Hey Tara! Welcome to Topic 5 - Chemical Energetics. Have you ever wondered why a matchstick feels hot when you strike it, or why an instant cold pack gets icy without being in a freezer? The answer lies in the energy changes during chemical reactions. In this topic, you will learn to classify reactions as exothermic or endothermic, draw and interpret energy level diagrams, understand activation energy, and use bond energies to calculate exactly how much energy a reaction releases or absorbs. This is one of the most practical topics in chemistry - it connects directly to cooking, engines, rockets, and even the biology happening inside your own cells right now. Let us dive in!
5.1 Exothermic and Endothermic Reactions
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What Are Energy Changes in Reactions?
Every chemical reaction involves a transfer of energy between the reacting chemicals and their surroundings. The surroundings means everything around the reaction - the solution, the container, the air, and even the thermometer you might use to measure the temperature.
There are two fundamental types of energy change:
- Energy can be released from the reaction to the surroundings (making the surroundings hotter)
- Energy can be absorbed from the surroundings into the reaction (making the surroundings cooler)
These two types give us the two key definitions you need to know.
Exothermic Reactions
An exothermic reaction is a reaction that transfers energy to the surroundings, usually as heat. The temperature of the surroundings increases.
Think of it this way: "exo" means "out" (like "exit"). Energy goes out of the reaction and into the surroundings.
Memory Trick
EXO = EXIT - energy EXITS the reaction. The surroundings warm up. If you hold the beaker and it feels hot, the reaction is exothermic.
Common Examples of Exothermic Reactions
| Reaction | Why it is Exothermic | Everyday Example |
|---|---|---|
| Combustion (burning) | Fuels react with oxygen, releasing large amounts of heat and light energy | Burning natural gas in a kitchen stove, petrol in a car engine, wood in a campfire |
| Neutralisation | Acid + alkali reactions release heat energy | Antacid tablets reacting with stomach acid (the stomach area may feel slightly warmer) |
| Respiration | Glucose reacts with oxygen inside living cells, releasing energy for life processes | Your body converts food into energy - this is why you feel warm even on a cold day |
| Oxidation of metals | Metals react with oxygen, releasing heat | Iron rusting (very slow), hand warmers (iron powder oxidation) |
| Dissolving concentrated acids in water | The hydration of acid molecules releases heat energy | Adding concentrated sulfuric acid to water causes significant heating (always add acid to water, never the reverse!) |
Combustion - The Most Important Exothermic Reaction
Combustion is the reaction of a substance with oxygen. It is always exothermic. Complete combustion of hydrocarbons produces carbon dioxide and water:
CH₄ + 2O₂ → CO₂ + 2H₂O
Methane burns in oxygen to produce carbon dioxide and water, releasing heat and light.
This is the reaction that heats homes (natural gas is mostly methane), powers gas stoves, and runs many power stations around the world.
Neutralisation - Exothermic Too!
When an acid reacts with an alkali, heat is always released:
HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
Hydrochloric acid + sodium hydroxide → sodium chloride + water (+ heat energy released)
If you measure the temperature of the solution during neutralisation, you will see it rise. A typical neutralisation might increase the temperature by 5-10 °C.
Respiration - Exothermic Inside You
Respiration is the process by which living cells break down glucose to release energy. It happens in every cell of your body, 24 hours a day:
C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O
Glucose + oxygen → carbon dioxide + water (+ energy released for life processes)
This is why your body maintains a temperature of about 37 °C - the exothermic reactions of respiration generate continuous heat.
Endothermic Reactions
An endothermic reaction is a reaction that takes in energy from the surroundings, usually as heat. The temperature of the surroundings decreases.
"Endo" means "in" (like "enter"). Energy goes into the reaction from the surroundings.
Memory Trick
ENDO = ENTER - energy ENTERS the reaction. The surroundings cool down. If you hold the beaker and it feels cold, the reaction is endothermic.
