Topic 12 · AS Level · Inorganic chemistry

What this chapter covers12.1

Nitrogen makes up almost four-fifths of the air, and for most purposes it does nothing at all. You breathe it in and out unchanged; food is packed under it because it will not react with anything in the packet. Yet the same element, once it has been forced to react, turns up in fertilisers, in the brown haze over a city at rush hour and in the acid that eats limestone buildings. This chapter is about that contrast: why the N₂ molecule is so reluctant to react, what happens to nitrogen once it is part of ammonia or an oxide, and how the oxides of nitrogen made in car engines and lightning go on to cause photochemical smog and acid rain — including the part they play in turning sulfur dioxide into sulfuric acid.

Read the verbs, because they decide what an answer must contain. 12.1.1 says explain, and names the two reasons it wants: the strength of the triple bond and the lack of polarity of the molecule. "Nitrogen has a triple bond" is not an explanation until it says why the triple bond matters. 12.1.2 says describe and explain three times and each time names the theory or the reaction type — Brønsted–Lowry, acid–base, acid–base — so every answer in that statement is an argument about protons. 12.1.3 has two halves, where oxides of nitrogen come from and how they are removed. 12.1.5 says both directly and catalytically, which is a signal that an answer naming only one of the two roles is half an answer.

The page is in three parts. Part one is the nitrogen molecule and why it is unreactive. Part two is ammonia and the ammonium ion, seen entirely through the Brønsted–Lowry theory. Part three follows the oxides of nitrogen from the engine to the atmosphere. Each part ends with examination practice taken from past papers, and the final part collects what examiners have reported about this topic, the mistakes that lose marks, a self-test and a knowledge organiser.

The nitrogen molecule12.1.1

A nitrogen atom has the electron configuration 1s² 2s² 2p³: five outer electrons, three of them unpaired. Two nitrogen atoms share three pairs of electrons, which gives each atom an octet — three bonding pairs and one lone pair — and gives the molecule its triple bond.

:N≡N:

A triple bond is not three identical bonds. It is one σ bond and two π bonds. Topic 13.3 describes the atom forming it as sp hybridised: the 2s orbital mixes with one 2p orbital to give two sp hybrid orbitals pointing in opposite directions along the axis of the molecule. One sp orbital on each atom overlaps end-on (head-on) with its partner to form the σ bond; the other holds the lone pair and points away from the bond. That leaves two unhybridised 2p orbitals on each atom, at right angles to each other and to the bond axis. Each overlaps sideways with the matching p orbital on the other atom, giving two π bonds — one with its electron density above and below the axis, the other in front of and behind it (Figure 12.1).

This arrangement puts six bonding electrons between two small nuclei, and the atoms are pulled closer than in any other nitrogen–nitrogen bond. The numbers show how much the third pair of electrons adds.

The triple bond is almost six times as strong as the single bond — far more than three single bonds added together. It is one of the strongest covalent bonds known, and that single number drives most of part one.

Why nitrogen is so unreactive12.1.1

12.1.1 names two reasons, and a complete answer uses both.

Reason one: the triple bond is very strong

Almost every reaction of nitrogen begins by breaking, or at least weakening, the N≡N bond. With a bond energy of 944 kJ mol⁻¹ this costs a great deal of energy, so the activation energy of the reaction is very high. At room temperature, only a vanishingly small fraction of collisions have enough energy to reach it, so the reaction is too slow to observe.

Notice what this argument is not saying. It is not saying that reactions of nitrogen are endothermic. Some are — making nitrogen monoxide from nitrogen and oxygen absorbs energy — but others are exothermic overall: forming ammonia from nitrogen and hydrogen releases energy, and magnesium burning in nitrogen releases a great deal. Nitrogen is unreactive in these reactions for a kinetic reason, not an energetic one. The energy released by forming new bonds only arrives after the old one has been broken, and the barrier is set by the bond that has to break first (Figure 12.2).

Reason two: the molecule has no polarity

Most reactions start with an attraction. A nucleophile is drawn to an atom carrying a partial positive charge; an electrophile is drawn to a region of high electron density or to an atom carrying a partial negative charge. In N₂ the two atoms are identical, so they have the same electronegativity and share the bonding electrons exactly equally. There is no permanent dipole — no δ+ end and no δ− end — so there is nothing for an attacking nucleophile or electrophile to be attracted to. The molecule is also small, with its electrons held tightly, so it is not easily polarised by an approaching species either.

