GCSE AQA Chemical Changes Revision Notes

Metal Oxides and the Reactivity Series

Ever wondered why some metals rust quickly whilst others stay shiny for years? It all comes down to how eagerly they react with oxygen and other substances.

When metals react with oxygen, they form metal oxides through oxidation reactions. Remember that oxidation means gaining oxygen, whilst reduction means losing oxygen. The speed and intensity of these reactions depends on where the metal sits in the reactivity series.

The reactivity series ranks metals from most to least reactive: Potassium > Sodium > Lithium > Calcium > Magnesium > Zinc > Iron > Copper. More reactive metals lose electrons more easily to form positive ions, which is why potassium explodes in water whilst copper barely reacts with anything!

You can work out this order by watching how metals react with water and dilute acids. Potassium fizzes violently and may even catch fire with a lilac flame, whilst copper just sits there doing nothing. The faster the bubbles form and the more heat produced, the more reactive the metal.

Key Tip: A more reactive metal can always displace a less reactive metal from its compound - this is how we can predict which reactions will actually happen.

Reactions with Water and Acids

Watching metals react with water and acids is like seeing a fireworks display - some metals go mental, others barely whisper!

With water at room temperature, the most reactive metals put on quite a show. Potassium and sodium float on the surface, fizz madly, and produce hydrogen gas plus metal hydroxides. Lithium reacts more steadily, calcium goes slowly, and magnesium barely bothers. Zinc, iron, and copper? They completely ignore cold water.

Dilute acids like hydrochloric acid get more metals involved in the action. Magnesium reacts rapidly with fizzing and heat, zinc goes at a moderate pace, iron takes its time, but copper still refuses to join the party. That's because copper sits below hydrogen in the reactivity series.

The secret behind all this excitement is electron transfer. Reactive metals desperately want to lose their outer electrons to form positive ions. Potassium, sodium, and lithium only have one outer electron, so they chuck it away easily. Copper holds onto its electrons tightly, making it stubbornly unreactive.

Quick Check: If you see violent fizzing and lots of heat, you're dealing with a highly reactive metal that's losing electrons like there's no tomorrow!

Metal Extraction and Electron Transfer

Getting pure metals from rocks isn't as simple as just digging them up - most metals are locked away in compounds that need some serious chemistry to break apart.

Unreactive metals like gold are found naturally as pure metal (which is why it's so valuable!), but most metals exist as compounds in the Earth's crust. We can extract metals that are less reactive than carbon by heating their oxides with carbon. This process is called reduction because the metal oxide loses oxygen.

Understanding electron transfer is crucial here, and there's a handy memory trick: OIL RIG (Oxidation Is Loss, Reduction Is Gain - of electrons). When zinc displaces copper from copper sulfate, zinc loses electrons (gets oxidised) whilst copper ions gain electrons (get reduced).

Half equations show exactly what happens to electrons. For example: Zn$s$ → Zn²⁺(aq) + 2e⁻ shows zinc losing two electrons (oxidation), whilst Cu²⁺(aq) + 2e⁻ → Cu$s$ shows copper ions gaining two electrons (reduction). These equations must balance - electrons lost always equal electrons gained.

Remember: In any reaction, if one substance gets oxidised (loses electrons), another must get reduced (gain electrons) - they're like dance partners that never separate!

Acid-Metal Reactions and Redox

When acids meet reactive metals, it's like a chemical conversation where electrons change hands and new substances are born.

The general pattern is simple: Metal + Acid → Salt + Hydrogen gas. Whether you use hydrochloric acid or sulfuric acid, you'll always get hydrogen bubbling off, but the salt will be different (chlorides from hydrochloric acid, sulfates from sulfuric acid).

Let's break down what's actually happening with electron transfer. Take magnesium reacting with hydrochloric acid: Mg$s$ + 2HCl(aq) → MgCl₂(aq) + H₂$g$. The magnesium atoms lose electrons to become Mg²⁺ ions (oxidation), whilst H⁺ ions from the acid gain electrons to form hydrogen gas (reduction).