Common Examples of Endothermic Reactions
| Reaction | Why it is Endothermic | Everyday Example |
|---|---|---|
| Photosynthesis | Plants absorb light energy from the Sun to convert CO₂ and water into glucose | Every green plant and tree you see - they are absorbing energy from sunlight |
| Thermal decomposition | Compounds are broken down by continuous heating - they need energy input | Heating limestone (CaCO₃) in a lime kiln to make quicklime (CaO) |
| Dissolving ammonium nitrate in water | The dissolving process absorbs heat from the water | Instant cold packs used for sports injuries - squeeze the pack, the salt dissolves, it gets cold |
| Citric acid + sodium hydrogen carbonate | This reaction absorbs heat energy | Sherbet powder fizzing on your tongue feels cold because it absorbs heat from your mouth |
| Electrolysis | Electrical energy is continuously needed to decompose compounds | Electrolysis of aluminium oxide in industry requires enormous electrical energy input |
Photosynthesis - The Key Endothermic Reaction
6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂
Carbon dioxide + water → glucose + oxygen (light energy absorbed from the Sun)
Photosynthesis is the reverse of respiration. It requires a continuous input of light energy from the Sun. Without this energy input, the reaction would not happen.
Thermal Decomposition - Needs Heat to Break Down
CaCO₃ → CaO + CO₂
Calcium carbonate → calcium oxide + carbon dioxide (heat energy must be continuously supplied)
If you stop heating, the decomposition stops. The reaction needs a continuous input of heat energy - this is the hallmark of an endothermic process.
Summary: Exothermic vs Endothermic
| Feature | Exothermic | Endothermic |
|---|---|---|
| Energy transfer | Energy released to surroundings | Energy absorbed from surroundings |
| Temperature change | Surroundings get hotter (temperature rises) | Surroundings get cooler (temperature falls) |
| Sign of ΔH | Negative (ΔH < 0) | Positive (ΔH > 0) |
| Products energy level | Lower than reactants | Higher than reactants |
| Examples | Combustion, neutralisation, respiration | Photosynthesis, thermal decomposition, dissolving NH₄NO₃ |
Energy Level Diagrams (Reaction Profile Diagrams)
An energy level diagram (also called a reaction profile diagram) is a graph that shows the energy of the reactants and products during a chemical reaction. The y-axis shows energy, and the x-axis shows the progress of the reaction.
Energy Level Diagram for an Exothermic Reaction
Energy level diagram for an exothermic reaction. Products are LOWER than reactants. ΔH is negative.
Key features of the exothermic energy level diagram:
- The reactants are drawn at a higher energy level
- The products are drawn at a lower energy level
- The difference in height between reactants and products represents the enthalpy change (ΔH), which is negative
- Energy has been released - that is why the products are lower (they have less energy, and the "missing" energy went to the surroundings)
Energy Level Diagram for an Endothermic Reaction
Energy level diagram for an endothermic reaction. Products are HIGHER than reactants. ΔH is positive.
Key features of the endothermic energy level diagram:
- The reactants are drawn at a lower energy level
- The products are drawn at a higher energy level
- The difference in height between reactants and products represents the enthalpy change (ΔH), which is positive
- Energy has been absorbed - the products have more energy because they took energy from the surroundings
Supplement
Activation Energy
Activation energy (Eₐ) is the minimum energy that colliding particles must have in order for a reaction to occur.
Think of it like a hill. Even in an exothermic reaction (which releases energy overall), the reactants first need a "push" to get over the energy barrier before they can roll down to the lower energy products. This initial energy push is the activation energy.
Key points about activation energy:
- ALL reactions have an activation energy - even exothermic ones
- The activation energy is shown as the "hump" or peak on the energy level diagram
- It is measured from the energy level of the reactants up to the top of the energy barrier
- A match needs to be struck (providing the activation energy) before it can burn exothermically
- A spark plug provides the activation energy for petrol to combust in a car engine
Memory Trick
Activation energy is like the effort to push a boulder to the top of a hill. Once it is at the top, it rolls down by itself (releasing energy). But you need that initial push first!
