Key Definitions
- Alkali Metals
- The Group 1 elements in the periodic table.
- Halogens
- The Group 7 elements in the periodic table.
- Rusting
- The corrosion of iron in the presence of oxygen and water.
- Alloy
- A mixture of a metal with other elements (usually metals or carbon).
- Indicator
- A chemical that changes colour depending on the pH of the solution.
- Acid
- A proton (H+) donor.
- Base
- A proton (H+) acceptor.
- Salt
- A compound formed when the hydrogen in an acid is replaced by a metal or ammonium ion.
Group 1 - Alkali Metals
The Group 1 elements (Lithium, Sodium, Potassium, Rubidium, Caesium, and Francium) are known as the alkali metals. They are highly reactive metals located in the leftmost column of the periodic table.
Physical Properties
Unlike typical transition metals (which are hard and dense), Group 1 metals show unique physical trends:
- Softness: They are very soft and can easily be cut with a knife. A freshly cut surface is shiny grey but tarnishes rapidly in air.
- Low Density: They have exceptionally low densities. Lithium, sodium, and potassium are less dense than water and will float on it.
- Melting & Boiling Points: They have low melting points which decrease down the group (as the metallic bonds get weaker due to larger atomic radii).
Reactions with Water
Alkali metals react vigorously with cold water to produce a metal hydroxide solution and hydrogen gas:
where M represents any Group 1 metal. The resulting metal hydroxide solution is highly alkaline (pH 12-14) because it dissociates to release hydroxide ions (OH⁻), turning universal indicator purple.
| Metal | Observations in Water | Chemical Equation |
|---|---|---|
| Lithium (Li) | Fizzes slowly, floats on the surface, moves around, and gradually disappears. | 2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g) |
| Sodium (Na) | Fizzes rapidly, melts into a silvery ball (due to heat released), floats and dashes across the surface, and disappears. | 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) |
| Potassium (K) | Reacts violently, melts into a ball, floats and moves very rapidly, ignites hydrogen gas to burn with a lilac flame, and may end with a small crackle/explosion. | 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) |
Reactions with Oxygen
Alkali metals react rapidly with oxygen in the air to form metal oxides. This is why they are stored under oil to prevent contact with air and moisture. When burned in air, they form oxides:
For example: 4Li(s) + O₂(g) → 2Li₂O(s) (Lithium oxide).
Why Reactivity Increases Down the Group Extended
All Group 1 atoms have one electron in their outer shell. During reactions, they lose this electron to form a 1+ ion (M⁺). Reactivity increases down the group because:
- Atomic radius increases: As you go down the group, atoms have more occupied electron shells.
- Shielding increases: There are more inner shells of electrons shielding the outer electron from the positive charge of the nucleus.
- Weaker nuclear attraction: The outer electron is further away from the nucleus and more shielded, so the attractive force holding it is weaker.
- Easier electron loss: Consequently, the outer electron is lost more easily, making the element more reactive.
Predictions for Rubidium & Caesium
Based on the reactivity trend, we can make predictions for the lower Group 1 elements:
- Rubidium (Rb): Will react explosively with water, sinking and melting immediately, causing a rapid release of energy and hydrogen gas that explodes instantly.
- Caesium (Cs): Will react even more violently than rubidium, detonating violently the moment it contacts water, even at sub-zero temperatures.
Flame Tests (Syllabus 2.46)
We can identify specific metal ions (cations) by the characteristic flame colours they produce when heated in a Bunsen burner flame:
| Metal Ion | Flame Colour |
|---|---|
| Lithium (Li⁺) | Crimson Red |
| Sodium (Na⁺) | Yellow-Orange |
| Potassium (K⁺) | Lilac |
Flame Test Procedure:
- Dip a clean nichrome or platinum wire into concentrated hydrochloric acid (HCl) to clean it.
- Hold the wire in a hot, blue Bunsen burner flame. Repeat this cleaning process until the wire produces no color in the flame.
- Dip the clean wire back into the concentrated HCl, then into the solid sample so that some solid sticks to it (or dip into a solution of the sample).
- Place the wire in the hot, roaring blue flame of the Bunsen burner.
- Observe and record the color of the flame.
Group 7 - Halogens
The Group 7 elements (Fluorine, Chlorine, Bromine, Iodine, and Astatine) are non-metals known as the halogens. They exist as covalent diatomic molecules (X₂) joined by a single covalent bond.
