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Year 8 UK
KS3 Chemistry
Energy Changes
⚡ Interactive
Bond Energy Calculations
Breaking bonds takes energy in — making bonds gives energy out
⚡ Interactive Energy Profile Diagram
Click the buttons below to see the energy profile for exothermic and endothermic reactions. Notice how bond breaking (energy IN) compares to bond making (energy OUT).
🔥 Energy IN = bond breaking
✅ Energy OUT = bond making
📐 ΔH = Energy OUT − Energy IN
📋 Table of Contents
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⚗️ What Are Chemical Bonds?
Section 1 · Chemical Bonds · KS3 Chemistry
A chemical bond is a force that holds atoms together in a molecule or compound. When atoms bond together, they share or transfer electrons in a way that makes them more stable. Without chemical bonds, molecules like water (H₂O), carbon dioxide (CO₂), and hydrogen (H₂) could not exist.
In Year 8 chemistry, we focus on covalent bonds — bonds where atoms share pairs of electrons. Covalent bonds are found in gases like hydrogen, oxygen, and methane, as well as in water. The strength of a covalent bond is described by its bond energy, which tells us how much energy is needed to break that bond.
🌟 Key Definition: A chemical bond is a force of attraction that holds atoms together. The bond energy is the amount of energy (in kJ/mol) needed to break one mole of that bond in the gas phase.
💡 Did you know? Different types of bonds have very different strengths. A C≡C triple bond (837 kJ/mol) is nearly three times stronger than a C–C single bond (347 kJ/mol) — which is why triple bonds are much harder to break!
🔗Bonds hold atoms together in molecules
⚡Bond energy tells us bond strength in kJ/mol
🔬Covalent bonds involve shared electron pairs
🌊Examples: H–H, O=O, O–H in water
Covalent bond Examples with Bonnd Energies
A visual showing common covalent bond examples — H–H (hydrogen), O=O (oxygen double bond), H–O–H (water) — with electron dot diagrams (Lewis structures) showing shared electron pairs.⚡ Quick Check — Section 1
1
What is a chemical bond?
A force that holds atoms together in a molecule or compound. In covalent bonds, atoms share pairs of electrons.
2
What does bond energy tell us?
Bond energy (in kJ/mol) tells us the amount of energy needed to break one mole of that bond — it is a measure of bond strength.
3
Which is stronger — a C–C single bond or a C=C double bond? How can you tell?
The C=C double bond is stronger. It has a higher bond energy (614 kJ/mol) compared to C–C (347 kJ/mol), meaning more energy is needed to break it.
⚡ Bond Breaking & Bond Making
Section 2 · Energy In / Energy Out · KS3 Chemistry
The key to understanding bond energy calculations is this fundamental rule: breaking bonds requires energy (energy must be put IN), and making bonds releases energy (energy comes OUT). This is the most important concept in this topic — if you remember nothing else, remember this!
When a chemical reaction happens, the bonds in the reactant molecules must first be broken. This takes in energy from the surroundings. New bonds then form in the product molecules, which releases energy back to the surroundings. The overall energy change of the reaction depends on how much energy was needed to break bonds versus how much was released when new bonds formed.
Bond Breaking
Energy IN ↑
Reactant bonds are broken
Energy absorbed from surroundings
Energy absorbed from surroundings
⇌
Bond Making
Energy OUT ↓
Product bonds are formed
Energy released to surroundings
Energy released to surroundings
🔑 The Golden Rule: BREAKING bonds = energy IN (endothermic step). MAKING bonds = energy OUT (exothermic step). The OVERALL reaction type depends on which is bigger.
⚠️ Common Mistake: Students often get this the wrong way round. Remember — breaking is like breaking up a friendship, it takes effort (energy in). Making is like making a new friendship, it feels good and releases energy!
Overall Energy Change (ΔH) = Energy IN (breaking) − Energy OUT (making)
Bond Breaking vs Bond Making — Energy Flow
Left: Energy is absorbed (IN) to break bonds. Right: Energy is released (OUT) when bonds form. Overall change: ΔH = energy in − energy out.⚡ Quick Check — Section 2
1
Does breaking a bond require energy to be taken IN or released OUT?
Breaking bonds requires energy to be taken IN — it is an endothermic process.
2
Does making a bond require energy to be taken in, or does it release energy?
