Assessed in Paper 1A (multiple choice), Paper 1B (data-based) and Paper 2 (short and extended response). IB Chemistry guide (first assessment 2025).
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R1.3.1 Combustion
Combustion
An exothermic reaction in which a substance reacts with oxygen when heated, releasing energy as heat and often light. Oxygen is a reactant: its atoms end up in the products. Reactive metals (e.g. 2Mg(s) + O₂(g) → 2MgO(s)), non-metals (e.g. S(s) + O₂(g) → SO₂(g); C(s) + O₂(g) → CO₂(g); 2H₂(g) + O₂(g) → 2H₂O(l); P₄(s) + 5O₂(g) → P₄O₁₀(s) in excess oxygen) and organic compounds all undergo combustion, forming oxides.
Complete combustion of hydrocarbons and alcohols
Combustion in a plentiful supply of oxygen, in which all the carbon in an organic compound is oxidized to carbon dioxide and all the hydrogen to water. Examples: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l); C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l). To deduce the O₂ coefficient, count the O atoms in the products and subtract any O atoms already in the fuel (one per –OH group in an alcohol); fractional coefficients such as 2CH₃OH + 3O₂ → 2CO₂ + 4H₂O (or CH₃OH + 1½O₂) are cleared by doubling.
Formulae in combustion equations
Elements that exist as diatomic molecules are written as molecules (O₂, H₂, N₂); formulae of ionic oxides are deduced from ion charges so that the total charge is zero (Mg²⁺ and O²⁻ give MgO; Al³⁺ and O²⁻ give Al₂O₃). An equation is balanced only by changing coefficients, never subscripts, because a formula identifies the substance.
Students often think All the oxygen atoms in the products come from O₂, so the alcohol's own oxygen atom can be ignored when working out the amount of O₂. In fact Yes. The O atom of the –OH group is already on the reactant side, so O₂ needs to supply only the rest: C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l), not 3½O₂.
Students often think Diatomic elements such as oxygen and hydrogen can be written and counted as single atoms (O, H), so the number of atoms needed is taken as the number of particles that react. In fact As O₂. Oxygen, like hydrogen, nitrogen and the halogens, exists as diatomic molecules, so each O₂ supplies two O atoms and the number of O₂ molecules is half the number of O atoms needed.
R1.3.2 Incomplete combustion
Incomplete combustion
Combustion of an organic compound in a limited supply of oxygen, in which some of the carbon is not fully oxidized and forms carbon monoxide, CO, or carbon (soot), C, instead of carbon dioxide; the hydrogen still forms water. Examples: 2CH₄(g) + 3O₂(g) → 2CO(g) + 4H₂O(l); CH₄(g) + O₂(g) → C(s) + 2H₂O(l); C₂H₅OH(l) + 2O₂(g) → 2CO(g) + 3H₂O(l); C₂H₅OH(l) + O₂(g) → 2C(s) + 3H₂O(l). Less energy is released per mole of fuel than in complete combustion.
Carbon monoxide and soot
Products of incomplete combustion. Carbon monoxide is a colourless, odourless, toxic gas: it binds to haemoglobin in the blood more strongly than oxygen, reducing oxygen transport. Soot is finely divided carbon; it blackens surfaces, gives a smoky yellow flame, and forms particulates that damage the lungs.
Students often think Incomplete combustion gives the same products as complete combustion, CO₂ and H₂O; it only means that some of the fuel is left unburnt. In fact The oxygen supply is too limited to oxidize all the carbon to CO₂, so the products change: carbon monoxide and/or carbon (soot) form as well as, or instead of, CO₂. Water still forms.
Students often think Soot and carbon monoxide come from impurities in the fuel, not from the fuel's own carbon; a pure fuel would not produce them. In fact From the carbon atoms of the fuel itself, which are not fully oxidized to CO₂ because the oxygen supply is limited.
R1.3.3 Fossil fuels
Fossil fuels
Coal, crude oil and natural gas, formed from the remains of organisms over millions of years. Coal is a solid with a high proportion of carbon; crude oil is a liquid mixture of hydrocarbons that is refined into fuels such as petrol (containing octane, C₈H₁₈); natural gas is mainly methane, CH₄. They are non-renewable because they form far more slowly than they are used.
Energy released per unit mass (specific energy)
The energy released when a unit mass of fuel burns completely, usually in kJ g⁻¹ or MJ kg⁻¹. It is found by dividing the energy released per mole of fuel (kJ mol⁻¹) by the molar mass (g mol⁻¹). A fuel that releases more energy per mole does not necessarily release more per gram: using the data 891 kJ mol⁻¹ for methane (M = 16.05 g mol⁻¹) and 5470 kJ mol⁻¹ for octane (M = 114.26 g mol⁻¹), methane releases 55.5 kJ g⁻¹ and octane 47.9 kJ g⁻¹.