How Catalysts Affect Activation Energy
A catalyst is a substance that increases the rate of a reaction without being chemically changed itself at the end of the reaction. It works by providing an alternative reaction pathway with a lower activation energy.
Effect of a catalyst: the activation energy is lowered, but the overall energy change (ΔH) remains the same.
Important points about catalysts and activation energy:
- A catalyst lowers the activation energy by providing an alternative pathway
- A catalyst does NOT change ΔH - the overall energy change of the reaction stays the same
- A catalyst does NOT change the amount of product formed - it only makes the reaction happen faster
- The catalyst is not used up in the reaction (it can be recovered unchanged at the end)
Understanding ΔH (Enthalpy Change)
The symbol ΔH (delta H) represents the enthalpy change of a reaction - the overall energy change measured in kilojoules per mole (kJ/mol).
| Type of Reaction | Sign of ΔH | What It Means |
|---|---|---|
| Exothermic | Negative (e.g. ΔH = −890 kJ/mol) | Energy has been released. Products have less energy than reactants. |
| Endothermic | Positive (e.g. ΔH = +178 kJ/mol) | Energy has been absorbed. Products have more energy than reactants. |
Exam Tip
The sign of ΔH tells you everything! If ΔH is negative, the reaction is exothermic. If ΔH is positive, the reaction is endothermic. Never forget the sign.
When drawing energy level diagrams, always label: (1) the reactants, (2) the products, (3) the enthalpy change ΔH with its sign, and (4) the activation energy Eₐ if asked for the extended tier.
Worked Examples
Worked Example 1
A student mixes hydrochloric acid with sodium hydroxide solution in a polystyrene cup. The temperature rises from 22 °C to 30 °C. Is this reaction exothermic or endothermic? Explain your answer.
Step 1: Observe
The temperature of the surroundings (the solution) increased from 22 °C to 30 °C. That is a rise of 8 °C.
Step 2: Apply the Definition
Since the temperature of the surroundings increased, energy has been transferred FROM the reaction TO the surroundings.
Step 3: Classify
A reaction that releases energy to the surroundings and causes the temperature to rise is an exothermic reaction.
This is an exothermic reaction because the temperature of the surroundings increased, showing that energy was transferred from the reaction to the surroundings. This is a neutralisation reaction (acid + alkali), which is always exothermic.
Worked Example 2
Ammonium nitrate is dissolved in water. The temperature drops from 25 °C to 17 °C. Draw an energy level diagram for this reaction and label the enthalpy change.
Step 1: Identify the Type
Temperature dropped by 8 °C. The surroundings got cooler. Therefore, energy was absorbed FROM the surroundings INTO the reaction. This is endothermic.
Step 2: Draw the Diagram
For an endothermic reaction: reactants go on the LEFT at a LOWER level. Products go on the RIGHT at a HIGHER level. The arrow for ΔH points UPWARDS and is labelled as positive.
Step 3: Label
- Label the y-axis "Energy"
- Label the x-axis "Progress of reaction"
- Label "Reactants: NH₄NO₃(s) + water" on the lower line
- Label "Products: NH₄⁺(aq) + NO₃⁻(aq)" on the upper line
- Draw a double-headed arrow between the two levels and label it "ΔH = +ve"
The energy level diagram shows reactants at a lower energy level and products at a higher energy level. ΔH is positive because the reaction is endothermic - energy has been absorbed from the surroundings.
Worked Example 3
The combustion of methane has ΔH = −890 kJ/mol. Explain what this value tells you and draw the energy level diagram with activation energy labelled.
Step 1: Interpret the Sign
ΔH = −890 kJ/mol. The negative sign tells us this is an exothermic reaction - energy is released to the surroundings.