Physical States & Colours
As you go down Group 7, the physical properties show a clear, gradual trend:
| Halogen | Formula | State at RTP | Colour | Vapour Colour |
|---|---|---|---|---|
| Fluorine | F₂ | Gas | Pale Yellow | Pale Yellow |
| Chlorine | Cl₂ | Gas | Pale Green | Pale Green |
| Bromine | Br₂ | Liquid | Red-Brown | Orange-Brown |
| Iodine | I₂ | Solid | Grey-Black | Purple (when sublimed) |
Physical Trends:
- Melting & Boiling Points: Increase down the group. This is because the molecules get larger (more electron shells and mass), which increases the strength of the intermolecular forces between molecules, requiring more energy to overcome.
- State at RTP: Changes from gas (F₂, Cl₂) to liquid (Br₂) to solid (I₂).
- Colour: Becomes darker down the group.
Why Reactivity Decreases Down the Group Extended
Halogen atoms have seven electrons in their outer shell. During reactions, they react by gaining one electron to achieve a stable full outer shell, forming 1- halide ions (X⁻). Reactivity decreases down the group because:
- Atomic radius increases: The atoms get larger as you go down the group.
- Shielding increases: There are more inner shells shielding the positive nucleus.
- Weaker nuclear attraction: The incoming electron is further away from the nucleus and more shielded, so the attractive force pulling it in is weaker.
- Harder electron gain: Consequently, it is harder for the atom to attract and gain the extra electron, making the element less reactive.
Displacement Reactions
A displacement reaction occurs when a more reactive halogen displaces a less reactive halogen from its halide solution.
| Halogen added | Potassium Chloride (Cl⁻) | Potassium Bromide (Br⁻) | Potassium Iodide (I⁻) |
|---|---|---|---|
| Chlorine Water (Cl₂) | No reaction | Orange solution formed (Br₂ displaced) | Brown solution formed (I₂ displaced) |
| Bromine Water (Br₂) | No reaction | No reaction | Brown solution formed (I₂ displaced) |
| Iodine Solution (I₂) | No reaction | No reaction | No reaction |
Equations:
- Chlorine + Potassium Bromide:
Molecular:Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
Ionic:Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq) - Chlorine + Potassium Iodide:
Molecular:Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)
Ionic:Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq) - Bromine + Potassium Iodide:
Molecular:Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)
Ionic:Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)
Redox Breakdown in terms of Electron Transfer:
Using the ionic equation: Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)
- Oxidation: Bromide ions lose electrons (OIL):
2Br⁻ → Br₂ + 2e⁻ - Reduction: Chlorine molecules gain electrons (RIG):
Cl₂ + 2e⁻ → 2Cl⁻ - Oxidizing Agent: Chlorine is the oxidizing agent because it accepts electrons, oxidizing the bromide ions.
- Reducing Agent: Bromide is the reducing agent because it donates electrons, reducing the chlorine molecules.
Predictions for Astatine
Astatine is located below iodine in Group 7. Based on the group trends, we can predict that Astatine will be:
- A black solid at room temperature, with an extremely high melting and boiling point.
- The least reactive of the halogens. It will not displace chloride, bromide, or iodide ions from their solutions.
Testing for Halide Ions (Syllabus 2.48)
To identify halide ions (Cl⁻, Br⁻, I⁻) in an unknown solution:
- Add dilute nitric acid (HNO₃) to the sample. This reacts with and removes any carbonate impurities that would produce a false positive. Do not use hydrochloric acid (HCl) as it introduces chloride ions.
- Add a few drops of silver nitrate solution (AgNO₃).
- Observe the color of the precipitate formed:
| Halide Ion | Precipitate Colour | Precipitate Formula | Ionic Equation |
|---|---|---|---|
| Chloride (Cl⁻) | White | AgCl(s) | Ag⁺(aq) + Cl⁻(aq) → AgCl(s) |
| Bromide (Br⁻) | Cream | AgBr(s) | Ag⁺(aq) + Br⁻(aq) → AgBr(s) |
| Iodide (I⁻) | Yellow | AgI(s) | Ag⁺(aq) + I⁻(aq) → AgI(s) |
Gases in the Atmosphere
The Current Atmosphere
The proportions of gases in the Earth’s atmosphere today have been relatively stable for about 200 million years:
| Gas | Formula | Proportion |
|---|---|---|
| Nitrogen | N₂ | ~78% |
| Oxygen | O₂ | ~21% |
| Argon | Ar | ~0.9% |
| Carbon dioxide | CO₂ | ~0.04% |
| Water vapour + trace gases | - | Variable, very small |
The proportions of gases in the Earth's atmosphere have been stable for around 200 million years, overwhelmingly dominated by nitrogen and oxygen.
Greenhouse Gases & the Greenhouse Effect
Some gases in the atmosphere absorb heat radiation (infrared) emitted by the Earth's surface and re-radiate it in all directions, including back towards the surface. This is the greenhouse effect, and it keeps the Earth warm enough to support life.