Making bonds releases energy OUT to the surroundings — it is an exothermic process.
3
What is the formula for calculating the overall energy change of a reaction?
ΔH = Energy IN (breaking bonds) − Energy OUT (making bonds)
🔢 Bond Energy Values
Section 3 · Data Tables · KS3 Chemistry
Each type of covalent bond has a specific bond energy value, measured in kilojoules per mole (kJ/mol). This value is the average amount of energy required to break one mole of that bond, measured in the gas phase. These values are experimentally determined and collected in data tables that chemists use in calculations.
It is important to note that the same type of bond (e.g. C–H) will always have the same bond energy value regardless of which molecule it appears in — this is why we call them average bond energies. The more energy required to break a bond, the stronger that bond is.
🌟 Key Definition: Bond energy (also called bond enthalpy) is measured in kJ/mol and represents the energy needed to break one mole of a specific bond in gaseous molecules. Higher bond energy = stronger bond.
| Bond | Bond Energy (kJ/mol) | Found in | Strength |
|---|---|---|---|
| H–H | 436 | Hydrogen gas (H₂) | Medium |
| O=O | 498 | Oxygen gas (O₂) | Strong |
| O–H | 463 | Water (H₂O) | Strong |
| C–H | 413 | Methane, hydrocarbons | Medium-strong |
| C–C | 347 | Alkanes, organic molecules | Medium |
| C=C | 614 | Alkenes (e.g. ethene) | Very strong |
| C–O | 360 | Alcohols, esters | Medium |
| C=O | 743 | Carbon dioxide, aldehydes | Very strong |
| N≡N | 945 | Nitrogen gas (N₂) | Extremely strong |
| Cl–Cl | 243 | Chlorine gas (Cl₂) | Weak |
| H–Cl | 432 | Hydrogen chloride (HCl) | Medium |
| H–O | 463 | Water (same as O–H) | Strong |
🎯 Tip: In an exam, you will almost always be given a bond energy table — you don't need to memorise exact values! What matters is knowing HOW to use the table in calculations. Focus on the method, not the numbers.
Comparing Bond Energy Values — Bar Chart
This chart compares common bond energies with horizontal bars. Colours range from light blue (weaker bonds) to deep red (stronger bonds), with each bond labelled along with its energy value in kJ/mol.⚡ Quick Check — Section 3
1
What are the units of bond energy?
Bond energy is measured in kJ/mol (kilojoules per mole).
2
Looking at the table, which bond is the strongest? How do you know?
The N≡N bond (945 kJ/mol) is the strongest — it has the highest bond energy, meaning it requires the most energy to break.
3
Why do chemists use 'average' bond energy values rather than exact values?
The exact energy of a bond can vary slightly depending on the molecule it is in. Average bond energies are used as representative values that apply across different molecules.
🧮 How to Calculate Energy Changes
Section 4 · The Method · KS3 Chemistry
Bond energy calculations follow a clear, repeatable method every time. Once you understand the steps, you can apply them to any reaction. The method has four key steps: identify all bonds broken, identify all bonds made, look up each bond energy, then subtract to find the overall energy change.
The key thing to remember is to count bonds carefully. For example, methane (CH₄) has four C–H bonds — so you must multiply the C–H bond energy by 4. Missing a bond or multiplying incorrectly is the most common error in exams.
🎯 The 4-Step Method:
1️⃣ Write out the balanced equation and draw all bonds in reactants and products.
2️⃣ List ALL bonds broken (reactants) — multiply bond energy × number of that bond.
3️⃣ List ALL bonds made (products) — multiply bond energy × number of that bond.
4️⃣ Calculate: ΔH = Total Energy IN − Total Energy OUT.
1️⃣ Write out the balanced equation and draw all bonds in reactants and products.
2️⃣ List ALL bonds broken (reactants) — multiply bond energy × number of that bond.
3️⃣ List ALL bonds made (products) — multiply bond energy × number of that bond.
4️⃣ Calculate: ΔH = Total Energy IN − Total Energy OUT.
ΔH = Σ (Bond Energies of bonds broken) − Σ (Bond Energies of bonds made)
⚠️ Common Mistake — Counting Bonds: Always draw the structural formula and count each bond individually. For H₂O (H–O–H), there are TWO O–H bonds. For CO₂ (O=C=O), there are TWO C=O double bonds. Never guess — always count from the structure!