What it takes to make nitrogen react12.1.1

Unreactive is not the same as inert. Nitrogen does react when enough energy is supplied, and the conditions in which it reacts tell the same story as section 3: every one of them is a way of getting past a very high activation energy.

In each case the temperature is high because enough molecules must have an energy at least equal to the activation energy (topic 8.2), or a catalyst provides a route with a lower activation energy (8.3). Nitrogen in the air around you, at around 290 K and with nothing to catalyse its reactions, simply has no way over the barrier.

Ammonia as a Brønsted–Lowry base12.1.2(a)

Ammonia is a colourless gas with a sharp, choking smell, and it is extremely soluble in water — a flask of ammonia gas will suck water up into it in a fountain. The solution turns red litmus blue and has a pH of about 11. Something in the solution is producing hydroxide ions, yet ammonia contains no oxygen at all. The Brønsted–Lowry theory explains where they come from.

To accept a proton, a species needs a lone pair of electrons with which to form a new bond to it: the proton has no electrons of its own to share. The nitrogen atom in NH₃ has exactly that — three bonding pairs to hydrogen and one lone pair. When ammonia dissolves, some of its molecules use that lone pair to take a proton from a water molecule:

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

Water has given up a proton, so here water is acting as the acid; ammonia has accepted it, so ammonia is the base. The hydroxide ion left behind is what makes the solution alkaline. The reaction is reversible, and the reverse reaction is also a proton transfer — NH₄⁺ giving its proton back to OH⁻ — so the two sides contain two conjugate acid–base pairs:

The members of a conjugate pair differ by exactly one H⁺. That is the test to apply to any equation: look for two species whose formulae differ by one hydrogen and one unit of charge.

Ammonia reacts in the same way with acids, but much more completely, because an acid such as hydrochloric acid is a far better proton donor than water. With hydrogen chloride gas, the product is a white smoke of solid ammonium chloride — the classic test for either gas:

NH₃(g) + HCl(g) → NH₄Cl(s)

and in solution the reaction is between ammonia and the oxonium ion, H₃O⁺, which is how a proton exists in water:

NH₃(aq) + H₃O⁺(aq) → NH₄⁺(aq) + H₂O(l)

In each case the product contains the ammonium ion, and in each case the reaction is a Brønsted–Lowry acid–base reaction: a proton moves from the acid to the lone pair on nitrogen.

How weak a base is ammonia?12.1.2(a)

The equilibrium in aqueous ammonia lies well to the left. In a typical bench solution only about one ammonia molecule in a hundred has accepted a proton at any moment; the rest are still NH₃ molecules, hydrogen-bonded to water. That is what makes ammonia a weak base: it is only partially converted into ions in solution. Sodium hydroxide, by contrast, is a strong base: it is completely ionic, so every formula unit supplies a hydroxide ion.

The difference shows up in the pH. The model below solves the equilibrium exactly for any concentration you choose, using the base dissociation constant of ammonia, and sets the result beside sodium hydroxide of the same concentration.

Forming the ammonium ion12.1.2(b)

When an ammonia molecule accepts a proton, the new N–H bond is formed from the nitrogen's lone pair. Both electrons in that bond come from the same atom, so it is a dative covalent (co-ordinate) bond. The proton has nothing to contribute but its positive charge, and the ion as a whole carries that charge.

NH₃ + H⁺ → NH₄⁺

Once the ion has formed, the dative bond cannot be told apart from the other three. All four N–H bonds are identical in length and strength, because once a pair of electrons is shared it makes no difference which atom it came from. The positive charge is not sitting on the new hydrogen: it belongs to the ion as a whole.

The shape of the ammonium ion12.1.2(b)

Electron-pair repulsion (topic 3.5) predicts both shapes. Nitrogen in ammonia has four pairs of electrons in its outer shell — three bonding pairs and one lone pair. Four pairs arrange themselves tetrahedrally, but only three of them have atoms at the end, so the molecule is trigonal pyramidal. The lone pair is held closer to the nitrogen than the bonding pairs are and repels them more strongly, squeezing the H–N–H angle down from 109.5° to about 107°.

