Welcome to Chemical Changes!

In this chapter, we are going to explore how substances transform into new ones. We will dive into the world of acids and alkalis, learn how to make salts (not just the kind you put on chips!), and discover how electricity can be used to split compounds apart in a process called electrolysis. Don't worry if it sounds like a lot—we will break it down into simple, easy-to-follow steps.

1. Acids, Alkalis, and the pH Scale

At the heart of this topic are two types of chemicals: acids and alkalis. To understand them, we look at the ions they release when dissolved in water.

What makes an Acid or an Alkali?

  • Acids: When acids dissolve in water, they produce hydrogen ions \( (H^+) \).
  • Alkalis: When alkalis dissolve in water, they produce hydroxide ions \( (OH^-) \).

The pH Scale

The pH scale measures how acidic or alkaline a solution is, usually from 0 to 14.

  • pH 0–6: Acidic (lower pH = stronger acid).
  • pH 7: Neutral (like pure water).
  • pH 8–14: Alkaline (higher pH = stronger alkali).

Indicators

Indicators are special dyes that change colour depending on the pH. You need to know these three:

  • Litmus: Red in acid, Blue in alkali.
  • Methyl Orange: Red in acid, Yellow in alkali.
  • Phenolphthalein: Colourless in acid, Pink in alkali.

Quick Review Box:
Higher concentration of \( H^+ \) = Lower pH (More Acidic)
Higher concentration of \( OH^- \) = Higher pH (More Alkaline)

The "Factor of 10" Rule: If the pH of a solution decreases by 1 (e.g., from pH 4 to pH 3), the concentration of hydrogen ions has actually increased by 10 times! If it moves 2 pH units, it's a 100-fold change (\( 10 \times 10 \)).

2. Strong vs. Weak and Dilute vs. Concentrated

Students often get these mixed up, but they mean very different things!

Strong vs. Weak (The "Splitting" factor)

This describes how well the acid molecules split into ions (dissociate) in water.

  • Strong Acids: Fully dissociate. Every single molecule breaks apart to release \( H^+ \) ions. Example: Hydrochloric acid \( (HCl) \).
  • Weak Acids: Only partially dissociate. Only a few molecules break apart; most stay stuck together. Example: Ethanoic acid (vinegar).

Dilute vs. Concentrated (The "Crowd" factor)

This describes how much of the acid is actually in the water.

  • Concentrated: A lot of acid particles in a small volume of water.
  • Dilute: A small amount of acid particles in a large volume of water.

Analogy: Think of tea. A "Strong" tea is like a strong acid (lots of flavour released). A "Concentrated" tea is like putting five tea bags in one tiny cup!

3. Reactions of Acids

Acids react with different substances to form salts. A salt is just a compound where the hydrogen in an acid has been replaced by a metal ion.

The General Equations

  1. Acid + Metal \(\rightarrow\) Salt + Hydrogen
  2. Acid + Metal Oxide \(\rightarrow\) Salt + Water
  3. Acid + Metal Hydroxide \(\rightarrow\) Salt + Water
  4. Acid + Metal Carbonate \(\rightarrow\) Salt + Water + Carbon Dioxide

Chemical Tests for Gases

When these reactions happen, you might see bubbles. Here is how to check which gas is which:

  • Hydrogen (\( H_2 \)): Use a lit splint. It will make a "squeaky pop" sound.
  • Carbon Dioxide (\( CO_2 \)): Bubble the gas through limewater. The limewater will turn cloudy/milky.

Neutralisation

When an acid reacts with a base (like a metal oxide or hydroxide), they neutralise each other. If it’s an acid reacting with an alkali, the ionic equation is always:
\( H^+(aq) + OH^-(aq) \rightarrow H_2O(l) \)

Key Takeaway: Acids react to form salts. The name of the salt depends on the acid: Hydrochloric acid makes Chlorides, Sulfuric acid makes Sulfates, and Nitric acid makes Nitrates.

4. Making Salts (Practical Chemistry)

How we make a salt depends on whether the starting materials dissolve in water.

Method A: Using an Insoluble Base (e.g., Making Copper Sulfate)

If you are reacting an acid with something that doesn't dissolve (like Copper Oxide powder):

  1. Add the solid to the acid until no more dissolves (this is called adding it in excess to ensure all acid is used up).
  2. Filter the mixture to remove the leftover unreacted powder.
  3. Evaporate the water from the remaining solution using a water bath until crystals start to form.
  4. Leave it to cool and crystallise.

Method B: Using a Soluble Base (Titration)

If both the acid and the alkali are liquids, you can't just "see" when the reaction is finished. You must use Titration.

  • Use a pipette to measure a fixed volume of alkali into a flask.
  • Add an indicator.
  • Use a burette to slowly add acid until the indicator changes colour (the end point).
  • Note the volume, then repeat without indicator to get a pure salt solution.

Predicting Solubility (The Rules)

Sometimes reactions form a solid that won't dissolve—this is called a precipitate. You need to know these rules:

  • Always Soluble: All Sodium, Potassium, Ammonium salts, and all Nitrates.
  • Mostly Soluble: Chlorides (except Silver and Lead) and Sulfates (except Lead, Barium, and Calcium).
  • Mostly Insoluble: Carbonates and Hydroxides (except Sodium, Potassium, and Ammonium).

5. Electrolysis

Electrolysis is using electricity to split an ionic compound into its elements. This only works if the compound is molten (melted) or dissolved in water, so the ions are free to move.

The Setup

We use two rods called electrodes:

  • Anode: The Positive electrode.
  • Cathode: The Negative electrode.

Memory Aid: Use the word PANIC (Positive Anode, Negative Is Cathode).

How Ions Move

Opposites attract!

  • Cations (Positive ions) move to the Cathode (Negative).
  • Anions (Negative ions) move to the Anode (Positive).

Oxidation and Reduction (OIL RIG)

  • Oxidation is Loss of electrons. (Happens at the Anode).
  • Reduction is Gain of electrons. (Happens at the Cathode).

Half Equations

These show what happens to electrons at the electrodes. For example, in the electrolysis of molten Lead Bromide \( (PbBr_2) \):

  • At the Cathode: \( Pb^{2+} + 2e^- \rightarrow Pb \) (Lead ions gain electrons - Reduction).
  • At the Anode: \( 2Br^- \rightarrow Br_2 + 2e^- \) (Bromide ions lose electrons - Oxidation).

Did you know? Electrolysis is used to purify Copper. Impure copper is used as the anode, and pure copper grows on the cathode as ions move through a copper sulfate solution!

Quick Review Box:
1. Electrolyte = the liquid being split.
2. Cations go to Cathode.
3. Anions go to Anode.
4. OIL RIG (Oxidation Is Loss, Reduction Is Gain).

Don't worry if half equations seem tricky! Just remember that the goal is to turn the charged ion back into a neutral atom by either adding or taking away electrons.