Step 2: Interpret the Value
890 kJ of energy is released for every mole of methane (CH₄) that burns completely. This is a lot of energy, which is why methane is such a useful fuel.
Step 3: Draw the Diagram
- Reactants (CH₄ + 2O₂) drawn at a HIGH energy level
- Products (CO₂ + 2H₂O) drawn at a LOW energy level
- The gap between them labelled "ΔH = −890 kJ/mol"
- A hump above the reactant level representing the activation energy (Eₐ) - this is the spark or flame needed to start the combustion
- Eₐ arrow from reactant level to the peak of the hump
ΔH = −890 kJ/mol means 890 kJ of heat energy is released to the surroundings per mole of methane burned. The reaction is highly exothermic. On the diagram, reactants are high, products are low, with a small activation energy hump (provided by a spark or match).
Exam Tips for 5.1
1. Temperature change tells you the type: Temperature goes UP = exothermic. Temperature goes DOWN = endothermic. This is the simplest way to identify the reaction type in an experiment.
2. Do NOT say "the reaction gets hot": Say "the temperature of the surroundings increases" or "energy is transferred to the surroundings." The examiner wants precise language.
3. Activation energy is always positive: Even exothermic reactions need activation energy. Do not confuse Eₐ with ΔH.
4. Catalysts do NOT change ΔH: A common mistake is saying catalysts make the reaction more exothermic. They do not - they only lower the activation energy.
5. Learn the classic examples: Combustion (exo), neutralisation (exo), respiration (exo), photosynthesis (endo), thermal decomposition (endo), dissolving NH₄NO₃ (endo). These appear in nearly every exam.
6. On diagrams, always use arrows: Show ΔH with a labelled double-headed arrow between reactant and product levels. Show Eₐ with an arrow from the reactant level to the peak of the hump.
Apply It: Real-World Chemistry
Chemistry is not just theory - it is everywhere around you. Here are real scenarios where exothermic and endothermic reactions matter.
1
A food chemist at Britannia Industries in Bangalore is developing a new self-heating lunch box for railway travellers. The box contains a separate compartment with calcium oxide (quicklime) and a water pouch. When the traveller pulls a tab, water mixes with the quicklime and the food above heats up within 10 minutes.
Why does the food heat up? What type of reaction is occurring, and why is quicklime chosen?
The Chemistry
CaO + H₂O → Ca(OH)₂. This is a highly exothermic reaction. When calcium oxide (quicklime) reacts with water, it releases a large amount of heat energy - enough to raise the temperature of the food compartment above 70 °C.
Why Quicklime?
Quicklime is chosen because: (1) it reacts vigorously with water, releasing heat rapidly; (2) it is cheap and widely available; (3) the product - calcium hydroxide - is a relatively safe alkali; (4) it is a dry powder, easy to store and package.
Chemistry Connection
This is a real-world exothermic reaction being used to solve a practical problem. The energy released by the CaO + H₂O reaction transfers as heat to the food, raising its temperature. No electricity or fire needed!
2
A sports physiotherapist at a Premier League football club in Birmingham keeps instant cold packs in the medical kit. When a player sprains an ankle during a match, she squeezes the pack and it becomes ice-cold within seconds - without any freezer or ice.
Explain the chemistry behind the instant cold pack. Why does it get cold?
The Chemistry
Inside the cold pack, there is a pouch of water and solid ammonium nitrate (NH₄NO₃) crystals, separated by a thin barrier. When the pack is squeezed, the barrier breaks and the ammonium nitrate dissolves in the water.
Why It Gets Cold
Dissolving ammonium nitrate in water is an endothermic process. The dissolving process absorbs heat energy from the water (and the surrounding area), causing the temperature to drop dramatically - often to around 2-5 °C.