Carbon dioxide added per unit energy
The mass (or amount) of CO₂ released per unit of energy obtained from a fuel, e.g. in g kJ⁻¹ or g per 1000 kJ; it is found from the mole ratio fuel : CO₂ in the balanced equation and the energy released per mole of fuel. It is the fair basis for comparing the contribution of different fuels to atmospheric CO₂. For fuels mainly composed of carbon and hydrogen, the higher the proportion of hydrogen, the less CO₂ per kJ: coal (modelled as carbon) releases the most and natural gas the least of the fossil fuels.
Greenhouse effect
The warming of the Earth's surface that occurs because greenhouse gases in the atmosphere, such as carbon dioxide, water vapour and methane, absorb infrared radiation emitted by the surface and re-emit it in all directions, some back towards the surface. Most incoming solar radiation (visible and ultraviolet) passes through these gases. The natural greenhouse effect keeps the surface habitable; increasing the CO₂ concentration by burning fossil fuels enhances it, contributing to global warming.
Advantages and disadvantages of fossil fuels
All three release a large amount of energy per unit mass and are readily available, but all are non-renewable and release CO₂. Coal: abundant and cheap, but releases the most CO₂ per unit energy, has the greatest tendency to burn incompletely (soot, CO) and often contains sulfur, forming SO₂. Crude oil: its liquid fractions are easily transported and are convenient fuels for vehicles, but it must be refined and spills pollute. Natural gas: releases the most energy per unit mass and the least CO₂ per unit energy and burns cleanly, but methane leaking unburnt is itself a greenhouse gas, and the gas is harder to store and transport.
Tendency to undergo incomplete combustion
How readily a fuel forms CO and soot when it burns. It increases with the proportion of carbon in the fuel and is greatest for solid fuels, whose oxygen supply is limited to the surface; gaseous fuels such as methane can mix thoroughly with air. Among fossil fuels the tendency is greatest for coal and least for natural gas, although any carbon-containing fuel burns incompletely if the oxygen supply is limited.
Students often think Natural gas is a clean fuel: it burns to water and produces no carbon dioxide, carbon monoxide or soot. In fact Yes. Methane contains carbon, so it forms CO₂ (CH₄ + 2O₂ → CO₂ + 2H₂O), and it forms CO and soot if the oxygen supply is limited. It releases less CO₂ per unit energy than coal or oil, but not zero.
Students often think The mass of CO₂ produced is equal to the mass of carbon in the fuel, since 'carbon emissions' and 'CO₂ emissions' mean the same thing. In fact No. Each carbon atom (12.01 g mol⁻¹) becomes a CO₂ molecule (44.01 g mol⁻¹), so the mass of CO₂ is 44.01/12.01 ≈ 3.7 times the mass of carbon burned.
R1.3.4 Renewable and non-renewable energy sources
Renewable and non-renewable energy sources
A renewable energy source is replenished by natural processes at least as fast as it is used (e.g. biofuels, solar, wind). A non-renewable source is used much faster than it is replenished, so its reserves are finite (e.g. coal, crude oil, natural gas, which formed over millions of years). The distinction concerns the rate of replenishment, not whether the source releases CO₂.
Biofuel
A fuel produced from recently living biomass, whose carbon was fixed from atmospheric carbon dioxide by photosynthesis over a short period of time (months to years). Examples: bioethanol, made by fermenting sugars from crops such as sugar cane; biodiesel, made from plant oils; biogas (mainly methane) from the anaerobic breakdown of organic waste.
Photosynthesis
The process in which green plants and other photosynthetic organisms use light energy to convert carbon dioxide and water into glucose and oxygen: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g). The reactants are carbon dioxide and water; the products are glucose and oxygen. It fixes atmospheric carbon into biomass, which is the source of the carbon in biofuels.
Advantages and disadvantages of biofuels
Advantages: renewable; the CO₂ released on burning was recently removed from the atmosphere by photosynthesis, so the net addition of CO₂ can be much lower than for fossil fuels; can be produced locally, including from waste. Disadvantages: not carbon neutral, because fossil fuels are used in farming (machinery, fertilizer), transport and processing; burning still releases CO₂ and can produce CO; crops for fuel compete with food crops for land and water, and clearing land (e.g. forest) releases stored carbon; bioethanol releases less energy per unit mass than petrol.
Students often think A renewable energy source is one that can be recycled or used again after it has released its energy. In fact No. A renewable energy source is replenished by natural processes at least as fast as it is used. A fuel's chemical energy is released once, when it burns; it cannot be reused.