Protonation turns the lone pair into a bonding pair. The ammonium ion has four bonding pairs and no lone pairs, all repelling equally, so it is a regular tetrahedron with every H–N–H angle 109.5° — the same shape as methane, with which it is isoelectronic. Rotate both below.

The ammonium ion as a weak acid12.1.2

The ammonium ion is the conjugate acid of ammonia, and it can give its proton away again. In water, a small proportion of ammonium ions donate a proton to water molecules:

NH₄⁺(aq) + H₂O(l) ⇌ NH₃(aq) + H₃O⁺(aq)

so a solution of ammonium chloride is slightly acidic, with a pH a little above 5 at 0.1 mol dm⁻³. The chloride ion plays no part: it is the conjugate base of a strong acid and has no tendency to take a proton back. The ammonium ion is a weak acid because only a very small fraction of the ions donate their proton — the position of equilibrium lies far to the left.

Displacing ammonia from ammonium salts12.1.2(c)

Warm any ammonium salt with a solution of sodium hydroxide and a sharp smell appears almost at once: ammonia gas is given off. A piece of damp red litmus paper held at the mouth of the tube turns blue. Ammonia is the only common gas that is alkaline, so this is the standard test for the ammonium ion.

Why it happens: an acid–base reaction

12.1.2(c) says by an acid–base reaction, and that is the key to the explanation. The ammonium ion is a proton donor; the hydroxide ion is a much stronger proton acceptor than ammonia. The hydroxide ion takes a proton from the ammonium ion:

NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)

This is the reverse of the reaction in section 5, and it goes almost to completion for two reasons: hydroxide is a stronger base than ammonia, and warming drives the ammonia out of solution as a gas, so it is continuously removed from the equilibrium. The same ionic equation describes every ammonium salt with every soluble hydroxide, because the anion of the salt and the cation of the base are spectator ions. A basic oxide does the same job, since the oxide ion is an even stronger base than hydroxide:

2NH₄⁺ + O²⁻ → 2NH₃ + H₂O

The full equation depends on the salt and the base. The model builds it for any pair, balances it, and gives the ionic equation that sits underneath.

Proton transfers, practised12.1.2

Every reaction in this part of the chapter is one proton moving from one species to another. The trainer gives you eight equations, several of them unfamiliar. For each, pick the species that acts as the Brønsted–Lowry acid; the page then identifies both conjugate pairs by finding the two pairs of species that differ by one H⁺. The answer is worked out from the formulae, not stored.

Examination practice: nitrogen and ammonia12.1.1–12.1.2

These questions are taken from past examination papers. Attempt each one before opening its worked answer. Where the examiners' report commented on a question, the insight is given with the answer. Structured-question answers are worked solutions that show the chemistry each mark needs.

Multiple choice

Structured questions

Where oxides of nitrogen come from12.1.3

Two oxides of nitrogen matter in this topic. Nitrogen monoxide, NO (also called nitrogen(II) oxide or nitric oxide), is a colourless gas. Nitrogen dioxide, NO₂ (nitrogen(IV) oxide), is a brown, toxic gas and is the colour of the haze over a polluted city. Together they are often written NO_x.

Both begin with the reaction section 3 said was so difficult: nitrogen and oxygen from the air combining.

N₂(g) + O₂(g) → 2NO(g)     ΔH = +181 kJ mol⁻¹

The reaction happens only at very high temperatures. There are two important places where air is heated that much.

Two points in the engine case are worth making explicit, because examiners report they are often missed. The nitrogen does not come from the fuel: petrol is a mixture of hydrocarbons. It comes from the air that is burned with the fuel. And the reaction does not need a catalyst: it is the high temperature in the cylinder that does it.

Once in the atmosphere, nitrogen monoxide is oxidised further by oxygen to nitrogen dioxide. This reaction is slow at the concentrations found in air, but it happens without any special conditions.

2NO(g) + O₂(g) → 2NO₂(g)

Why high temperature? A model of the equilibrium

The formation of NO is endothermic and, strictly, reversible. The model below computes how much NO air would contain if it were held at a given temperature until equilibrium was reached. It uses ideas from A Level topic 23 (entropy and Gibbs free energy) that you do not need for this topic; what matters here is the result, which shows why only lightning and combustion make significant amounts.