✅ Sign Conventions: If ΔH is negative (−), the reaction is exothermic — more energy was released making bonds than was used breaking them. If ΔH is positive (+), the reaction is endothermic — more energy was used breaking bonds than was released.
The 4-Step Bond Energy Calculation Method
Step 1: Draw structural formula → Step 2: List bonds broken (energy IN) → Step 3: List bonds made (energy OUT) → Step 4: Calculate ΔH = IN − OUT.⚡ Quick Check — Section 4
1
How many C–H bonds are in one molecule of methane (CH₄)?
There are 4 C–H bonds in methane (one carbon bonded to four hydrogen atoms).
2
If energy IN > energy OUT, is the reaction exothermic or endothermic?
Endothermic — more energy was used breaking bonds than was released making bonds, so the overall reaction absorbs energy from the surroundings.
3
What is the first step in any bond energy calculation?
Write out the balanced equation and draw out all the structural formulas of reactants and products, clearly showing all bonds.
🔥 Exothermic Reactions
Section 5 · Energy Released · KS3 Chemistry
A reaction is exothermic if the total energy released when new bonds form in the products is greater than the total energy needed to break the bonds in the reactants. In other words, energy OUT > energy IN. The extra energy is released to the surroundings — usually as heat, which is why the temperature of the surroundings rises.
The energy change (ΔH) for an exothermic reaction is always negative (e.g. −284 kJ/mol). This makes sense because energy has left the reaction system and gone into the surroundings. Common examples include combustion (burning), neutralisation, and respiration.
🌟 Key Fact: In an exothermic reaction: Energy OUT (bond making) > Energy IN (bond breaking). ΔH is negative. Temperature of surroundings rises. The products are at a lower energy level than the reactants.
🔥Combustion of fuels (burning methane, wood)
🧪Neutralisation (acid + alkali)
🫁Respiration in living cells
💥Explosions (rapid exothermic reactions)
Example: Combustion of hydrogen: H₂ + ½O₂ → H₂O
Bonds broken: 1 × H–H (436) + ½ × O=O (½ × 498 = 249)
Energy IN = 436 + 249 = 685 kJ/mol
Bonds made: 2 × O–H (2 × 463 = 926)
Energy OUT = 926 kJ/mol
ΔH = 685 − 926 = −241 kJ/mol ✅ EXOTHERMIC
Negative ΔH → exothermic → energy released to surroundings → temperature rises
⚠️ Common Mistake: Students sometimes think that because bond breaking is endothermic (takes energy in), the whole reaction must be endothermic. This is wrong! Most reactions are exothermic overall because the bond making step releases more energy than the breaking step uses.
Energy Level Diagram — Endothermic Reaction
Reactants are at lower energy than products. The activation energy hump is shown, and ΔH is positive indicating energy is absorbed.⚡ Quick Check — Section 5
1
In an exothermic reaction, which is larger: energy IN or energy OUT?
Energy OUT (bond making) is larger. More energy is released making product bonds than is needed to break reactant bonds.
2
A reaction has ΔH = −350 kJ/mol. Is it exothermic or endothermic? What happens to the temperature?
Exothermic — ΔH is negative. The temperature of the surroundings rises as heat is released.
3
Give two everyday examples of exothermic reactions.
Any two from: burning fuels (combustion), neutralisation reactions, respiration in cells, hand warmers, explosions.
❄️ Endothermic Reactions
Section 6 · Energy Absorbed · KS3 Chemistry
A reaction is endothermic if the total energy needed to break the bonds in the reactants is greater than the energy released when new bonds form in the products. In other words, energy IN > energy OUT. The reaction must absorb energy from the surroundings — which causes the temperature of the surroundings to fall.
The energy change (ΔH) for an endothermic reaction is always positive (e.g. +178 kJ/mol). This means energy has entered the reaction system from the surroundings. Examples include thermal decomposition, dissolving ammonium nitrate, and photosynthesis.
🌟 Key Fact: In an endothermic reaction: Energy IN (bond breaking) > Energy OUT (bond making). ΔH is positive. Temperature of surroundings falls. The products are at a higher energy level than the reactants.