Students often think Renewable fuels are carbon neutral: the CO₂ released when they burn is exactly balanced by the CO₂ the plants absorbed, so their net effect on atmospheric CO₂ is zero. In fact No. A renewable fuel such as bioethanol releases CO₂ when it burns, and fossil fuels used to grow, transport and process it add more. Whether a source is renewable depends on how fast it is replenished, not on its emissions.
R1.3.5 Fuel cell
Fuel cell
An electrochemical cell that converts chemical energy from a continuously supplied fuel (such as hydrogen or methanol) and oxygen directly into electrical energy, without burning the fuel. The fuel is oxidized at the anode and oxygen is reduced at the cathode; electrons flow through the external circuit and ions through the electrolyte. Unlike a primary or secondary cell, it runs as long as fuel and oxygen are supplied and is refuelled, not recharged.
Anode and cathode
In any electrochemical cell, oxidation (loss of electrons) occurs at the anode and reduction (gain of electrons) occurs at the cathode. In a fuel cell the fuel is oxidized at the anode, which is the negative electrode, and oxygen is reduced at the cathode, which is the positive electrode.
Half-equations in a hydrogen fuel cell (acidic electrolyte)
Anode (oxidation): H₂(g) → 2H⁺(aq) + 2e⁻. Cathode (reduction): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l). Overall: 2H₂(g) + O₂(g) → 2H₂O(l); the only product is water.
Half-equations in a methanol fuel cell (acidic electrolyte)
Anode (oxidation): CH₃OH(aq) + H₂O(l) → CO₂(g) + 6H⁺(aq) + 6e⁻; water supplies the extra oxygen atom needed to form CO₂. Cathode (reduction): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l), or 1½O₂ + 6H⁺ + 6e⁻ → 3H₂O to match the anode. Overall: CH₃OH(aq) + 1½O₂(g) → CO₂(g) + 2H₂O(l); unlike the hydrogen cell, CO₂ is produced.
Students often think Reduction occurs at the anode (the anode and cathode processes are the other way round). In fact At the cathode. Oxidation (loss of electrons) occurs at the anode, where the fuel reacts; reduction (gain of electrons) occurs at the cathode, where oxygen reacts.
Students often think Hydrogen is a non-metal, and non-metals gain electrons to form negative ions, so hydrogen forms H⁻ ions in the cell. In fact No. At the anode hydrogen loses electrons and forms H⁺ ions: H₂ → 2H⁺ + 2e⁻. Hydrogen is oxidized.
Diagnostic a bearings check, not a test
10 questions, one per part of the topic where we can. Answer them, then see which statements you own and which to read.
1 Which equation correctly represents the combustion of an element in excess oxygen?
Answer and reasoning
2Al(s) + O₂(g) → 2AlO(s) — A student who assumes a metal oxide is always one metal to one oxygen, like MgO, picks this. Aluminium forms Al³⁺ ions and oxide ions are O²⁻, so the formula is Al₂O₃: 4Al(s) + 3O₂(g) → 2Al₂O₃(s).
P₄(s) + 5O₂(g) → P₄O₁₀(s) — Phosphorus is a non-metal that burns in excess oxygen to form phosphorus(V) oxide, P₄O₁₀. Every formula is correct (oxygen as O₂) and the equation is balanced: 4 P and 10 O atoms on each side.
H₂(g) + O₂(g) → H₂O₂(l) — A student who balances the oxygen atoms by changing the product's formula to H₂O₂, which is hydrogen peroxide, a different substance, picks this. Hydrogen burns to form water; balance with coefficients: 2H₂(g) + O₂(g) → 2H₂O(l).
S(s) + 2O(g) → SO₂(g) — A student who writes oxygen as single O atoms picks this. Oxygen gas consists of O₂ molecules, so sulfur burns as S(s) + O₂(g) → SO₂(g).
2 Which equation represents an incomplete combustion of ethanol?
Answer and reasoning
C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l) — This is the equation for complete combustion. You may think 'incomplete' simply means that some of the fuel is left unburnt, with the same products. Incomplete combustion is defined by its products: carbon (soot) or carbon monoxide forms instead of carbon dioxide.
C₂H₅OH(l) + 1½O₂(g) → 2C(s) + 3H₂O(l) — A student who chooses the right products but takes all 3 of their O atoms from O₂ picks this. Ethanol already contains one O atom, so O₂ supplies only 2 O atoms: C₂H₅OH(l) + O₂(g) → 2C(s) + 3H₂O(l). As written, this equation has 4 O atoms on the left and 3 on the right.