The greenhouse effect occurs when short-wavelength solar radiation passes through the atmosphere, but the long-wavelength infrared radiation emitted by the Earth is absorbed and re-emitted by greenhouse gases like CO₂, CH₄, and H₂O.
The Main Greenhouse Gases
| Gas | Formula | Main human sources |
|---|---|---|
| Carbon dioxide | CO₂ | Burning fossil fuels, deforestation |
| Methane | CH₄ | Cattle farming, rice paddies, landfill, natural gas leaks |
| Water vapour | H₂O | Natural evaporation (not directly from human activity) |
Climate Change
Since the Industrial Revolution (mid-1800s), human activities have significantly increased the concentration of greenhouse gases. This enhanced greenhouse effect is causing global temperatures to rise - known as global warming.
Consequences of Climate Change
- Polar ice caps melting → sea level rise → flooding of low-lying areas
- More extreme weather events (droughts, floods, storms)
- Changes to ecosystems and habitats
- Impact on agriculture and food production
- Migration and redistribution of species
Why is There a Debate?
While the vast majority of scientists agree that human activity is the main cause, the evidence involves very complex climate models and long-term data. Media, politics, and economics also influence public perception.
Human Activities Driving Climate Change
- Burning fossil fuels (coal, oil, gas) for energy, transport, and industry → releases CO₂.
- Deforestation → fewer trees to absorb CO₂ by photosynthesis, and burning trees releases stored carbon.
- Agriculture → cattle produce methane; rice paddies release methane; fertilisers release nitrous oxide.
- Landfill → decomposing waste releases methane.
Carbon Footprint
A carbon footprint is the total amount of carbon dioxide and other greenhouse gases emitted over the full life cycle of a product, service, or event.
Ways to Reduce Carbon Footprint
- Using renewable energy sources (solar, wind, tidal) instead of fossil fuels.
- Improving energy efficiency of buildings and vehicles.
- Carbon capture and storage (CCS) technology.
- Reducing waste, reusing materials, and recycling.
- Planting trees to absorb CO₂.
- Using public transport or cycling instead of driving.
Comparing carbon footprints
A school is comparing two options for heating: natural gas boiler or a ground-source heat pump powered by renewable electricity. Discuss which has the lower carbon footprint.
Gas boiler: Burns natural gas (fossil fuel) → releases CO₂ directly. Manufacturing and transporting gas also contributes. Carbon footprint = high.
Heat pump (renewable): Uses electricity from renewable sources → no CO₂ during operation. However, manufacturing the pump, drilling boreholes, and transporting materials all produce some CO₂. Carbon footprint = lower, but not zero.
Conclusion: The heat pump has a significantly lower carbon footprint over its lifetime, but the full life cycle (manufacturing + installation + disposal) must be considered.
Atmospheric Pollutants
Burning fossil fuels releases several harmful pollutants:
Burning fossil fuels releases a variety of atmospheric pollutants, each with distinct negative consequences for human health and the environment.
| Pollutant | Source | Harm caused | Solution |
|---|---|---|---|
| Carbon monoxide (CO) | Incomplete combustion | Toxic - binds to haemoglobin, prevents O₂ transport. Colourless/odourless. | Catalytic converters (CO → CO₂) |
| Sulfur dioxide (SO₂) | Burning fuels with sulfur impurities | Acid rain → damages buildings (limestone), kills aquatic life, harms plants | Flue gas desulfurisation; use low-sulfur fuels |
| Nitrogen oxides (NOₓ) | N₂ + O₂ at high engine temps | Acid rain, photochemical smog, respiratory problems | Catalytic converters (NOₓ → N₂) |
| Particulates (soot) | Incomplete combustion | Respiratory problems, global dimming, darkens buildings | Particulate filters in vehicles |
Experiments to Determine the Percentage of Oxygen in Air
Air is a mixture of gases, with nitrogen (~78%) and oxygen (~21%) being the main components. We can determine the percentage of oxygen using reactive substances that consume oxygen when heated or rusted.
Experiment 1: Using Copper Turnings
Setup: Two gas syringes are connected by a glass tube containing copper turnings. One syringe contains 100 cm3 of air, and the other is empty. The tube containing copper is heated strongly with a Bunsen burner.
Method: The air is passed back and forth over the hot copper turnings using the syringe plungers. Copper reacts with oxygen in the air to form black copper(II) oxide:
2Cu(s) + O2(g) → 2CuO(s)
Result: As the oxygen is consumed, the volume of gas decreases. The heating is continued until the volume stops changing. After cooling to room temperature, the final volume of gas is approximately 79 cm3. This shows that 21 cm3 (21%) of the original air was oxygen.