Chemistry Connection
The cold pack demonstrates an endothermic process perfectly. Energy is absorbed from the surroundings (the water and the injured area), so the surroundings cool down. No ice or refrigeration needed - just chemistry!
3
A catalytic converter manufacturer in Stuttgart, Germany, designs platinum-rhodium catalysts for car exhaust systems. Their latest model reduces harmful CO and NOₓ emissions by 95% while operating at lower temperatures than older designs.
How does the catalyst in a catalytic converter help reduce emissions? Does it change the energy released by the combustion reactions?
The Chemistry
The platinum-rhodium catalyst provides an alternative pathway with lower activation energy for the conversion of harmful gases. For example: 2CO + 2NO → 2CO₂ + N₂. Without the catalyst, these reactions would need much higher temperatures to occur at a useful rate.
Does It Change ΔH?
No! The catalyst lowers the activation energy, making the reaction happen faster at lower temperatures. But the overall enthalpy change (ΔH) remains exactly the same. The same amount of energy is released - it just happens more easily.
Chemistry Connection
This perfectly illustrates the role of catalysts: they lower Eₐ without changing ΔH. The reaction still releases the same energy, but it occurs at a lower temperature and faster rate, making it practical for cleaning car exhaust.
4
A team of agricultural scientists at the Indian Agricultural Research Institute (IARI) in Delhi is studying why rice paddy yields in Punjab are highest during the summer monsoon season when sunlight is abundant and temperatures are warm.
How does the endothermic nature of photosynthesis explain why rice yields depend on sunlight availability?
The Chemistry
Photosynthesis is endothermic: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂. The reaction requires a continuous input of light energy from the Sun. Without sufficient light energy, the rate of photosynthesis drops, and less glucose is produced for the plant to grow.
The Link to Yield
During the monsoon season in Punjab, there are long daylight hours interspersed with rain, providing both the light energy (for the endothermic photosynthesis) and water (a reactant). More photosynthesis means more glucose, more growth, and higher rice yields.
Chemistry Connection
Because photosynthesis is endothermic, it absolutely depends on an external energy source (sunlight). No energy input = no reaction = no plant growth. This is why cloudy, dark winters produce less plant growth than bright, sunny seasons.
5
NASA engineers at the Johnson Space Center in Houston are testing a new hand warmer design for astronauts conducting spacewalks. The gloves contain iron powder, salt, activated carbon, and vermiculite in small packets. When exposed to oxygen, the packets generate heat for up to 8 hours at −250 °C outside the space station.
What exothermic reaction powers these hand warmers? Why does the reaction need oxygen, and how does the salt act as a catalyst?
The Chemistry
The exothermic reaction is the oxidation of iron: 4Fe + 3O₂ → 2Fe₂O₃. This is essentially a controlled, slow rusting reaction. It releases significant heat energy.
Role of Oxygen and Salt
Oxygen is a reactant - without it, the iron cannot oxidise. The salt (NaCl) and water vapour act as a catalyst, speeding up the oxidation by lowering the activation energy. The activated carbon helps distribute heat evenly and absorb moisture.
Chemistry Connection
This shows two Topic 5 concepts together: (1) the iron oxidation is exothermic (releases heat to the astronaut's hands), and (2) the salt lowers the activation energy (acting as a catalyst) so the reaction proceeds fast enough to be useful, even at very low temperatures.
Practice Questions: 5.1
20 multiple choice questions. Click an option to check your answer.
Your Score
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Question 1
In an exothermic reaction, the temperature of the surroundings:
In an exothermic reaction, energy is released to the surroundings, so the temperature of the surroundings increases.
Question 2
Which of the following is an endothermic reaction?
Thermal decomposition requires continuous heating - it absorbs energy from the surroundings, making it endothermic. Combustion, neutralisation, and respiration are all exothermic.
Question 3
The enthalpy change (ΔH) for an exothermic reaction is:
Exothermic reactions release energy, so the products have less energy than the reactants. This means ΔH is negative (e.g. ΔH = −890 kJ/mol for methane combustion).