C₂H₅OH(l) + 4O(g) → 2CO(g) + 3H₂O(l) — Carbon monoxide is a correct product of incomplete combustion, but you wrote oxygen as single O atoms. Oxygen reacts as O₂ molecules: C₂H₅OH(l) + 2O₂(g) → 2CO(g) + 3H₂O(l).
C₂H₅OH(l) + O₂(g) → 2C(s) + 3H₂O(l) — In a very limited supply of oxygen the carbon of ethanol is deposited as carbon (soot), while the hydrogen still forms water. The equation is balanced: the products contain 3 O atoms, 1 from ethanol and 2 from O₂.
3 Burning fossil fuels increases the concentration of carbon dioxide in the atmosphere. How does this increase lead to warming of the Earth's surface?
Answer and reasoning
CO₂ absorbs visible and ultraviolet light coming from the Sun, heating the air. — A student who places the absorption on the incoming sunlight picks this. CO₂ absorbs little visible or ultraviolet radiation; most sunlight passes through it. CO₂ absorbs the infrared radiation emitted by the warmed surface.
CO₂ absorbs infrared radiation emitted by the surface and re-emits some of it back. — Sunlight passes through the atmosphere and warms the surface, which emits infrared radiation. CO₂ molecules absorb this infrared and re-emit it in all directions, some back towards the surface, so more CO₂ means a warmer surface (the enhanced greenhouse effect).
CO₂ destroys the ozone layer, so more of the Sun's radiation reaches the surface. — A student who merges the greenhouse effect with ozone depletion picks this. CO₂ does not destroy ozone; ozone depletion (caused mainly by CFCs) concerns ultraviolet radiation. CO₂ warms the surface by absorbing and re-emitting infrared radiation.
CO₂ forms a layer high in the atmosphere that reflects the surface's heat back down. — A student who pictures CO₂ as a mirror-like layer picks this. CO₂ is mixed throughout the atmosphere, and its molecules absorb infrared radiation and re-emit it; they do not reflect it.
4 Biofuels are renewable energy sources, whereas fossil fuels are non-renewable. Which statement gives the defining feature of a renewable energy source?
Answer and reasoning
It can be collected and used again after it has released its energy. — A student who reads 'renewable' as 'reusable' or 'recyclable' picks this. The chemical energy of a fuel is released once, when it burns. Renewable means replenished by natural processes at least as fast as it is used.
It adds no carbon dioxide to the atmosphere overall when used. — A student who links 'renewable' to zero net emissions picks this. Renewable biofuels still add CO₂ overall, from the fossil fuels used to produce them, and a low-carbon source can be non-renewable. The definition concerns the rate of replenishment.
It comes from a natural source, such as the remains of plants or animals. — A student who takes natural or biological origin as the test picks this. Coal, oil and natural gas also come from the remains of organisms, but they are non-renewable because they formed over millions of years.
It is replenished naturally at least as fast as it is used. — Renewable is about the rate of replacement: new crops for biofuels grow within months or years, whereas coal, oil and gas took millions of years to form and are used far faster than they form.
5 In a hydrogen fuel cell with an acidic electrolyte, hydrogen gas is supplied to the anode. Which half-equation represents the reaction at the anode?
Answer and reasoning
H₂ → 2H⁺ + 2e⁻ — Oxidation takes place at the anode: each hydrogen molecule loses two electrons, forming two H⁺ ions that move through the acidic electrolyte to the cathode. The electrons flow through the external circuit.
2H⁺ + 2e⁻ → H₂ — This is a reduction, the reverse of the anode reaction. You placed reduction at the anode; reduction occurs at the cathode, and at the anode the fuel is oxidized: H₂ → 2H⁺ + 2e⁻.
H₂ + 2e⁻ → 2H⁻ — A student who applies 'non-metals gain electrons' to hydrogen picks this. In the fuel cell hydrogen is the fuel and is oxidized: it loses electrons and forms H⁺ ions, which the acidic electrolyte carries.
2H → 2H⁺ + 2e⁻ — The process is right but hydrogen is written as atoms. Hydrogen gas consists of H₂ molecules, so the half-equation is H₂ → 2H⁺ + 2e⁻.
6 Ethanol, C₂H₅OH, undergoes complete combustion in oxygen. What amount of oxygen, O₂, reacts with 1.00 mol of ethanol?
Answer and reasoning
3.00 mol — C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l). The products contain 2 × 2 + 3 × 1 = 7 O atoms; ethanol supplies 1 of them, so O₂ supplies 6 O atoms, which is 3 O₂. So 1.00 mol of ethanol reacts with 3.00 mol of O₂.