Experiment 2: Using Iron Wool
Setup: Wet iron wool is pushed to the bottom of a test tube. The tube is inverted and placed inside a beaker of water. The initial volume of air inside the tube is marked.
Method: The apparatus is left for about a week. The iron wool reacts slowly with oxygen and water in the tube to form hydrated iron(III) oxide (rust). As oxygen is consumed, the water level rises to fill the space left by the oxygen.
Result: The water level rises by approximately 20-21% of the original height of the air column in the tube, showing that oxygen makes up about 21% of air.
Experiment 3: Using Phosphorus
Setup: A piece of phosphorus is placed in an evaporating basin floating on water, covered by a bell jar. The initial water level inside the bell jar is marked.
Method: The phosphorus is ignited using a hot wire. It burns rapidly, reacting with oxygen to form white clouds of phosphorus pentoxide (P4O10), which dissolve in the water. As oxygen is consumed, water rises inside the bell jar.
P4(s) + 5O2(g) → P4O10(s)
Result: The final water level rises by approximately 21% of the original volume of air inside the bell jar.
Reactivity Series
4.4.1 Metal Oxides
When metals react with oxygen, they form metal oxides. The vigour of this reaction depends on the metal's position in the reactivity series.
Magnesium burns in air with a bright white flame to produce the white powder magnesium oxide.
At this stage, we define oxidation as the gain of oxygen and reduction as the loss of oxygen. In the reaction above, magnesium is oxidised (it gains oxygen).
Here, copper oxide is reduced (it loses oxygen) and carbon is oxidised (it gains oxygen). The carbon acts as a reducing agent.
The Reactivity Series
The reactivity series ranks metals in order of how vigorously they react. A metal's reactivity is determined by how readily it loses its outer shell electrons to form positive ions (cations).
Order (most to least reactive)
The reactivity series showing common extraction methods. Carbon and hydrogen are included for reference.
Reactions with water and dilute acid
- Potassium: reacts violently with cold water, ignites with a lilac flame. Too dangerous for acid reactions.
- Sodium: vigorous fizzing with cold water, melts into a ball on the surface. Too dangerous for acid reactions.
- Calcium: steady bubbling with cold water, producing a cloudy solution of calcium hydroxide.
- Magnesium: very slow reaction with cold water. Reacts vigorously with steam and with dilute acid (rapid fizzing).
- Zinc: no reaction with cold water. Slow, steady bubbling with dilute acid.
- Iron: no reaction with cold water. Very slow reaction with dilute acid.
- Copper: no reaction with cold water or dilute acid.
| Metal | Reaction with Water | Reaction with Dilute Acid |
|---|---|---|
| Potassium | Violent - lilac flame | Too dangerous |
| Sodium | Vigorous fizzing, melts | Too dangerous |
| Calcium | Steady bubbling, cloudy solution | Vigorous fizzing |
| Magnesium | Very slow with cold water | Rapid fizzing |
| Zinc | No reaction | Slow, steady bubbles |
| Iron | No reaction | Very slow |
| Copper | No reaction | No reaction |
See where each metal sits on our Interactive Periodic Table and compare their properties.
Displacement Reactions
A more reactive metal can displace a less reactive metal from a compound in solution. This happens because the more reactive metal has a greater tendency to form positive ions.
When an iron nail is placed in blue copper sulfate solution, the solution gradually fades from blue to pale green (iron(II) sulfate), and a reddish-brown coating of copper metal appears on the nail.
Predicting displacement
Will zinc react with magnesium chloride solution?
Step 1: Check the reactivity series. Magnesium is above zinc.
Step 2: A less reactive metal cannot displace a more reactive metal. Therefore, no reaction occurs.
Predicting displacement (reaction occurs)
What happens when magnesium ribbon is placed in zinc sulfate solution?
Step 1: Magnesium is above zinc in the reactivity series.
Step 2: A more reactive metal displaces a less reactive metal. So magnesium will displace zinc.
Equation: Mg(s) + ZnSO₄(aq) → MgSO₄(aq) + Zn(s)
Observations: The magnesium dissolves. A grey coating of zinc metal appears. The solution warms up (exothermic).
Oxidation & Reduction (OIL RIG)
At a basic level, oxidation means gaining oxygen and reduction means losing oxygen. These complementary processes always happen together in a redox reaction.
Redox in Terms of Electron Transfer Extended
For Higher Tier, oxidation and reduction are defined more precisely in terms of electron transfer:
- Oxidation Is Loss of electrons (OIL).
- Reduction Is Gain of electrons (RIG).
A substance that is oxidised is the reducing agent (it causes another substance to be reduced by donating electrons). A substance that is reduced is the oxidising agent.