Question 4
On an energy level diagram for an endothermic reaction, the products are drawn:
In an endothermic reaction, energy is absorbed, so the products end up at a higher energy level than the reactants. The difference represents the positive ΔH.
Question 5
Photosynthesis is endothermic because:
Photosynthesis absorbs light energy from the Sun to convert CO₂ and H₂O into glucose and O₂. Since energy is absorbed from the surroundings, it is endothermic.
Question 6
A student dissolves ammonium nitrate in water. The beaker feels cold. Which statement is correct?
The beaker feels cold because the dissolving process absorbed heat energy from the surroundings (including the student's hand). Energy flowed from the surroundings into the reaction, making it endothermic.
Question 7
Which of these reactions is exothermic?
Respiration (C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O) releases energy for life processes. It is the reverse of photosynthesis and is always exothermic.
Question 8
What is the activation energy of a reaction?
Activation energy (Eₐ) is the minimum energy that colliding particles must have in order for a reaction to occur. It is shown as the "hump" on the energy level diagram.
Question 9
A catalyst works by:
A catalyst provides an alternative reaction pathway with a lower activation energy. It does NOT change the overall ΔH or make the reaction more exothermic. It is not used up in the reaction.
Question 10
When a catalyst is added to a reaction, the ΔH of the reaction:
A catalyst does NOT change ΔH. The overall energy change of the reaction is exactly the same. The catalyst only lowers the activation energy, allowing the reaction to occur faster.
Question 11
Which energy level diagram represents an exothermic reaction?
In an exothermic reaction, energy is released, so products are at a lower energy level than reactants, and ΔH is negative.
Question 12
Burning wood in a fireplace is which type of reaction?
Burning (combustion) is always exothermic. While it needs activation energy (the match) to start, the overall reaction releases much more energy than it absorbs. The room gets warmer!
Question 13
In an energy level diagram, the activation energy is measured from:
Activation energy (Eₐ) is the energy needed to get the reaction started. It is measured from the energy level of the reactants up to the peak (top) of the energy barrier.
Question 14
An instant hot pack uses the reaction between iron and oxygen. This is an example of:
Iron reacting with oxygen (4Fe + 3O₂ → 2Fe₂O₃) is an oxidation reaction that is exothermic - it releases heat. This is why hand warmers use iron powder.
Question 15
Which statement about catalysts is INCORRECT?
Catalysts do NOT increase the amount of product formed - they only make the reaction reach equilibrium faster. The yield remains the same. A, B, and D are all correct statements about catalysts.
Question 16
A reaction has ΔH = +56 kJ/mol. This means:
A positive ΔH means the reaction is endothermic - 56 kJ of energy is absorbed from the surroundings per mole of reactant. The products are at a higher energy level than the reactants.
Question 17
During neutralisation, the temperature of the solution:
Neutralisation (acid + alkali → salt + water) is always exothermic. The temperature of the solution rises because energy is released from the reaction to the surroundings.
Question 18
Why does thermal decomposition stop when heating is removed?
Thermal decomposition is endothermic - it needs a continuous input of heat energy. Once you stop heating, there is no energy input, so the reaction cannot continue.
Question 19
Which pair correctly matches the reaction to its type?
Combustion releases heat (exothermic). Photosynthesis absorbs light energy (endothermic). This is the correct pairing. Respiration is exothermic, and neutralisation is exothermic too.
Question 20
On an energy level diagram, a catalyst is shown as:
A catalyst provides an alternative pathway with lower activation energy. On the diagram, this appears as a second, smaller hump between reactants and products, while the original higher hump is also shown for comparison.
5.2 Calculating Energy Changes Using Bond Energies
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Supplement - Extended Tier Only
This entire section is part of the Extended syllabus (Supplement content). It builds on the ideas from Section 5.1 and introduces quantitative calculations using bond energy data.