3.50 mol — A student who counts the 7 O atoms in the products (2CO₂ + 3H₂O) but takes them all from O₂, giving 3½O₂, picks this. Ethanol already contains one O atom in its –OH group, so O₂ supplies only 6 O atoms: 3.00 mol.
6.00 mol — A student who finds that O₂ must supply 6 O atoms and takes that as 6 mol of O₂ picks this. Oxygen is diatomic: each O₂ molecule supplies two O atoms, so 6 O atoms need 3 O₂, i.e. 3.00 mol.
1.50 mol — A student who balances only the carbon, forming 2CO₂ (4 O atoms, 3 of them from O₂), and leaves the hydrogen unoxidized picks this. The hydrogen in ethanol also burns, forming 3H₂O, which needs 3 more O atoms: 3.00 mol of O₂ in total.
Working C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l). O atoms in products = 2 × 2 + 3 × 1 = 7; O atoms in 1 C₂H₅OH = 1; O atoms from O₂ = 7 − 1 = 6; amount of O₂ = 6/2 = 3 per mol ethanol, so n(O₂) = 3.00 mol.
7 Methane burns in a limited supply of oxygen. Which equation represents an incomplete combustion of methane?
Answer and reasoning
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) — This is complete combustion. You may think incomplete combustion gives the same products with some methane left over. Incomplete combustion is defined by its products: carbon monoxide or carbon (soot) forms instead of CO₂.
CH₄(g) + 3O(g) → CO(g) + 2H₂O(l) — Carbon monoxide is a correct product of incomplete combustion, but you wrote oxygen as single O atoms. Oxygen reacts as O₂ molecules: 2CH₄(g) + 3O₂(g) → 2CO(g) + 4H₂O(l).
CH₄(g) + O₂(g) → C(s) + 2H₂O(l) — In a very limited supply of oxygen the carbon of methane is deposited as carbon (soot), while the hydrogen still forms water; 1 C, 4 H and 2 O atoms on each side.
2CH₄(g) + 3O₂(g) → 2CO(g) + H₈O₄(l) — A student who makes the atoms balance by changing the formula of water to H₈O₄ picks this. A formula identifies the substance and cannot be changed; balance with the numbers in front of formulae instead: 2CH₄(g) + 3O₂(g) → 2CO(g) + 4H₂O(l).
8 Coal has a greater tendency than natural gas to undergo incomplete combustion. Which statement gives a reason for this?
Answer and reasoning
Coal contains impurities, and these, not its carbon, form the soot and carbon monoxide. — A student who attributes soot to impurities in the coal picks this. Soot is carbon and carbon monoxide contains carbon; both come from the fuel's own carbon when oxygen is limited. Impurities such as sulfur form other pollutants, such as SO₂.
Coal needs more oxygen per gram than methane, as only the carbon in a fuel uses up oxygen. — A student who counts only the carbon as needing oxygen picks this. Hydrogen is oxidized to water too, so methane needs more oxygen per gram: 4.0 g of O₂ per gram of CH₄ compared with 2.7 g per gram of carbon. Coal burns incompletely because it is a solid that mixes poorly with air.
Methane is a clean fuel, so its carbon cannot form carbon monoxide or soot. — A student who treats natural gas as a fuel with no carbon emissions picks this. Methane contains carbon and forms CO or soot whenever its oxygen supply is limited, for example in a Bunsen flame with the air hole closed; it is simply less prone to this than coal.
Coal is a solid, so oxygen reaches only its surface, whereas gas mixes fully with air. — A gaseous fuel such as methane can mix thoroughly with air before it burns, so every molecule has oxygen nearby. Solid coal, which is mostly carbon, reacts only where oxygen reaches its surface, so the oxygen supply to the carbon is easily limited and CO and soot form.
9 Bioethanol is made by fermenting sugar from sugar cane and is burned as a fuel for vehicles. Which statement about bioethanol as a fuel is correct?
Answer and reasoning
It burns to form water only, so using it as a fuel adds no CO₂ to the atmosphere. — A student who assumes a 'green' fuel produces no CO₂ picks this. Bioethanol is ethanol: C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O. Its possible advantage lies in where its carbon came from, not in its combustion products.
The CO₂ it releases is exactly balanced by photosynthesis, so its net emission is zero. — A student who applies the simple 'plants absorb it, burning returns it' cycle picks this. That cycle leaves out the fossil fuels used in farming, fertilizer production, transport, fermentation and distillation, so the net emission is lower than for petrol but not zero.
