Welcome to Chemical Changes!
Hello! In this chapter, we are going to explore how substances transform into new ones. We’ll look at why some metals are more "energetic" than others, how to make your own crystals (salts), and how we use electricity to "split" chemicals apart. Understanding these changes is the secret to how we get metals like aluminum for soda cans and how we make medicines. Don’t worry if some of this seems like a lot to take in at first—we’ll break it down into simple, bite-sized pieces!
5.4.1 Reactivity of Metals
Metal Oxides
When metals react with oxygen, they create metal oxides.
Example: When iron rusts, it is reacting with oxygen to form iron oxide.
- Oxidation is when a substance gains oxygen.
- Reduction is when a substance loses oxygen.
The Reactivity Series
Some metals are very reactive (like potassium), and others are very "lazy" or unreactive (like gold). We arrange them in a list called the reactivity series based on how easily they form positive ions.
Memory Aid (Mnemonic):
Please Stop Lions Calling Me A Careless Zebra, Instead Try Learning How Copper Saves Gold.
(Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, Carbon, Zinc, Iron, Tin, Lead, Hydrogen, Copper, Silver, Gold)
Displacement Reactions: A more reactive metal is like a stronger person—it can "push out" (displace) a less reactive metal from its compound.
Analogy: Imagine a more popular kid at a lunch table taking the seat of a less popular kid.
Extracting Metals
How we get metal out of the ground depends on where it sits in the reactivity series:
1. Unreactive metals (Gold): Found as pure metal in the Earth. Just dig it up!
2. Metals less reactive than Carbon: Can be extracted by heating them with carbon (Reduction). The carbon "steals" the oxygen.
3. Metals more reactive than Carbon: Carbon isn't strong enough to steal the oxygen, so we must use electrolysis.
HT ONLY: Oxidation and Reduction in terms of Electrons
Quick Review: In the Higher Tier, we define Redox reactions by electron transfer.Use the mnemonic OIL RIG:
Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)
Key Takeaway: Metals form positive ions. Their place in the reactivity series tells us how they react and how we extract them from ores.
5.4.2 Reactions of Acids
Acids + Metals
When an acid reacts with a metal, it produces a salt and hydrogen gas.
\( \text{metal} + \text{acid} \rightarrow \text{salt} + \text{hydrogen} \)
Quick Review: You can test for hydrogen using a "squeaky pop" test with a lit splint!
Neutralisation
Acids can be cancelled out (neutralised) by alkalis (soluble bases) or bases (insoluble).
1. \( \text{acid} + \text{metal hydroxide} \rightarrow \text{salt} + \text{water} \)
2. \( \text{acid} + \text{metal oxide} \rightarrow \text{salt} + \text{water} \)
3. \( \text{acid} + \text{metal carbonate} \rightarrow \text{salt} + \text{water} + \text{carbon dioxide} \) (this one fizzes!)
The pH Scale
The pH scale (0-14) measures how acidic or alkaline a solution is.
- pH 0-6: Acidic (contains \( H^+ \) ions)
- pH 7: Neutral (pure water)
- pH 8-14: Alkaline (contains \( OH^- \) ions)
Did you know? A pH probe is more accurate than universal indicator paper because it gives a numerical value instead of just a color change.
Required Practical 8: Making a Soluble Salt
To make pure crystals of a salt (like Copper Sulfate):
1. Heat the acid gently.
2. Add the metal oxide/carbonate in excess (keep adding until it stops reacting) to make sure all the acid is used up.
3. Filter the mixture to remove the leftover solid.
4. Evaporate the water from the solution to leave behind crystals.
HT ONLY: Strong vs Weak Acids
- Strong Acids: Fully ionise in water (e.g., Hydrochloric, Sulfuric, Nitric acids). Every molecule splits up to release \( H^+ \).- Weak Acids: Only partially ionise in water (e.g., Ethanoic, Citric acids).
Fact: As the pH decreases by 1 unit, the concentration of \( H^+ \) ions increases by 10 times! (e.g., pH 1 has 10x more \( H^+ \) than pH 2).
Key Takeaway: Acids react to form salts. The type of salt depends on the acid (Hydrochloric makes Chlorides; Sulfuric makes Sulfates; Nitric makes Nitrates).
5.4.3 Electrolysis
The Process
Electrolysis means "splitting with electricity." We pass an electric current through a liquid called an electrolyte to break compounds down into elements.
Key Terms:
- Cathode: The negative electrode. Positive ions go here.
- Anode: The positive electrode. Negative ions go here.
Memory Aid: PANIC (Positive Anode, Negative Is Cathode).
Electrolysis of Molten Compounds
When an ionic solid is melted (molten), the ions are free to move.
Example: If you electrolyse molten Lead Bromide, you get Lead at the cathode and Bromine at the anode.
Extracting Aluminium
Aluminium is too reactive to be extracted with carbon. Instead, we use electrolysis.
- Cryolite: We mix aluminium oxide with cryolite to lower the melting point, which saves lots of money and energy!
- The Anode Problem: The positive anodes are made of carbon. They react with oxygen to form \( CO_2 \), so they gradually "burn away" and must be replaced regularly.
Electrolysis of Aqueous Solutions (Dissolved in Water)
When you dissolve a salt in water, you also have \( H^+ \) and \( OH^- \) ions from the water. This makes it a bit more complicated. Only one ion is discharged at each electrode.
The Rules:
1. At the Cathode (-): Hydrogen is produced IF the metal is more reactive than hydrogen.
2. At the Anode (+): Oxygen is produced UNLESS the solution contains halide ions (Chloride, Bromide, Iodide), in which case the Halogen is produced.
HT ONLY: Half Equations
These show what happens at each electrode.- At the Cathode (Reduction): \( 2H^+ + 2e^- \rightarrow H_2 \)
- At the Anode (Oxidation): \( 4OH^- \rightarrow O_2 + 2H_2O + 4e^- \) or \( 2Cl^- \rightarrow Cl_2 + 2e^- \)
Common Mistake: Don't confuse the names! Remember Anions (negative) go to the Anode, and Cations (positive) go to the Cathode.
Key Takeaway: Electrolysis is used to extract reactive metals or separate dissolved compounds. The product depends on the reactivity of the ions involved.