What is Bond Energy?
Bond energy (also called bond dissociation energy) is the amount of energy needed to break one mole of a particular covalent bond. It is measured in kJ/mol (kilojoules per mole).
For example, the bond energy of the O–H bond is 463 kJ/mol. This means it takes 463 kJ of energy to break one mole of O–H bonds (that is, to break 6.02 × 10²³ individual O–H bonds).
The Two Key Rules
| Process | Energy Change | Type |
|---|---|---|
| Breaking bonds | Energy is absorbed (required) | Endothermic |
| Making bonds | Energy is released | Exothermic |
Memory Trick
BREAK = TAKE (energy in) and MAKE = RELEASE (energy out). Breaking bonds needs energy input (like snapping a stick requires effort). Making bonds releases energy (like magnets snapping together releases energy).
Think about it this way: bonds are like springs holding atoms together. To pull atoms apart (break the bond), you need to put energy in. When atoms come together and form a new bond, energy is released.
The Bond Energy Calculation Formula
ΔH = Energy to break bonds − Energy released making bonds
Or equivalently:
ΔH = Σ(bond energies of reactants) − Σ(bond energies of products)
Σ means "sum of"
If ΔH is negative → exothermic
If ΔH is positive → endothermic
Why does this formula work?
- First, ALL bonds in the reactants must be broken. This requires energy (endothermic step).
- Then, NEW bonds in the products must be formed. This releases energy (exothermic step).
- The overall ΔH is the difference between these two values.
- If more energy is released making bonds than was absorbed breaking bonds → ΔH is negative → exothermic
- If more energy is absorbed breaking bonds than is released making bonds → ΔH is positive → endothermic
Common Bond Energies Reference Table
You will be given bond energy data in the exam, but it helps to know the common values:
| Bond | Bond Energy (kJ/mol) | Bond | Bond Energy (kJ/mol) |
|---|---|---|---|
| C–H | 412 | O=O | 496 |
| C–C | 348 | C=O (in CO₂) | 803 |
| C=C | 612 | O–H | 463 |
| C–O | 360 | H–H | 436 |
| C–N | 305 | N–H | 388 |
| C–Cl | 338 | H–Cl | 431 |
| H–F | 567 | H–Br | 366 |
| N≡N | 944 | Cl–Cl | 242 |
| H–I | 298 | Br–Br | 193 |
Step-by-Step Method for Bond Energy Calculations
- Write out the balanced equation with structural formulae (showing all bonds)
- List ALL bonds broken in the reactants and multiply each by its bond energy
- Add up the total energy needed to break all bonds
- List ALL bonds formed in the products and multiply each by its bond energy
- Add up the total energy released when all bonds are formed
- Calculate ΔH = Total energy to break bonds − Total energy to form bonds
- State whether the reaction is exothermic (ΔH negative) or endothermic (ΔH positive)
Worked Examples
Worked Example 1
Calculate the enthalpy change for the complete combustion of methane. CH₄ + 2O₂ → CO₂ + 2H₂O. Bond energies: C–H = 412, O=O = 496, C=O = 803, O–H = 463 kJ/mol.
Step 1: Draw the Structural Formulae
Reactants: H–C(–H)(–H)–H + O=O + O=O
CH₄ has 4 × C–H bonds. 2O₂ has 2 × O=O bonds.
Products: O=C=O + H–O–H + H–O–H
CO₂ has 2 × C=O bonds. 2H₂O has 4 × O–H bonds.
CH₄ has 4 × C–H bonds. 2O₂ has 2 × O=O bonds.
Products: O=C=O + H–O–H + H–O–H
CO₂ has 2 × C=O bonds. 2H₂O has 4 × O–H bonds.