All of its carbon was fixed from atmospheric CO₂ by photosynthesis in the cane. — The cane fixed its carbon from atmospheric CO₂ by photosynthesis (6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂) within the last year or two, so burning the bioethanol returns recently removed carbon. Even so, fossil fuels used in farming, fertilizer manufacture, transport and processing add further CO₂, so bioethanol is not carbon neutral.
Its carbon was taken from the soil by the roots of the cane, not from CO₂ in the air. — A student who places the source of plant carbon in the soil picks this. Roots take up water and mineral ions; the carbon comes from atmospheric CO₂ through photosynthesis, 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂.
10 In a methanol fuel cell with an acidic electrolyte, an aqueous solution of methanol is supplied to the anode. Which equation represents the reaction that occurs at the anode?
Answer and reasoning
2CH₃OH(aq) + O₂(g) → 2CO₂(g) + 8H⁺(aq) + 8e⁻ — A student who reasons that oxidation means reaction with oxygen, and so puts O₂ at the anode, picks this. Oxygen is supplied to the cathode, where it is reduced; at the anode methanol loses electrons and water supplies the extra O atom: CH₃OH + H₂O → CO₂ + 6H⁺ + 6e⁻.
CH₃OH(aq) + H₂O(l) → CO₂(g) + 6H⁺(aq) + 6e⁻ — Methanol is oxidized at the anode. Water supplies the second O atom needed for CO₂; balancing H gives 6H⁺ (4 from CH₃OH, 2 from H₂O), and balancing charge gives 6e⁻. The electrons flow to the cathode through the external circuit.
2CH₃OH(aq) + 3O₂(g) → 2CO₂(g) + 4H₂O(l) — This is the overall (combustion) equation. You pictured the cell as burning the methanol. In a fuel cell the fuel and oxygen react at separate electrodes; the anode half-equation shows the electrons released: CH₃OH + H₂O → CO₂ + 6H⁺ + 6e⁻.
CO₂(g) + 6H⁺(aq) + 6e⁻ → CH₃OH(l) + H₂O(l) — This is a reduction: electrons are gained and methanol is formed. You placed reduction at the anode. Reduction occurs at the cathode, where oxygen is reduced; at the anode methanol is oxidized: CH₃OH + H₂O → CO₂ + 6H⁺ + 6e⁻.
Read the ones marked not yet in Learn, then Verify.
Verify confirm before you go
6 more questions. Every wrong answer here is a real misconception, and you see why it is wrong straight away.
1 Octane, C₈H₁₈, is a component of petrol, which is refined from crude oil. Use the following data: 5470 kJ of energy is released when 1 mol of octane burns completely; A_r: H 1.01, C 12.01, O 16.00. What mass of CO₂ is added to the atmosphere for every 1.00 × 10³ kJ of energy released by burning octane?
Answer and reasoning
64.4 g — 2C₈H₁₈ + 25O₂ → 16CO₂ + 18H₂O, so 1 mol of octane gives 8 mol of CO₂ (8 × 44.01 = 352.1 g) and releases 5470 kJ. Per 1.00 × 10³ kJ: 352.1 × 1000/5470 = 64.4 g of CO₂.
8.05 g — A student who takes one mole of CO₂ per mole of octane picks this: 44.01 × 1000/5470 = 8.05 g. Each octane molecule contains 8 carbon atoms, so it forms 8 CO₂ molecules; the answer is 8 times larger, 64.4 g.
20.9 g — A student who takes the mass of CO₂ to be equal to the mass of octane burned, 114.26 × 1000/5470 = 20.9 g, as if the fuel simply turned into CO₂, picks this. The CO₂ also contains oxygen atoms taken from the air, so its mass is greater: 8 × 44.01 × 1000/5470 = 64.4 g.
17.6 g — A student who calculates the mass of carbon burned, 8 × 12.01 × 1000/5470 = 17.6 g, picks this. Each carbon atom leaves as a CO₂ molecule of molar mass 44.01 g mol⁻¹, so the mass of CO₂ is 8 × 44.01 × 1000/5470 = 64.4 g.
Working M(CO₂) = 12.01 + 2 × 16.00 = 44.01 g mol⁻¹. 2C₈H₁₈ + 25O₂ → 16CO₂ + 18H₂O, so n(CO₂) : n(C₈H₁₈) = 8 : 1. Per 1 mol octane: m(CO₂) = 8 × 44.01 = 352.08 g for 5470 kJ. Per 1.00 × 10³ kJ: 352.08 × (1000/5470) = 64.37 g ≈ 64.4 g.
2 Use the following data for the energy released when 1 mol of each fuel burns completely: coal (modelled as pure carbon, C) 394 kJ; natural gas (methane, CH₄) 891 kJ; petrol (modelled as octane, C₈H₁₈) 5470 kJ. A_r: H 1.01, C 12.01, O 16.00. Which conclusion is supported by these data?