Removing NO_x: the catalytic converter12.1.3

The exhaust from a petrol engine contains three pollutants that a catalytic converter is designed to remove: carbon monoxide from incomplete combustion, oxides of nitrogen from the reaction in section 13, and unburnt hydrocarbons. The converter is a stainless-steel canister in the exhaust pipe, containing a ceramic honeycomb coated with a very thin layer of precious metals — platinum, palladium and rhodium. The honeycomb gives the metal a very large surface area for a very small mass of expensive catalyst, so the gases pass over as much catalyst as possible in the short time they spend in the converter.

The central reaction for 12.1.3 is a redox reaction between two of the pollutants. Carbon monoxide is a reducing agent and nitrogen monoxide is an oxidising agent, so on the catalyst surface each gets rid of the other:

2CO(g) + 2NO(g) → 2CO₂(g) + N₂(g)

Carbon is oxidised from +2 to +4; nitrogen is reduced from +2 to 0. Both products are harmless components of air. Nitrogen dioxide is removed in the same way, and the unburnt hydrocarbons are oxidised too, either by the oxides of nitrogen or by oxygen:

4CO(g) + 2NO₂(g) → 4CO₂(g) + N₂(g)

C₈H₁₈(g) + 25NO(g) → 8CO₂(g) + 9H₂O(g) + 12½N₂(g)

These equations are not to be memorised one by one. Every one of them is built the same way: the electrons lost by the carbon-containing species equal the electrons gained by the nitrogen oxide. The model does this for any pair.

Photochemical smog and PAN12.1.4

In a sunny city with heavy traffic, the air can fill with a brownish haze that irritates the eyes and lungs: photochemical smog. The word "photochemical" says that sunlight drives the reactions that make it. It forms from two groups of primary pollutants released from vehicle exhausts — oxides of nitrogen and unburnt hydrocarbons (often called volatile organic compounds, VOCs) — reacting in sunlight with oxygen to form secondary pollutants. One of those secondary pollutants is named by the syllabus: peroxyacetyl nitrate, PAN.

The name can be read off the structure: acetyl is the old name for the ethanoyl group, CH₃CO–, the one found in ethanoic (acetic) acid; peroxy refers to the –O–O– link; and the –NO₂ group on the end, joined through an oxygen atom, makes the nitrate part. The acetyl part comes from the hydrocarbon, and the nitrogen and extra oxygen come from nitrogen dioxide and the air. That is the whole of what 12.1.4 asks you to understand:

NO / NO₂ + unburnt hydrocarbons + O₂  —(sunlight)→  PAN (and other secondary pollutants, including ozone)

No equation is required. The reactions that make PAN involve free radicals and many steps, and the syllabus deliberately asks only that you understand that oxides of nitrogen react with unburned hydrocarbons to form PAN and that PAN is a component of photochemical smog. PAN is harmful because it irritates the eyes and respiratory system and damages plants.

Acid rain: the direct role of NO and NO₂12.1.5

Rain is always slightly acidic, because carbon dioxide dissolves in it. Acid rain is rain made considerably more acidic by oxides of sulfur and nitrogen. Its effects include damage to limestone and marble buildings and statues, acidification of lakes and rivers so that fish and other organisms die, damage to trees and forests, and leaching of nutrients from soils.

Oxides of nitrogen contribute to it directly, by forming nitric acid. Nitrogen monoxide is first oxidised to nitrogen dioxide (section 13), and nitrogen dioxide then reacts with water and more oxygen:

4NO₂(g) + O₂(g) + 2H₂O(l) → 4HNO₃(aq)

Nitrogen dioxide also reacts with water on its own, giving a mixture of two acids — nitric acid and nitrous acid, HNO₂:

2NO₂(g) + H₂O(l) → HNO₃(aq) + HNO₂(aq)

The second equation is a disproportionation: nitrogen starts at +4 and ends at +5 in HNO₃ and +3 in HNO₂. Examiners have reported that very few answers mention both acids, and that HNO₂ is rarely seen.

Where the sulfur dioxide comes from

Fossil fuels — coal especially, and crude oil before it is refined — contain sulfur compounds as impurities. When the fuel is burned, the sulfur is oxidised to sulfur dioxide:

S + O₂ → SO₂

Sulfur dioxide dissolves in rain to form sulfurous acid (sulfuric(IV) acid). No oxidation number changes in this step — sulfur stays at +4 — so it is an acid–base reaction, not a redox one:

SO₂(g) + H₂O(l) → H₂SO₃(aq)

But the main acid in acid rain is sulfuric acid, H₂SO₄, in which sulfur is +6. Getting from +4 to +6 means oxidising SO₂ to SO₃, and that is the reaction in which the oxides of nitrogen play their second, catalytic role.