Half-equations
In the displacement of copper by magnesium:
This can be split into two half-equations showing the electron transfer:
Magnesium loses two electrons (oxidised). Copper ions gain two electrons (reduced). The electrons lost by one species are gained by the other.
Extraction & Uses of Metals
Extraction of Metals
Most metals are found in the Earth's crust as ores, which are rocks containing enough metal to make extraction economically worthwhile.
Method depends on reactivity
- Metals below carbon (zinc, iron, copper): extracted by reduction with carbon in a blast furnace.
- Metals above carbon (aluminium, sodium, potassium): too reactive for carbon reduction, so must be extracted by electrolysis.
- Very unreactive metals (gold, platinum): found native in the Earth's crust and need no chemical extraction.
Reduction with carbon
The carbon removes oxygen from the metal oxide (reducing it), and is itself oxidised to carbon dioxide. This is a redox reaction.
Biological methods
Phytomining uses plants that absorb metal compounds from the soil. The plants are harvested and burned, and copper is extracted from the ash. This is useful for low-grade ores that are uneconomical to mine traditionally.
Bioleaching uses bacteria to produce a leachate solution containing dissolved metal compounds. The copper ions can then be recovered from the leachate, for example by displacement with scrap iron or by electrolysis.
Choosing the extraction method
How would you extract (a) iron, (b) aluminium, (c) gold from their ores?
(a) Iron is below carbon in the reactivity series → extract by reduction with carbon (blast furnace). 2Fe₂O₃ + 3C → 4Fe + 3CO₂
(b) Aluminium is above carbon → too reactive for carbon reduction, so extract by electrolysis of Al₂O₃ dissolved in molten cryolite.
(c) Gold is very unreactive → found native (uncombined) in the Earth's crust. No chemical extraction needed.
Rusting & Prevention
Rusting is the specific corrosion of iron (and steel). It requires both oxygen and water. Rust is a chemical compound called hydrated iron(III) oxide.
Iron + Oxygen + Water → Hydrated Iron(III) Oxide
Rust Prevention Methods
- Barrier Methods: Coating the iron with paint, grease/oil, or plastic. This physically prevents oxygen and water from reaching the surface of the iron. If the barrier is scratched, however, the iron will rust.
- Galvanising: Coating the iron with a layer of zinc. Zinc is more reactive than iron, so it reacts with oxygen and water first. Even if the zinc coating is scratched, the zinc continues to protect the iron by reacting preferentially.
- Sacrificial Protection: Attaching blocks of a more reactive metal (such as zinc or magnesium) to the iron structure. The more reactive metal oxidizes (loses electrons) in preference to the iron:
Zn → Zn2+ + 2e-
This is commonly used on ship hulls, oil rigs, and underground pipelines.
Steel Types & Alloys
Steel is an alloy of iron and carbon, often with other metals added to modify its properties.
| Steel Type | Composition | Properties | Uses |
|---|---|---|---|
| Mild Steel (Low-carbon steel) | Iron + ~0.25% Carbon | Strong, malleable, ductile, rusts easily | Car bodies, machinery, ship building, structural beams |
| High-carbon Steel | Iron + ~0.6-1.5% Carbon | Very hard, strong, brittle | Cutting tools, chisels, drill bits, masonry nails |
| Stainless Steel | Iron + Chromium + Nickel | Hard, extremely resistant to corrosion | Cutlery, kitchen sinks, chemical reaction vessels |
Uses of Other Metals
- Aluminium: Used for aircraft bodies (low density and strong when alloyed) and food containers/cans (resistant to corrosion due to protective oxide layer).
- Copper: Used for electrical wiring (excellent electrical conductor, ductile) and water pipes (highly unreactive, malleable).
Extraction of Aluminium Extended
Aluminium is too reactive to be extracted from its ore (bauxite) by reduction with carbon. Instead, it must be extracted using electrolysis of molten aluminium oxide (alumina, Al₂O₃).
The Role of Cryolite
Pure aluminium oxide has an extremely high melting point of about 2050°C. Melting it directly requires an enormous amount of thermal energy, which is very expensive.
To overcome this, aluminium oxide is dissolved in molten cryolite (Na₃AlF₆). This serves two vital purposes:
- It lowers the operating temperature of the cell to about 950°C, significantly reducing heating costs.
- It improves the electrical conductivity of the electrolyte, making the process more efficient.