This pattern repeats for zinc and iron reactions too. The metal always gets oxidised (loses electrons), and the hydrogen ions always get reduced (gain electrons). It's like a predictable chemical dance where metals give up electrons and hydrogen ions grab them.

The key insight? Only metals above hydrogen in the reactivity series will react with acids to produce hydrogen gas. That's why copper won't react - it's too stubborn to give up its electrons to hydrogen ions.

Test Tip: Remember that in acid-metal reactions, the metal always loses electrons (oxidation) and hydrogen ions always gain electrons (reduction) - no exceptions!

Neutralisation and Salt Formation

Neutralisation reactions are like chemical peacekeepers - they take aggressive acids and feisty alkalis and help them settle their differences to make useful salts.

When acids meet alkalis or bases, they produce salts and water through neutralisation. The specific salt depends on two things: which acid you use (hydrochloric acid makes chlorides, nitric acid makes nitrates, sulfuric acid makes sulfates) and what positive ions are in your base or alkali.

Metal carbonates add an extra twist - they react with acids to produce salts, water, AND carbon dioxide gas. So the reaction 2HCl(aq) + Na₂CO₃(aq) → 2NaCl(aq) + H₂O$l$ + CO₂$g$ gives you fizzing as the CO₂ escapes.

Predicting salt formulae is straightforward once you know the pattern. Identify the negative ion from your acid (Cl⁻ from HCl, SO₄²⁻ from H₂SO₄) and the positive ion from your base (Na⁺ from NaOH, Ca²⁺ from Ca(OH)₂), then combine them to balance the charges.

Pro Tip: If you see fizzing when acid meets a white powder, you've probably got a carbonate releasing CO₂ gas - a dead giveaway for this type of reaction!

Making Salts and Understanding pH

Creating pure, dry salts in the lab is like following a recipe - get the steps right and you'll have perfect crystals every time.

To make soluble salts, react your acid with a solid base, metal oxide, or carbonate until no more dissolves. Filter off the excess solid, then evaporate the water to leave behind your salt solution. Cool it down to form crystals, filter them out, and dry them - job done!

The pH scale from 0 to 14 tells you how acidic or alkaline a solution is. pH 7 is neutral (like pure water), below 7 is acidic, and above 7 is alkaline. Universal indicator gives you a colour code: red for strong acids $pH 1-3$, orange for weak acids $pH 4-6$, green for neutral (pH 7), blue for weak alkalis $pH 8-11$, and purple for strong alkalis $pH 12-14$.

Here's the crucial bit: acids produce H⁺ ions in water, whilst alkalis contain OH⁻ ions. During neutralisation, these ions react together: H⁺(aq) + OH⁻(aq) → H₂O$l$. It's like they're made for each other!

As pH drops by one unit, the hydrogen ion concentration increases by a factor of 10. So pH 2 has ten times more H⁺ ions than pH 3.

Lab Safety: Always add acid to water, never the other way around - "Do as you oughta, add acid to water" could save you from nasty splashes!

Titrations - Measuring Reactions Precisely

Titrations are like chemical detective work - you're finding out exactly how much acid and alkali react together by watching for that crucial colour change.

The basic setup involves a burette filled with alkali, a pipette to measure acid into a conical flask, and a few drops of indicator to signal when you've hit the endpoint. Place everything on a white tile so you can spot colour changes easily.

Phenolphthalein changes from colourless (acidic) to pale pink (neutral), whilst methyl orange switches from red (acidic) to yellow (alkaline). The trick is stopping the instant you see a permanent colour change - that's your endpoint where acid and alkali have completely neutralised each other.

During the titration, swirl the flask continuously and add alkali drop by drop near the end. Record your initial and final burette readings to calculate exactly how much alkali was needed. This precision lets you work out unknown concentrations using the balanced equation.