NO₂ as a catalyst for oxidising SO₂12.1.5

Sulfur dioxide reacts with oxygen only very slowly in clean air. Nitrogen dioxide speeds this up by a two-step cycle. In the first step it oxidises sulfur dioxide to sulfur trioxide, and is itself reduced to nitrogen monoxide:

SO₂(g) + NO₂(g) → SO₃(g) + NO(g)

In the second step, oxygen in the air oxidises the nitrogen monoxide back to nitrogen dioxide:

NO(g) + ½O₂(g) → NO₂(g)    or    2NO(g) + O₂(g) → 2NO₂(g)

Add the two steps and the NO₂ and NO cancel: the nitrogen dioxide used in step 1 is regenerated in step 2. The overall change is simply

SO₂(g) + ½O₂(g) → SO₃(g)

and the sulfur trioxide then reacts with water in the rain:

SO₃(g) + H₂O(l) → H₂SO₄(aq)

Nitrogen dioxide meets every part of the definition of a catalyst. It increases the rate of the oxidation of sulfur dioxide, by providing an alternative route with a lower activation energy; it takes part in the reaction; and it is regenerated, so it is not used up and one molecule can oxidise many molecules of sulfur dioxide. Watch that happen in the model.

Nitrogen's oxidation numbers12.1

Every nitrogen species in this chapter can be placed on a single scale. Nitrogen has five outer electrons, so its oxidation number runs from −3, when it has in effect gained three (as in ammonia), to +5, when it has in effect lost all five (as in nitric acid). Seeing them together makes the redox chemistry of part three easy to check: in the catalytic converter nitrogen moves down the scale, in the atmosphere it moves up, and in the ammonia chemistry of part two it does not move at all — which is exactly why those reactions are acid–base reactions rather than redox.

Examination practice: oxides of nitrogen12.1.3–12.1.5

These questions are taken from past examination papers. Attempt each one before opening its worked answer; where the examiners' report commented on a question, the insight is given with the answer.

Multiple choice

Structured questions

Examiner's overall observation12

Examiners' reports on the structured AS papers have commented on this topic in most recent sessions. Taken together, they give a consistent picture.

What candidates generally do well. Lightning is well known as the natural source of oxides of nitrogen, and the catalytic converter is readily named as the device that removes them. The consequences of acid rain, and the equations for sulfuric acid formation, are usually known. The dot-and-cross diagram of the ammonium ion, with three covalent bonds and one dative bond, is usually correct, as are its tetrahedral shape and 109.5° bond angle. Recall of VOCs and PAN as the ingredients and product of photochemical smog has been good, and in the multiple-choice papers questions on what a catalytic converter removes, and on which pollutant contributes to acid rain, were among those candidates found easiest.

Where marks are lost on nitrogen itself. "Triple bond" appears in almost every answer on the unreactivity of nitrogen, but often without the word strong or any reference to the energy needed to break it. Examiners have pointed out that "strong triple bond" and "high activation energy" are in effect one point, so a two-mark answer needs a different second idea — the lack of polarity. Questions on how the triple bond forms expose weak knowledge of orbitals: few candidates identify sp hybridisation of nitrogen, the modes of overlap are described imprecisely, and diagrams of a π bond often do not show sideways overlap of p orbitals. Formulae for magnesium nitride are frequently wrong or left unbalanced.

Where marks are lost on ammonia and the ammonium ion. Definitions of a weak acid or base confuse "partially dissolving" with partially dissociating. The equation for the ammonium ion acting as a weak acid in water is poorly done: the equilibrium sign is rare, charges are omitted, H₃O⁺ is given the state (l), and hydroxide ions are written in place of water. Some diagrams show nitrogen with five bonds, as though NH₅ existed, or fail to distinguish the dative bond and the charge.