Electrode Reactions
The electrolysis cell consists of a steel shell lined with carbon (which acts as the negative cathode) and carbon blocks dipped into the electrolyte (which act as the positive anodes):
- At the Cathode (Negative Electrode): Aluminium ions are reduced by gaining electrons to form molten aluminium metal:
Al³⁺ + 3e⁻ → Al(l)
Since molten aluminium is denser than the electrolyte, it sinks to the bottom of the cell, where it is periodically siphoned/tapped off. - At the Anode (Positive Electrode): Oxide ions are oxidized by losing electrons to form oxygen gas:
2O²⁻ → O₂(g) + 4e⁻
Anode Wear & Replacement
At the high operating temperature of the cell (~950°C), the oxygen gas produced at the anodes reacts with the carbon anodes themselves:
This reaction burns away the carbon anodes over time, producing carbon dioxide gas. As a result, the carbon anodes must be regularly replaced, adding to the cost of the process.
Aluminium Electrolysis Cell
Schematic of the industrial electrolysis of aluminium oxide dissolved in molten cryolite, operating at around 950°C.
Acids, Alkalis & Titrations
4.4.2 Reactions of Acids
Acids react with metals, metal oxides, metal hydroxides, and metal carbonates. The type of salt produced depends on which acid is used.
Acid + Metal → Salt + Hydrogen
Only metals above hydrogen in the reactivity series react with dilute acids. The hydrogen gas can be tested with a lighted splint, which produces a squeaky pop.
Acid + Metal Oxide → Salt + Water
Acid + Metal Hydroxide → Salt + Water
Acid + Metal Carbonate → Salt + Water + CO₂
Carbon dioxide can be tested by bubbling it through limewater, which turns milky (cloudy white).
Naming salts
Struggle with salt names? Common naming errors are covered in 10 Most Common Mistakes in GCSE Chemistry Exams.
Predicting the products: Zn + H₂SO₄
Step 1: Identify the acid. Sulfuric acid → produces a sulfate salt.
Step 2: Identify the metal. Zinc. So the salt is zinc sulfate (ZnSO₄).
Step 3: Metal + acid → salt + hydrogen.
Balanced equation: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
Observation: Zinc dissolves, bubbles of gas produced (squeaky pop with lighted splint).
Neutralisation
Neutralisation is the reaction between an acid and a base (or alkali) to produce a salt and water. A base is any substance that neutralises an acid. An alkali is a soluble base that produces hydroxide ions (OH⁻) in solution.
In terms of H⁺ and OH⁻ ions:
The hydrogen ions from the acid combine with the hydroxide ions from the base to form water. This is why the solution becomes less acidic (or less alkaline) during neutralisation.
pH Scale & Indicators
The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is.
- pH 0 to 6: Acidic (lower pH = stronger acid). Acids produce H⁺ ions in solution.
- pH 7: Neutral.
- pH 8 to 14: Alkaline (higher pH = stronger alkali). Alkalis produce OH⁻ ions in solution.
Universal indicator is a mixture of dyes that changes colour across the full pH range (red for strong acid, green for neutral, purple for strong alkali). A pH probe connected to a pH meter gives a more precise numerical reading.
Strong & Weak Acids
Full Ionisation vs Partial Ionisation Extended
Strong acids
Strong acids fully ionise (dissociate) in water. Every molecule splits to release H⁺ ions. This is an irreversible process.
Examples: hydrochloric acid (HCl), sulfuric acid (H₂SO₄), nitric acid (HNO₃).
Weak acids
Weak acids only partially ionise in water. A reversible equilibrium exists, with most molecules remaining undissociated.
Examples: ethanoic acid (CH₃COOH), citric acid, carbonic acid.
Strong vs concentrated
These terms mean different things:
- Strong/weak refers to the degree of ionisation (how much the acid splits into ions).
- Concentrated/dilute refers to the amount of acid dissolved per unit volume of solution.
A dilute strong acid and a concentrated weak acid are both valid concepts.
pH and hydrogen ion concentration
The pH scale is logarithmic. As the pH decreases by 1 unit, the hydrogen ion concentration increases by a factor of 10. A decrease of 2 pH units means the H⁺ concentration is 100 times greater.
At the same concentration, a strong acid will have a lower pH than a weak acid because more H⁺ ions are present in solution.
Comparing H⁺ concentration
Two acids both have a concentration of 0.1 mol/dm³. Acid A has pH 1, Acid B has pH 3. Which is the strong acid?
Step 1: pH 1 is lower than pH 3, so Acid A has a higher H⁺ concentration.
Step 2: The pH difference is 2 units. The pH scale is logarithmic, so Acid A has 10² = 100 times more H⁺ ions than Acid B.
Step 3: At the same concentration, the acid with more H⁺ ions is the strong acid. Acid A is strongly ionised; Acid B is weakly ionised.
Titrations Extended
Required Practical 2: Neutralisation by Titration Extended
A titration is used to find the exact volume of acid needed to neutralise a known volume of alkali (or vice versa). This is Required Practical 2.
Method
- Use a pipette to measure a fixed volume of alkali (e.g. 25.0 cm³) into a conical flask.