🌿Photosynthesis (absorbs light energy)
🔬Thermal decomposition (heating to break down)
❄️Dissolving ammonium nitrate (cold packs)
🍳Cooking food (energy absorbed from heat)
Example: Decomposition of hydrogen iodide: 2HI → H₂ + I₂
Bonds broken: 2 × H–I (2 × 297 = 594)
Energy IN = 594 kJ/mol
Bonds made: 1 × H–H (436) + 1 × I–I (151)
Energy OUT = 436 + 151 = 587 kJ/mol
ΔH = 594 − 587 = +7 kJ/mol ✅ ENDOTHERMIC
Positive ΔH → endothermic → energy absorbed from surroundings → temperature falls
💡 Did you know? Cold packs used in sports medicine work because the chemical reaction inside (dissolving ammonium nitrate in water) is endothermic — it absorbs heat from your skin, making the pack feel ice-cold. No refrigeration needed!
Energy Level Diagram — Endothermic Reaction
This diagram shows an endothermic reaction where energy is absorbed. Reactants start at a lower energy level, rise over the activation energy hump, and form products at a higher energy level. The upward arrow represents ΔH (positive).⚡ Quick Check — Section 6
1
In an endothermic reaction, what happens to the temperature of the surroundings?
The temperature falls — the reaction absorbs heat energy from the surroundings, making the surroundings feel colder.
2
A reaction has ΔH = +210 kJ/mol. What type of reaction is this?
Endothermic — ΔH is positive, meaning the reaction absorbs energy from the surroundings.
3
Why is photosynthesis classed as endothermic?
Photosynthesis absorbs light energy from the sun to break bonds in carbon dioxide and water. More energy is needed to break reactant bonds than is released making product bonds — so it is endothermic.
📋 Worked Examples
Section 7 · Step-by-Step Calculations · KS3 Chemistry
The best way to master bond energy calculations is to practise worked examples. Here are two fully worked calculations — one exothermic and one endothermic — showing every step of the method. Study these carefully before attempting the quiz questions.
🔥 Worked Example 1 — Combustion of Methane (Exothermic)
Reaction: CH₄ + 2O₂ → CO₂ + 2H₂O
Reaction: CH₄ + 2O₂ → CO₂ + 2H₂O
Step 1 — Identify all bonds BROKEN (reactants): CH₄ + 2O₂
CH₄ has 4 × C–H bonds: 4 × 413 = 1652 kJ
2O₂ has 2 × O=O bonds: 2 × 498 = 996 kJ
Total Energy IN = 1652 + 996 = 2648 kJ/mol
Step 2 — Identify all bonds MADE (products): CO₂ + 2H₂O
CO₂ (O=C=O) has 2 × C=O bonds: 2 × 743 = 1486 kJ
2H₂O has 4 × O–H bonds: 4 × 463 = 1852 kJ
Total Energy OUT = 1486 + 1852 = 3338 kJ/mol
Step 3 — Calculate ΔH:
ΔH = Energy IN − Energy OUT = 2648 − 3338
ΔH = −690 kJ/mol → EXOTHERMIC ✅
Negative → energy released → surroundings get hotter
❄️ Worked Example 2 — Formation of Nitrogen Monoxide (Endothermic)
Reaction: N₂ + O₂ → 2NO
Reaction: N₂ + O₂ → 2NO
Step 1 — Identify all bonds BROKEN (reactants): N₂ + O₂
N₂ has 1 × N≡N bond: 1 × 945 = 945 kJ
O₂ has 1 × O=O bond: 1 × 498 = 498 kJ
Total Energy IN = 945 + 498 = 1443 kJ/mol
Step 2 — Identify all bonds MADE (products): 2NO
2NO has 2 × N=O bonds: 2 × 607 = 1214 kJ
Total Energy OUT = 1214 kJ/mol
Step 3 — Calculate ΔH:
ΔH = Energy IN − Energy OUT = 1443 − 1214
ΔH = +229 kJ/mol → ENDOTHERMIC ✅
Positive → energy absorbed → surroundings get cooler
⚠️ Common Mistake — Forgetting to Multiply: In Example 1, there are 2 molecules of H₂O in the products, each containing 2 O–H bonds. Students often forget to multiply by 2 (for 2 molecules) × 2 (for 2 bonds each) = 4 bonds total. Always check your multiplications carefully!