Answer and reasoning
Octane releases the most energy per gram of the three fuels. — A student who compares the energy released per mole, where octane's 5470 kJ is the largest, picks this. Per gram, the large molar mass of octane (114.26 g mol⁻¹) must be divided in: 47.9 kJ g⁻¹, less than methane's 55.5 kJ g⁻¹.
Octane adds the least CO₂ per kJ of energy released of the three. — A student who assumes one mole of CO₂ per mole of fuel, which makes octane's per-kJ figure look smallest, picks this. Octane forms 8 mol of CO₂ per mole: 8 × 44.01/5470 = 0.0644 g kJ⁻¹, more than methane's 44.01/891 = 0.0494 g kJ⁻¹. Coal gives 44.01/394 = 0.112 g kJ⁻¹.
Methane releases the most energy per gram of the three fuels. — Energy per gram = energy per mole ÷ molar mass: carbon 394/12.01 = 32.8 kJ g⁻¹; methane 891/16.05 = 55.5 kJ g⁻¹; octane 5470/114.26 = 47.9 kJ g⁻¹. Methane releases the most energy per gram.
Methane adds no CO₂ to the atmosphere when it is burned as a fuel. — A student who treats natural gas as a fuel with no carbon emissions picks this. Methane contains carbon: CH₄ + 2O₂ → CO₂ + 2H₂O. It releases the least CO₂ per kJ of the three fuels (0.0494 g kJ⁻¹), but not zero.
Working M: C 12.01; CH₄ 12.01 + 4 × 1.01 = 16.05; C₈H₁₈ 8 × 12.01 + 18 × 1.01 = 114.26 g mol⁻¹. Energy per gram: C 394/12.01 = 32.8; CH₄ 891/16.05 = 55.5; C₈H₁₈ 5470/114.26 = 47.9 kJ g⁻¹. CO₂ per kJ (M(CO₂) = 44.01 g mol⁻¹): C 1 × 44.01/394 = 0.112; CH₄ 1 × 44.01/891 = 0.0494; C₈H₁₈ 8 × 44.01/5470 = 0.0644 g kJ⁻¹. Methane releases the most energy per gram and the least CO₂ per kJ.
3 Plants such as sugar cane, grown to make bioethanol, carry out photosynthesis. Which statement gives the reactants and products of photosynthesis?
Answer and reasoning
Glucose reacts with oxygen to form carbon dioxide and water. — This is aerobic respiration, the overall reverse of photosynthesis; you interchanged the two processes. In photosynthesis carbon dioxide and water are the reactants and glucose and oxygen are the products.
Water and soil minerals react to form glucose and oxygen. — A student who places the plant's carbon in the soil picks this. Soil minerals supply elements such as nitrogen, not carbon; the carbon in glucose comes from CO₂ in the air, which reacts with water to form glucose and oxygen.
Carbon dioxide and water react to form glucose and oxygen gas. — Photosynthesis uses light energy to convert carbon dioxide and water into glucose and oxygen: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g). The carbon in glucose, and so in a biofuel made from it, comes from atmospheric CO₂.
Carbon dioxide reacts to form carbon and oxygen gas. — A student who pictures photosynthesis as splitting CO₂, with the plant keeping the carbon, picks this. The carbon becomes part of glucose, C₆H₁₂O₆, whose hydrogen comes from water, a reactant. The products are glucose and oxygen.
4 When the equation for the complete combustion of octane, C₈H₁₈, is balanced with a coefficient of 2 for octane, what is the coefficient of O₂?
Answer and reasoning
25 — 2C₈H₁₈ forms 16CO₂ and 18H₂O, which contain 16 × 2 + 18 = 50 O atoms. Each O₂ supplies two O atoms, so 25 O₂ are needed: 2C₈H₁₈(l) + 25O₂(g) → 16CO₂(g) + 18H₂O(l).
50 — 50 is the number of O atoms in the products (16 × 2 + 18). You took each O atom as one oxygen particle. Oxygen reacts as O₂ molecules, each supplying two O atoms, so 50/2 = 25 O₂.
16 — A student who oxidizes only the carbon picks this: 16CO₂ needs 32 O atoms, i.e. 16 O₂, and the hydrogen is left as H₂. Hydrogen also burns, to 18H₂O, which needs 18 more O atoms: (32 + 18)/2 = 25 O₂.
11 — A student who writes one CO₂ per octane molecule (2CO₂) with 18H₂O, giving (4 + 18)/2 = 11 O₂, picks this. Each octane molecule contains 8 carbon atoms and forms 8 CO₂, so 2C₈H₁₈ gives 16CO₂ and needs (32 + 18)/2 = 25 O₂.