- Add a few drops of a suitable indicator (e.g. Phenolphthalein or methyl orange).
- Fill a burette with the acid and record the starting volume.
- Add the acid dropwise to the alkali, swirling the flask continuously.
- Stop adding acid when the indicator permanently changes colour (the end point).
- Record the final burette reading and calculate the volume of acid used (the titre).
- Repeat until you get concordant results (titres within 0.10 cm³ of each other).
Key points for exam questions
- Use a white tile under the flask to see colour changes more clearly.
- The titre is: final burette reading minus initial burette reading.
- Take the mean of concordant results only (ignore anomalous values).
- You must read the burette at the bottom of the meniscus.
Titration calculations use moles. Need a refresher? Read The Mole Explained: The One Concept That Unlocks All of GCSE Chemistry.
Acids, Bases & Salt Preparations
Making Soluble Salts Extended
To make a pure, dry sample of a soluble salt from an insoluble reactant (such as a metal oxide, hydroxide, or carbonate):
- Warm dilute acid in a beaker using a Bunsen burner.
- Add the insoluble base (metal oxide, hydroxide, or carbonate) in small amounts, stirring after each addition, until no more dissolves and excess solid remains.
- Filter the mixture to remove the excess unreacted solid. The filtrate is the salt solution.
- Pour the filtrate into an evaporating basin. Heat gently until about half of the water has evaporated.
- Leave the solution to cool and crystallise slowly. Pat the crystals dry with filter paper.
Naming the salt: What salt is produced when magnesium oxide reacts with hydrochloric acid?
Step 1: The metal comes from material MgO → magnesium.
Step 2: Hydrochloric acid produces chloride salts.
Step 3: The salt is magnesium chloride (MgCl₂).
Full equation: MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)
The Haber Process Extended
The Haber process is used to manufacture ammonia (NH₃) on an industrial scale.
Conditions
- Temperature: ~450°C (a compromise - low enough for a reasonable yield, high enough for a reasonable rate).
- Pressure: ~200 atmospheres (high pressure favours the forward reaction as there are fewer moles of gas on the right).
- Catalyst: Iron (speeds up the reaction without being consumed).
- Cooling: The mixture is cooled to -33°C to liquefy and remove the ammonia. Yield is ~15%. Unreacted gases are recycled.
The Haber process creates ammonia efficiently by recycling unreacted nitrogen and hydrogen gases, establishing a continuous industrial loop.
NPK Fertilisers
Ammonia is used to make ammonium salts (e.g., ammonium nitrate, NH₄NO₃) which are used as fertilisers. Fertilisers provide nitrogen (N), phosphorus (P), and potassium (K) to help crops grow.
Le Chatelier’s Principle applied to the Haber Process
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ/mol (exothermic forward)
Temperature (450°C): The forward reaction is exothermic. A lower temperature would increase the yield (Le Chatelier’s shifts to oppose cooling → forward). BUT rate would be too slow. 450°C is a compromise - reasonable yield and acceptable rate.
Pressure (200 atm): 4 moles of gas on the left → 2 moles on the right. High pressure shifts equilibrium to the side with fewer moles → forward → more NH₃. BUT very high pressures are expensive and dangerous. 200 atm is a compromise.
Iron catalyst: Speeds up both forward and reverse reactions equally - does NOT change yield but reaches equilibrium FASTER.
Recycling: Unreacted N₂ and H₂ are recycled back over the catalyst to improve overall conversion.
The Contact Process
The Contact Process is the industrial manufacture of sulfuric acid (H2SO4). The raw materials are sulfur (or metal sulfides like zinc blende), air (to provide oxygen), and water.
Steps in the Manufacture of Sulfuric Acid
- Step 1: Production of Sulfur Dioxide
Sulfur is burned in air to form sulfur dioxide gas:S(s) + O2(g) → SO2(g) - Step 2: Conversion of Sulfur Dioxide to Sulfur Trioxide
Sulfur dioxide is reacted with more oxygen in a reversible, exothermic reaction:2SO2(g) + O2(g) ⇌ 2SO3(g) (ΔH = -197 kJ/mol)
Reaction Conditions:- Catalyst: Vanadium(V) oxide (V2O5)
- Temperature: 450°C (compromise temperature: high enough for a fast rate, but low enough to maintain a good yield of SO3 since the forward reaction is exothermic)
- Pressure: 1-2 atm (normal pressure is sufficient because the yield of SO3 is already very high, and high pressures are expensive and dangerous)
- Step 3: Dissolving SO3 into Acid
Sulfur trioxide is dissolved in concentrated sulfuric acid to form a liquid called oleum (H2S2O7):SO3(g) + H2SO4(l) → H2S2O7(l)Why not dissolve SO3 directly in water? The reaction is extremely exothermic and produces a dense, highly acidic mist of sulfuric acid that is dangerous and very difficult to condense into a liquid. - Step 4: Dilution of Oleum
Oleum is carefully mixed with water to form concentrated sulfuric acid:H2S2O7(l) + H2O(l) → 2H2SO4(l)
Uses of Sulfuric Acid
- Manufacture of fertilizers (such as ammonium sulfate).