Where marks are lost on the oxides of nitrogen. Questions on how oxides of nitrogen are formed in engines are frequently answered as if they were about catalytic converters, and "catalytic converter" is given as a man-made source. The equation for formation is often written with NO rather than N₂ as the reactant, and the fact that the nitrogen comes from the air is often missing. Equations for the catalytic converter frequently show only the nitrogen change and leave out the simultaneous oxidation of CO to CO₂, use the wrong oxide of nitrogen, or are not balanced. For photochemical smog, descriptions of how PAN forms lack precision, "VOCs" or "CO" is given when the product is asked for, and the name of a gas is given when the type of pollution is asked for; some answers claim hydrocarbons damage the ozone layer.

The weakest area: the catalytic role of NO₂. Only a small number of candidates give all the relevant equations and identify the catalytic role. Many give only the formation of sulfuric acid, or write about nitric acid instead. The regeneration equation causes the most difficulty and is often wrongly balanced; some answers write the first step backwards. Statements that NO₂ is "not used up" are common but unsupported by an equation, and many candidates do not state that a catalyst increases the rate. On the direct role, very few answers mention both HNO₃ and HNO₂. Examiners have also noted that dissolving SO₂ in water to form H₂SO₃ involves no change in oxidation number, and that SO₃ + H₂ and H₂SO₄ are common wrong products for that step.

What a successful response demonstrates. It answers the question actually asked (formation versus removal; product versus primary pollutant; type of pollution versus a gas); it supports every explanation with the specific chemistry — a strong bond and a high activation energy, a lone pair accepting a proton, a catalyst regenerated by a stated equation; and it writes complete, balanced equations with the right species, charges, state symbols and ⇌ where the reaction is an equilibrium.

Things that lose marks in this topic12

Each row is a misconception reported by examiners or one that leads directly to a reported error. Read across: the incorrect idea, why it fails, the correct model, and what it costs.

Self-test12

Twelve questions written for this page, covering the whole unit. Each gives the reasoning behind the answer once you have chosen.

Definitions to learn12

Summary and examination checklist12

Essential knowledge

• N₂ is unreactive because its N≡N bond is very strong (944 kJ mol⁻¹), giving reactions a very high activation energy, and because the molecule is non-polar, offering no site for electrophiles or nucleophiles.
• Ammonia is a weak Brønsted–Lowry base: the lone pair on nitrogen accepts a proton. In water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, far to the left.
• NH₄⁺ forms when that lone pair makes a dative covalent bond to H⁺. It is tetrahedral, 109.5°; all four bonds are identical once formed.
• NH₄⁺ is a weak acid; ammonium chloride solution is slightly acidic.
• Warming an ammonium salt with a base releases NH₃: NH₄⁺ + OH⁻ → NH₃ + H₂O. Damp red litmus turns blue.
• NO forms from N₂ and O₂ at high temperature: naturally in lightning; man-made in internal combustion engines (and other high-temperature combustion). NO is oxidised to NO₂ in air.
• Catalytic converters (Pt/Pd/Rh on a honeycomb) remove NO_x by cross-redox with CO (and hydrocarbons): 2CO + 2NO → 2CO₂ + N₂.
• NO_x + unburnt hydrocarbons in sunlight → PAN, a component of photochemical smog.
• Directly: NO₂ forms HNO₃ (and HNO₂) in rain. Catalytically: SO₂ + NO₂ → SO₃ + NO; NO + ½O₂ → NO₂; then SO₃ + H₂O → H₂SO₄.

Examination checklist

• Did I say the triple bond is strong, and add non-polarity for the second mark?
• Does every Brønsted–Lowry answer mention a proton being donated or accepted, and the lone pair on N?
• Do my equilibrium equations for NH₃ and NH₄⁺ in water have ⇌, charges and (aq)/(l)?
• Is my NH₄⁺ diagram showing a distinct dative pair, brackets and the + charge?
• Did I answer formation questions about the engine and removal questions about the converter?
• Does my converter equation contain both a nitrogen oxide and CO (or a hydrocarbon), and balance?
• For smog, did I name PAN as the product and "photochemical smog" as the type of pollution?
• For the catalytic role, did I give both equations, the regeneration, and the words "increases the rate"?
• Did I use the data given in the question rather than remembered values?

Knowledge organiser12

Data used on this page12

These are the values the models on this page compute with. They are reference values chosen to be consistent with one another; they are not the official data booklet, and values in other sources differ by a few units. In an examination, always use the data printed in the question or the data booklet supplied.

Bond energies

Thermochemical and equilibrium data at 298 K