The process might seem fiddly, but it's incredibly useful for quality control in industry and for making sure medicines contain exactly the right amount of active ingredient.

Technique Tip: Near the endpoint, add the alkali drop by drop whilst swirling constantly - rushing here will overshoot your target and ruin your results!

Calculations and Acid Strength

Turning titration results into useful numbers might look scary, but it's just organised thinking with a clear method to follow.

Start with your balanced equation, then use the formula: moles = concentration × volume (remembering to convert mL to dm³ by dividing by 1000). Use the stoichiometry from your equation to work out moles of the unknown solution, then calculate its concentration using: concentration = moles ÷ volume.

Understanding strong vs weak acids is crucial for predictions. Strong acids like hydrochloric, nitric, and sulfuric acids completely ionise in water, releasing all their H⁺ ions. Weak acids like ethanoic, citric, and carbonic acids only partially ionise, keeping most of their H⁺ ions locked up.

This explains why equal concentrations of strong and weak acids have different pH values - the strong acid releases more H⁺ ions, making it more acidic. Don't confuse strength with concentration though! You can have a dilute strong acid or a concentrated weak acid.

The pH scale is logarithmic, so each unit represents a 10-fold change in H⁺ concentration. pH 1 has 10 times more H⁺ ions than pH 2, which has 10 times more than pH 3, and so on.

Calculation Hack: Always convert mL to dm³ first (÷1000), use the balanced equation to find the mole ratio, then work systematically through each step - no shortcuts!

Electrolysis - Splitting Compounds with Electricity

Electrolysis is like using electrical energy as a crowbar to split ionic compounds apart and force chemical reactions that wouldn't normally happen.

When ionic compounds melt or dissolve in water, their ions become free to move around, creating electrolytes that conduct electricity. Pass a current through them, and the ions migrate to oppositely charged electrodes - remember PANIC: Positive ions to Anode, Negative ions to Cathode (wait, that's backwards - Positive ions actually go to the Cathode!).

At the cathode (negative electrode), reduction happens as positive ions gain electrons. For example: Cu²⁺(aq) + 2e⁻ → Cu$s$. At the anode (positive electrode), oxidation occurs as negative ions lose electrons: 2Cl⁻(aq) → Cl₂$g$ + 2e⁻.

Half equations must balance in terms of both atoms and charge. Add electrons to balance charge, water molecules to balance oxygen atoms, and H⁺ or OH⁻ ions to balance hydrogen atoms if needed.

Molten ionic compounds give predictable results - metals form at the cathode, non-metals at the anode. Lead bromide produces lead metal and bromine gas, whilst sodium chloride gives sodium metal and chlorine gas.

Memory Aid: Think of electrolysis as an electrical tug-of-war where positive and negative ions get pulled to opposite corners and transformed into new substances!

Metal Extraction by Electrolysis

Sometimes carbon just isn't strong enough to extract metals from their compounds - that's when we bring out the big guns of electrolysis.

Electrolysis for metal extraction is essential when metals are too reactive to be reduced by carbon, or when the metal would react with carbon itself. Reactive metals like sodium, magnesium, and aluminium all need the electrical treatment because they hold onto oxygen more stubbornly than carbon can pull it away.

The process involves melting the metal compound to free up the ions, then passing electricity through to force the metal ions to gain electrons at the cathode. Meanwhile, non-metal ions lose electrons at the anode, often forming gases that can be captured or safely released.

Predicting electrolysis products follows clear patterns. For binary ionic compounds (just two elements), the metal always appears at the cathode, and the non-metal at the anode. Pb²⁺ ions become lead metal, whilst Br⁻ ions become bromine gas.

The downside? Massive energy costs for melting compounds and generating electrical current make this extraction method expensive. That's why we only use electrolysis when absolutely necessary, and why metals like aluminium were once more valuable than gold!

Industrial Reality: The huge energy requirements for electrolytic metal extraction explain why aluminium recycling is so important - it uses 95% less energy than extracting new aluminium from ore!