Worked Example — Combustion of Methane with Structural Formulas
Combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O) showing bonds broken (orange/red), bonds formed (blue/green), and bond energy calculation with final ΔH.⚡ Quick Check — Section 7
1
In the combustion of methane example, why are there 4 O–H bonds made (not 2)?
There are 2 molecules of H₂O produced, and each H₂O molecule contains 2 O–H bonds. So 2 molecules × 2 bonds = 4 O–H bonds total.
2
For the reaction H₂ + Cl₂ → 2HCl: how many bonds are broken and how many are made?
Bonds broken: 1 × H–H and 1 × Cl–Cl (2 bonds total broken). Bonds made: 2 × H–Cl (2 bonds made). The reaction forms HCl twice.
3
A student gets ΔH = +50 kJ/mol. What does this tell us about the reaction?
The reaction is endothermic — more energy was absorbed breaking bonds than was released making bonds. The surroundings would cool down.
🧠 Multiple Choice Quiz
10 questions — calculations, concepts, and applications
0 of 10 answered
Question 1 of 10
What is the energy change when breaking a chemical bond?
✅ Correct Answer: B) Energy is absorbed from the surroundings
Question 2 of 10
Which formula correctly calculates the overall energy change (ΔH) of a reaction?
✅ Correct Answer: C) ΔH = Energy IN − Energy OUT
Question 3 of 10
For the reaction H₂ + Cl₂ → 2HCl, using bond energies: H–H = 436, Cl–Cl = 243, H–Cl = 432 kJ/mol. What is ΔH?
✅ Correct Answer: A) −185 kJ/mol
Question 4 of 10
A reaction has ΔH = +320 kJ/mol. What type of reaction is this and what happens to the temperature of the surroundings?
✅ Correct Answer: D) Endothermic — temperature falls
Question 5 of 10
How many O–H bonds are broken when 2 molecules of water (H₂O) are split apart?
✅ Correct Answer: C) 4
Question 6 of 10
In the combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O), how many C–H bonds are broken?
✅ Correct Answer: D) 4
Question 7 of 10
Bond energies: N≡N = 945 kJ/mol, H–H = 436 kJ/mol, N–H = 391 kJ/mol. For N₂ + 3H₂ → 2NH₃, what is the total energy IN (bonds broken)?
✅ Correct Answer: A) 1381 kJ/mol
Question 8 of 10
Which of the following is an example of an endothermic reaction?
✅ Correct Answer: C) Photosynthesis in plants
Question 9 of 10
A student calculates: energy IN = 1200 kJ/mol and energy OUT = 1800 kJ/mol. What is ΔH and what type of reaction is this?
✅ Correct Answer: B) ΔH = −600 kJ/mol, exothermic
Question 10 of 10
Why is the N≡N bond in nitrogen gas (N₂) particularly difficult to break?
✅ Correct Answer: B) It is a triple bond with very high bond energy (945 kJ/mol)
✍️ Long Answer Questions
Detailed responses — model answers and mark schemes included
4 marks
Q1. Explain, using the terms 'bond breaking' and 'bond making', why the combustion of fuels releases energy to the surroundings.
💡 Hint: Think about which step takes energy in, which step releases energy, and which step is bigger in combustion.
✅ Model Answer
In any chemical reaction, energy is first required to break the bonds in the reactant molecules. This is an endothermic step — energy is taken in from the surroundings.
When the new bonds form in the product molecules, energy is released back to the surroundings. This is an exothermic step.
In combustion reactions, the energy released when making the new bonds in the products (such as C=O in CO₂ and O–H in H₂O) is greater than the energy needed to break the bonds in the reactants (such as C–H in the fuel and O=O in oxygen).
Because energy out > energy in, the overall reaction is exothermic. The excess energy is released to the surroundings as heat, causing the temperature of the surroundings to rise.
📋 Mark Scheme
1 mark: Bond breaking requires energy to be taken in (endothermic step).
1 mark: Bond making releases energy to the surroundings (exothermic step).
1 mark: In combustion, energy released making bonds (products) is greater than energy needed to break bonds (reactants).
1 mark: Overall reaction is exothermic — energy released as heat — temperature of surroundings rises.
5 marks
Q2. Use the bond energy data below to calculate the overall energy change (ΔH) for the reaction:
H₂ + Cl₂ → 2HCl
Bond energies: H–H = 436 kJ/mol, Cl–Cl = 243 kJ/mol, H–Cl = 432 kJ/mol
State whether the reaction is exothermic or endothermic and explain what this means for the temperature of the surroundings.
