Chemical changes are everywhere around you - from the rust...
Chemistry118 views·Updated 18 Sept 2026·3 pagesGCSE Chemistry: Guide to Unit 4 - Chemical Changes
Erin@erin_agwa
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1
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Reactions of Metals
Metal reactivity determines how vigorously metals react with oxygen, water, and acids. When metals react with oxygen, they form metal oxides - think about how iron rusts when exposed to air and moisture.
The most reactive metals like potassium and sodium react explosively with water, producing metal hydroxides (which are alkaline) and hydrogen gas. You can test for hydrogen because it burns with a distinctive "pop" sound when lit.
Acids react with metals to produce salts and hydrogen gas. The type of salt depends on the acid used: hydrochloric acid makes chlorides, sulfuric acid makes sulfates, and nitric acid makes nitrates. For example, zinc + hydrochloric acid → zinc chloride + hydrogen.
Remember: The reactivity series goes potassium, sodium, lithium, calcium, magnesium, zinc, iron, copper (most to least reactive). Carbon and hydrogen fit between zinc and iron.
Metal compounds also react with acids. Metal oxides and carbonates both react with acids, but carbonates also produce carbon dioxide gas, which turns limewater cloudy - a useful test you'll need to remember.
2
of 3
Metal Extraction and Key Reactions
Unreactive metals like gold are found naturally as pure elements in the Earth. More reactive metals need to be extracted from their compounds using reduction with carbon, which displaces oxygen from metal oxides.
OILRIG is your key memory tool: Oxidation Is Loss, Reduction Is Gain (of electrons). When carbon removes oxygen from a metal oxide, the metal is reduced whilst carbon is oxidised.
Neutralisation happens when acids and alkalis react together, moving the pH closer to 7. Strong acids like HCl completely ionise in water, whilst weak acids like citric acid only partially ionise. The key reaction is: H⁺ + OH⁻ → H₂O.
Essential equation: At the heart of neutralisation is hydrogen ions from acids combining with hydroxide ions from alkalis to make water.
Electrolysis uses electricity to split ionic compounds. The compound (called the electrolyte) must be molten or dissolved so ions can move. Positive ions (cations) go to the negative electrode (cathode), whilst negative ions (anions) go to the positive electrode (anode).
3
of 3
Aluminium Extraction and Energy Changes
Aluminium extraction is a perfect example of electrolysis in action. Aluminium oxide is mixed with cryolite to help disrupt the lattice structure, making the process more efficient.
This process is incredibly expensive because it requires massive amounts of energy to melt the aluminium oxide and run the electrical current. The carbon electrodes also wear away and need replacing regularly.
Energy transfers happen in every chemical reaction, and energy is always conserved. Exothermic reactions release energy to the surroundings (heating them up) - think combustion, oxidation, and neutralisation reactions. These are used in self-heating cans and hand warmers.
Energy tip: If a reaction makes things warmer, it's exothermic. If it makes things cooler, it's endothermic.
Endothermic reactions absorb energy from the surroundings, cooling them down. Examples include thermal decomposition (like heating copper carbonate), mixing citric acid with sodium hydrogencarbonate, and photosynthesis in plants.
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Chemistry118 views·Updated 18 Sept 2026·3 pagesGCSE Chemistry: Guide to Unit 4 - Chemical Changes
Erin@erin_agwa
Chemical changes are everywhere around you - from the rust forming on bikes to the energy released when you burn fuel. Understanding how metals react, how we extract them, and how energy moves during reactions will help you make sense...
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Reactions of Metals
Metal reactivity determines how vigorously metals react with oxygen, water, and acids. When metals react with oxygen, they form metal oxides - think about how iron rusts when exposed to air and moisture.
The most reactive metals like potassium and sodium react explosively with water, producing metal hydroxides (which are alkaline) and hydrogen gas. You can test for hydrogen because it burns with a distinctive "pop" sound when lit.
Acids react with metals to produce salts and hydrogen gas. The type of salt depends on the acid used: hydrochloric acid makes chlorides, sulfuric acid makes sulfates, and nitric acid makes nitrates. For example, zinc + hydrochloric acid → zinc chloride + hydrogen.
Remember: The reactivity series goes potassium, sodium, lithium, calcium, magnesium, zinc, iron, copper (most to least reactive). Carbon and hydrogen fit between zinc and iron.
Metal compounds also react with acids. Metal oxides and carbonates both react with acids, but carbonates also produce carbon dioxide gas, which turns limewater cloudy - a useful test you'll need to remember.
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Metal Extraction and Key Reactions
Unreactive metals like gold are found naturally as pure elements in the Earth. More reactive metals need to be extracted from their compounds using reduction with carbon, which displaces oxygen from metal oxides.
OILRIG is your key memory tool: Oxidation Is Loss, Reduction Is Gain (of electrons). When carbon removes oxygen from a metal oxide, the metal is reduced whilst carbon is oxidised.
Neutralisation happens when acids and alkalis react together, moving the pH closer to 7. Strong acids like HCl completely ionise in water, whilst weak acids like citric acid only partially ionise. The key reaction is: H⁺ + OH⁻ → H₂O.
Essential equation: At the heart of neutralisation is hydrogen ions from acids combining with hydroxide ions from alkalis to make water.
Electrolysis uses electricity to split ionic compounds. The compound (called the electrolyte) must be molten or dissolved so ions can move. Positive ions (cations) go to the negative electrode (cathode), whilst negative ions (anions) go to the positive electrode (anode).
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Aluminium Extraction and Energy Changes
Aluminium extraction is a perfect example of electrolysis in action. Aluminium oxide is mixed with cryolite to help disrupt the lattice structure, making the process more efficient.
This process is incredibly expensive because it requires massive amounts of energy to melt the aluminium oxide and run the electrical current. The carbon electrodes also wear away and need replacing regularly.
Energy transfers happen in every chemical reaction, and energy is always conserved. Exothermic reactions release energy to the surroundings (heating them up) - think combustion, oxidation, and neutralisation reactions. These are used in self-heating cans and hand warmers.
Energy tip: If a reaction makes things warmer, it's exothermic. If it makes things cooler, it's endothermic.
Endothermic reactions absorb energy from the surroundings, cooling them down. Examples include thermal decomposition (like heating copper carbonate), mixing citric acid with sodium hydrogencarbonate, and photosynthesis in plants.
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