Working Complete combustion: all carbon → CO₂, all hydrogen → H₂O. 2C₈H₁₈ contains 16 C and 36 H atoms, so the products are 16CO₂ + 18H₂O. O atoms in the products = 16 × 2 + 18 × 1 = 50, supplied by 50/2 = 25 O₂: 2C₈H₁₈(l) + 25O₂(g) → 16CO₂(g) + 18H₂O(l). (For one octane molecule the coefficient would be 12½, which is why the equation is written for two.)
5 Use the following data for the energy released when 1 mol of each fuel burns completely: coal (modelled as pure carbon, C) 394 kJ; petrol (modelled as octane, C₈H₁₈) 5470 kJ. A_r: H 1.01, C 12.01, O 16.00. How many times greater is the mass of CO₂ added to the atmosphere per kJ of energy released when coal burns than when octane burns?
Answer and reasoning
13.9 — A student who takes one mole of CO₂ per mole of each fuel, so that the ratio becomes 5470/394 = 13.9, picks this. Each octane molecule forms 8 CO₂, so octane gives 352.08 g of CO₂ per 5470 kJ (0.0644 g kJ⁻¹), and the ratio is 0.1117/0.0644 = 1.74.
1.74 — Coal: 44.01 g of CO₂ per 394 kJ = 0.1117 g kJ⁻¹. Octane: 8 × 44.01 = 352.08 g of CO₂ per 5470 kJ = 0.06437 g kJ⁻¹. Ratio 0.1117/0.06437 = 1.74. Coal adds more CO₂ per kJ because none of its energy comes from oxidizing hydrogen to water, which forms no CO₂.
1.46 — A student who uses the mass of fuel burned per kJ (coal 12.01/394, octane 114.26/5470) as the mass of CO₂, as if each fuel simply turned into CO₂, picks this. The CO₂ also contains oxygen from the air: 12.01 g of carbon forms 44.01 g of CO₂. With the CO₂ masses the ratio is 1.74.
1.19 — A student who compares the CO₂ formed per gram of fuel (coal 44.01/12.01 = 3.66 g; octane 352.08/114.26 = 3.08 g) picks this. The fuels release different energies per gram (32.8 and 47.9 kJ g⁻¹), so compare per kJ: 0.1117/0.06437 = 1.74.
Working M(CO₂) = 12.01 + 2 × 16.00 = 44.01 g mol⁻¹. Coal: C + O₂ → CO₂, 1 mol CO₂ per 394 kJ: 44.01/394 = 0.1117 g kJ⁻¹. Octane: 2C₈H₁₈ + 25O₂ → 16CO₂ + 18H₂O, 8 mol CO₂ per 5470 kJ: 8 × 44.01/5470 = 352.08/5470 = 0.06437 g kJ⁻¹. Ratio = 0.1117/0.06437 = 1.74 (no unit).
6 In a hydrogen fuel cell with an acidic electrolyte, which equation represents the reaction that occurs at the cathode?
Answer and reasoning
4H₂O(l) → 2O₂(g) + 8H⁺(aq) + 8e⁻ — This is an oxidation (electrons are lost and oxygen gas is formed), the reverse of the cathode reaction. You placed oxidation at the cathode. Oxidation occurs at the anode, where hydrogen is oxidized; at the cathode oxygen is reduced: O₂ + 4H⁺ + 4e⁻ → 2H₂O.
O(g) + 2H⁺(aq) + 2e⁻ → H₂O(l) — The process is right, but oxygen is written as single atoms. Oxygen gas consists of O₂ molecules, so the half-equation is O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l).
H₂(g) + ½O₂(g) → H₂O(l) + heat — This is the equation for burning hydrogen. You pictured the cell as burning its fuel and using the heat. In a fuel cell hydrogen and oxygen react at separate electrodes and no heat-to-electricity step is involved; at the cathode oxygen is reduced: O₂ + 4H⁺ + 4e⁻ → 2H₂O.
O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) — Reduction takes place at the cathode: each O₂ molecule gains four electrons and, with four H⁺ ions from the acidic electrolyte, forms two water molecules. O, H and charge all balance.
That was your twenty minutes. Real practice on R1.3 is past-paper questions marked against the mark scheme.
Paper 1A tests it as multiple choice; Paper 1B through data you have not seen before; Paper 2 with short answers and, at HL, extended responses. Look for the command words — outline, explain, compare, evaluate — and give exactly what each asks for.
Compiled from the IB Chemistry guide (first assessment 2025) and our question bank · Specialist review in progress. How these pages are made · Free, no account ·
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