H₂ + Cl₂ → 2HCl
Bond energies: H–H = 436 kJ/mol, Cl–Cl = 243 kJ/mol, H–Cl = 432 kJ/mol
State whether the reaction is exothermic or endothermic and explain what this means for the temperature of the surroundings.
💡 Hint: Step 1: list bonds broken. Step 2: list bonds made. Step 3: calculate ΔH = IN − OUT. Step 4: state reaction type.
✅ Model Answer
Step 1 — Bonds broken (reactants):
1 × H–H bond: 1 × 436 = 436 kJ/mol
1 × Cl–Cl bond: 1 × 243 = 243 kJ/mol
Total energy IN = 436 + 243 = 679 kJ/mol
Step 2 — Bonds made (products):
2 × H–Cl bonds: 2 × 432 = 864 kJ/mol
Total energy OUT = 864 kJ/mol
Step 3 — Calculate ΔH:
ΔH = Energy IN − Energy OUT = 679 − 864 = −185 kJ/mol
The reaction is exothermic (ΔH is negative). This means that more energy is released making new bonds in HCl than is needed to break the bonds in H₂ and Cl₂. The excess energy is released to the surroundings as heat, so the temperature of the surroundings rises.
📋 Mark Scheme
1 mark: Correct total energy IN = 679 kJ/mol (436 + 243).
1 mark: Correct total energy OUT = 864 kJ/mol (2 × 432).
1 mark: Correct ΔH = 679 − 864 = −185 kJ/mol.
1 mark: States reaction is exothermic (negative ΔH).
1 mark: States temperature of surroundings rises / energy released as heat.
6 marks
Q3. Use the bond energy values below to calculate ΔH for the combustion of methane:
CH₄ + 2O₂ → CO₂ + 2H₂O
Bond energies: C–H = 413, O=O = 498, C=O = 743, O–H = 463 kJ/mol
Show all working and state whether the reaction is exothermic or endothermic.
CH₄ + 2O₂ → CO₂ + 2H₂O
Bond energies: C–H = 413, O=O = 498, C=O = 743, O–H = 463 kJ/mol
Show all working and state whether the reaction is exothermic or endothermic.
💡 Hint: Count bonds carefully: CH₄ has 4 C–H bonds, 2O₂ has 2 O=O bonds, CO₂ has 2 C=O bonds, and 2H₂O has 4 O–H bonds.
✅ Model Answer
Step 1 — Bonds broken (reactants: CH₄ + 2O₂):
4 × C–H bonds: 4 × 413 = 1652 kJ/mol
2 × O=O bonds: 2 × 498 = 996 kJ/mol
Total energy IN = 1652 + 996 = 2648 kJ/mol
Step 2 — Bonds made (products: CO₂ + 2H₂O):
2 × C=O bonds: 2 × 743 = 1486 kJ/mol
4 × O–H bonds: 4 × 463 = 1852 kJ/mol
Total energy OUT = 1486 + 1852 = 3338 kJ/mol
Step 3 — Calculate ΔH:
ΔH = 2648 − 3338 = −690 kJ/mol
The reaction is exothermic. ΔH is negative, confirming that more energy is released making the new bonds in CO₂ and H₂O than is needed to break the bonds in CH₄ and O₂. The combustion of methane releases 690 kJ of energy per mole of methane burned.
📋 Mark Scheme
1 mark: Energy from C–H bonds = 4 × 413 = 1652 kJ/mol.
1 mark: Energy from O=O bonds = 2 × 498 = 996 kJ/mol.
1 mark: Energy from C=O bonds = 2 × 743 = 1486 kJ/mol.
1 mark: Energy from O–H bonds = 4 × 463 = 1852 kJ/mol.
1 mark: Correct ΔH = 2648 − 3338 = −690 kJ/mol.
1 mark: States exothermic and gives explanation (energy OUT > energy IN).
4 marks
Q4. The reaction N₂ + 3H₂ → 2NH₃ is used to make ammonia in the Haber Process. Bond energies: N≡N = 945, H–H = 436, N–H = 391 kJ/mol.
Calculate ΔH and determine whether the reaction is exothermic or endothermic. Explain what this means for industrial use.