- Manufacture of detergents.
- Used in paints, pigments, and dyes.
Chemical Tests
Flame Tests Extended
Different metal ions produce characteristic flame colours when heated in a Bunsen flame.
| Metal ion | Flame colour |
|---|---|
| Lithium (Li⁺) | Crimson red |
| Sodium (Na⁺) | Yellow |
| Potassium (K⁺) | Lilac |
| Calcium (Ca²⁺) | Orange-red |
| Copper (Cu²⁺) | Green |
Each metal ion produces a unique, characteristic flame colour when heated. This allows identification of unknown metal ions in a sample.
Method
- Clean a nichrome wire loop by dipping it in hydrochloric acid and holding it in a blue Bunsen flame until no colour is seen.
- Dip the clean wire into the sample.
- Hold the sample in the flame and observe the colour.
Metal Hydroxide Precipitates Extended
Adding sodium hydroxide (NaOH) solution to solutions containing metal ions produces coloured precipitates that identify the ion.
| Metal ion | Precipitate colour | Formula |
|---|---|---|
| Calcium (Ca²⁺) | White | Ca(OH)₂ |
| Magnesium (Mg²⁺) | White | Mg(OH)₂ |
| Aluminium (Al³⁺) | White (dissolves in excess NaOH) | Al(OH)₃ |
| Copper(II) (Cu²⁺) | Blue | Cu(OH)₂ |
| Iron(II) (Fe²⁺) | Green | Fe(OH)₂ |
| Iron(III) (Fe³⁺) | Brown | Fe(OH)₃ |
Adding sodium hydroxide solution produces coloured precipitates that can identify the metal ion present. Three ions give white precipitates, but aluminium's dissolves in excess NaOH.
Testing for Carbonates Extended
To test for carbonate ions (CO₃²⁻), add dilute hydrochloric acid. Carbonates fizz (effervesce) as they decompose to produce carbon dioxide gas.
Confirm CO₂ by passing it through limewater - it turns milky (cloudy).
Testing for Halides Extended
Add dilute nitric acid then silver nitrate solution (AgNO₃).
| Halide ion | Precipitate colour | Formula |
|---|---|---|
| Chloride (Cl⁻) | White | AgCl |
| Bromide (Br⁻) | Cream | AgBr |
| Iodide (I⁻) | Yellow | AgI |
Adding silver nitrate solution (after acidifying with dilute nitric acid) produces coloured precipitates to identify halide ions: white for chloride, cream for bromide, yellow for iodide.
Testing for Sulfates Extended
Add dilute hydrochloric acid then barium chloride solution (BaCl₂).
A white precipitate of barium sulfate (BaSO₄) confirms sulfate ions are present.
Tests for Gases
| Gas | Test | Positive result |
|---|---|---|
| Hydrogen (H₂) | Burning splint | Squeaky pop |
| Oxygen (O₂) | Glowing splint | Splint relights |
| Carbon dioxide (CO₂) | Bubble through limewater | Turns milky (cloudy) |
| Chlorine (Cl₂) | Damp litmus paper | Bleaches paper white |
The four required gas tests: hydrogen produces a squeaky pop, oxygen relights a glowing splint, carbon dioxide turns limewater milky, and chlorine bleaches damp litmus paper white.
Chemical & Physical Tests for Water
1. Chemical Test for the Presence of Water
To identify if a liquid contains water, we use anhydrous copper(II) sulfate.
- Anhydrous copper(II) sulfate is a white powder.
- When water is added, it turns into hydrated copper(II) sulfate, which is blue.
- Equation: CuSO4(s) + 5H2O(l) → CuSO4•5H2O(s)
- Reverse reaction: Heating blue copper(II) sulfate crystals drives off the water, turning it white again (dehydration).
2. Physical Test for Pure Water
The copper(II) sulfate test only checks for the presence of water, not its purity (e.g., salt water will also turn it blue). To check if water is chemically pure, we measure its physical properties:
- Boiling Point: Pure water boils at exactly 100°C at 1 atm pressure.
- Melting/Freezing Point: Pure water freezes at exactly 0°C.
- Impurities: The presence of dissolved impurities (like salt) raises the boiling point and lowers/depresses the freezing point.