3. Solubility Trends of Group 2

This is a common exam topic! Think of it as two opposite patterns:
- Hydroxides (\(OH^-\)): Become MORE soluble as you go down the group. (Barium hydroxide is very soluble; Magnesium hydroxide is not).
- Sulfates (\(SO_4^{2-}\)): Become LESS soluble as you go down the group. (Barium sulfate is famously insoluble—we use it in "Barium meals" for X-rays because it won't dissolve in the body).

4. Thermal Stability (The "Distortion" Story)

Why do some carbonates decompose (break down with heat) easier than others? It depends on the size and charge of the metal ion (cation).
Small, highly charged cations (like \(Mg^{2+}\)) are "polarising." They pull on the electron cloud of the carbonate or nitrate ion, making it easier for the molecule to break apart. As you go down the group, the cations get larger, their "pulling power" decreases, and the compounds become more thermally stable (they need more heat to break).

5. Flame Colours

When you heat these compounds, electrons jump to higher energy levels and then fall back down, releasing light. You need to know these "signatures":
- Lithium: Red/Crimson
- Sodium: Yellow/Orange
- Potassium: Lilac
- Calcium: Brick Red
- Strontium: Red
- Barium: Pale Green/Apple Green

Did you know? Fireworks get their beautiful colors from these metal salts! Barium creates the greens, and Strontium creates the deep reds.

Key Takeaway: Down Groups 1 and 2, reactivity increases while ionisation energy decreases. Group 2 sulfate solubility decreases, but hydroxide solubility increases.

Section 8C: Inorganic Chemistry of Group 7 (The Halogens)

The Halogens (Chlorine, Bromine, Iodine) are non-metals that exist as diatomic molecules (like \(Cl_2\)). They are "electron hunters" that want to gain one electron.

1. Physical Trends

As you go down Group 7:
- Melting and Boiling points increase: Molecules get larger, so they have more electrons, which leads to stronger London forces (temporary dipoles) between molecules.
- Appearance: Chlorine is a green gas; Bromine is a red-brown liquid; Iodine is a grey-black solid that turns into a purple vapor.
- Electronegativity decreases: The atoms get larger, so the nucleus has less "pull" on shared electrons.

2. Reactivity and Displacement

Unlike the metals, Halogens get LESS reactive as you go down the group. Chlorine is the "strongest" hunter, and Iodine is the "weakest."
A more reactive halogen will displace a less reactive halide ion from its compound.
Analogy: Imagine a stronger person (Chlorine) coming and taking the seat of a weaker person (Iodide ion) at a table.
\(Cl_2 + 2KI \rightarrow 2KCl + I_2\)
(The solution would turn brown because Iodine is formed).

3. Chlorine and Water Treatment

When chlorine is added to water, it undergoes a disproportionation reaction:
\(Cl_2 + H_2O \rightarrow HCl + HClO\)
The \(HClO\) (chloric(I) acid) kills bacteria, which is why we put chlorine in drinking water and pools. Even though chlorine is toxic, the benefits of clean water outweigh the risks!

4. Testing for Halide Ions (Silver Nitrate Test)

To find out if a mystery solution contains \(Cl^-\), \(Br^-\), or \(I^-\), we use Acidified Silver Nitrate (\(AgNO_3\)):
- 1. Add dilute Nitric Acid (to remove impurities).
- 2. Add Silver Nitrate solution.
- Chloride (\(Cl^-\)): White precipitate. (Dissolves in dilute ammonia).
- Bromide (\(Br^-\)): Cream precipitate. (Dissolves in concentrated ammonia).
- Iodide (\(I^-\)): Yellow precipitate. (Does not dissolve in ammonia).

Common Mistake: Don't forget the ammonia step! Cream and white can look very similar in a dark lab. The solubility in ammonia is the only way to be 100% sure.

5. Reaction with Concentrated Sulfuric Acid

This reaction shows the "reducing power" of the halides. As you go down the group, halide ions find it easier to lose an electron (they become better reducing agents).
- \(NaCl\) only produces \(HCl\) (steamy fumes).
- \(NaBr\) produces \(HBr\), but also reduces the sulfur to produce \(SO_2\) (orange Bromine fumes).
- \(NaI\) is so strong it reduces the sulfur all the way to \(H_2S\), which smells like rotten eggs!

Key Takeaway: Group 7 reactivity decreases down the group. Use Silver Nitrate followed by Ammonia to identify halide ions. Down the group, halides become better at giving away electrons (reducing agents).

Final Summary for Success

- Redox: Keep track of electrons using OIL RIG and oxidation numbers.
- Group 1 & 2: Reactivity increases down the group. Remember your flame colors and the sulfate/hydroxide solubility "X" pattern.
- Group 7: Reactivity decreases down the group. Displacement reactions happen when a "stronger" halogen kicks out a "weaker" one.

Don't worry if this seems tricky at first! Try practicing a few oxidation number calculations and drawing the "Displacement Table" for the Halogens. You've got this!