Calculate ΔH and determine whether the reaction is exothermic or endothermic. Explain what this means for industrial use.
💡 Hint: 2NH₃ contains 2 molecules each with 3 N–H bonds = 6 N–H bonds total in products.
✅ Model Answer
Bonds broken (reactants):
1 × N≡N: 1 × 945 = 945 kJ/mol
3 × H–H: 3 × 436 = 1308 kJ/mol
Total energy IN = 945 + 1308 = 2253 kJ/mol
Bonds made (products: 2NH₃):
6 × N–H bonds: 6 × 391 = 2346 kJ/mol
Total energy OUT = 2346 kJ/mol
ΔH = 2253 − 2346 = −93 kJ/mol
The reaction is exothermic (ΔH negative). For industrial use in the Haber Process, this means the reaction releases heat energy. In practice, the process is carried out at high temperatures (around 450°C) to increase the rate of reaction — despite this making the equilibrium less favourable. The exothermic nature means the reactor must be cooled to maintain the optimal temperature.
📋 Mark Scheme
1 mark: Energy IN (N≡N + 3H–H) = 945 + 1308 = 2253 kJ/mol.
1 mark: Energy OUT (6 × N–H) = 2346 kJ/mol.
1 mark: ΔH = 2253 − 2346 = −93 kJ/mol.
1 mark: States exothermic; explains industrial implication (releases heat / reactor must be cooled / carried out at high temperature for rate).
4 marks
Q5. Compare exothermic and endothermic reactions in terms of: (a) bond making and breaking, (b) the sign of ΔH, (c) the effect on the temperature of the surroundings, and (d) give one example of each.
💡 Hint: Draw a comparison table in your head — what is different about each type for each category?
✅ Model Answer
(a) Bond making and breaking:
In exothermic reactions, the energy released making bonds in the products is greater than the energy needed to break bonds in the reactants (energy out > energy in). In endothermic reactions, it is the reverse — more energy is needed to break bonds than is released making new bonds (energy in > energy out).
(b) Sign of ΔH:
Exothermic reactions have a negative ΔH (e.g. −300 kJ/mol), indicating energy has been released. Endothermic reactions have a positive ΔH (e.g. +200 kJ/mol), indicating energy has been absorbed.
(c) Effect on surroundings:
In exothermic reactions, heat is released to the surroundings, so the temperature of the surroundings rises. In endothermic reactions, heat is absorbed from the surroundings, so the temperature of the surroundings falls.
(d) Examples:
Exothermic: combustion of fuels / neutralisation of an acid with alkali.
Endothermic: photosynthesis / thermal decomposition / dissolving ammonium nitrate.
📋 Mark Scheme
1 mark: Exothermic — energy out > energy in; endothermic — energy in > energy out.
1 mark: Exothermic ΔH negative; endothermic ΔH positive.
1 mark: Exothermic — temperature rises; endothermic — temperature falls.
1 mark: One correct example of each (exothermic and endothermic).
⚡ One Word Quiz
20 quick-fire questions — type your answer and press Check!
Q1. What are the units of bond energy?
Q2. Does breaking a chemical bond require energy to go IN or come OUT?
Q3. Does making a chemical bond release energy OUT or require energy IN?
Q4. What is the sign of ΔH for an exothermic reaction?
Q5. What is the sign of ΔH for an endothermic reaction?
Q6. In an exothermic reaction, what happens to the temperature of the surroundings?
Q7. In an endothermic reaction, what happens to the temperature of the surroundings?
Q8. What is the bond energy of the H–H bond (kJ/mol)?
Q9. What is the bond energy of the O=O bond (kJ/mol)?
Q10. How many C–H bonds are in one molecule of methane (CH₄)?
Q11. Give one example of an exothermic reaction.
Q12. Give one example of an endothermic reaction.
Q13. What is the formula for calculating the overall energy change ΔH?
Q14. Which bond is stronger: C–C or C=C?
Q15. What symbol is used for the overall energy change of a reaction?
Q16. How many O–H bonds are in two molecules of water (2H₂O)?
Q17. Is photosynthesis exothermic or endothermic?
Q18. What type of bond do atoms share in a covalent bond?
Q19. In the Haber Process reaction N₂ + 3H₂ → 2NH₃, how many N–H bonds are formed?
Q20. What process in living cells is an example of an exothermic reaction?