Step 2: Calculate Energy to Break Bonds (Reactants)
| Bond | Number | Bond Energy | Total |
|---|---|---|---|
| C–H | 4 | 412 | 4 × 412 = 1648 kJ |
| O=O | 2 | 496 | 2 × 496 = 992 kJ |
| Total energy to break bonds | 1648 + 992 = 2640 kJ | ||
Step 3: Calculate Energy Released Making Bonds (Products)
| Bond | Number | Bond Energy | Total |
|---|---|---|---|
| C=O | 2 | 803 | 2 × 803 = 1606 kJ |
| O–H | 4 | 463 | 4 × 463 = 1852 kJ |
| Total energy released making bonds | 1606 + 1852 = 3458 kJ | ||
Step 4: Calculate ΔH
ΔH = Energy to break bonds − Energy to make bonds
ΔH = 2640 − 3458 = −818 kJ/mol
ΔH = 2640 − 3458 = −818 kJ/mol
ΔH = −818 kJ/mol. The reaction is exothermic because ΔH is negative. More energy is released making bonds in the products than is needed to break bonds in the reactants. (Note: the actual value is −890 kJ/mol. Bond energy calculations give approximate values because bond energies are averages.)
Worked Example 2
Calculate ΔH for the reaction: 2H₂ + O₂ → 2H₂O. Bond energies: H–H = 436, O=O = 496, O–H = 463 kJ/mol.
Step 1: Identify All Bonds
Reactants: 2 × H–H + 1 × O=O
Products: 2 × H₂O, each with 2 × O–H bonds = 4 × O–H total
Products: 2 × H₂O, each with 2 × O–H bonds = 4 × O–H total
Step 2: Energy to Break Bonds
| Bond | Number | Bond Energy | Total |
|---|---|---|---|
| H–H | 2 | 436 | 2 × 436 = 872 kJ |
| O=O | 1 | 496 | 1 × 496 = 496 kJ |
| Total energy to break bonds | 872 + 496 = 1368 kJ | ||
Step 3: Energy Released Making Bonds
| Bond | Number | Bond Energy | Total |
|---|---|---|---|
| O–H | 4 | 463 | 4 × 463 = 1852 kJ |
| Total energy released making bonds | 1852 kJ | ||
Step 4: Calculate ΔH
ΔH = 1368 − 1852 = −484 kJ/mol
ΔH = −484 kJ/mol. The reaction is exothermic. The energy released forming 4 O–H bonds (1852 kJ) is greater than the energy needed to break 2 H–H bonds and 1 O=O bond (1368 kJ). This is why hydrogen is being explored as a clean fuel - it releases a lot of energy when burned!
Worked Example 3
Calculate ΔH for: H₂ + Cl₂ → 2HCl. Bond energies: H–H = 436, Cl–Cl = 242, H–Cl = 431 kJ/mol.
Step 1: Identify All Bonds
Reactants: 1 × H–H + 1 × Cl–Cl
Products: 2 × H–Cl
Products: 2 × H–Cl
Step 2: Energy to Break Bonds
H–H: 1 × 436 = 436 kJ
Cl–Cl: 1 × 242 = 242 kJ
Total = 436 + 242 = 678 kJ
Cl–Cl: 1 × 242 = 242 kJ
Total = 436 + 242 = 678 kJ
Step 3: Energy Released Making Bonds
H–Cl: 2 × 431 = 862 kJ
Total = 862 kJ
Total = 862 kJ
Step 4: Calculate ΔH
ΔH = 678 − 862 = −184 kJ/mol
ΔH = −184 kJ/mol. The reaction is exothermic. The two H–Cl bonds formed (862 kJ released) are stronger than the H–H and Cl–Cl bonds broken (678 kJ absorbed), so the overall reaction releases energy.
Worked Example 4
Calculate ΔH for the combustion of ethane: C₂H₆ + 3½O₂ → 2CO₂ + 3H₂O. Bond energies: C–H = 412, C–C = 348, O=O = 496, C=O = 803, O–H = 463 kJ/mol. (Use 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O to avoid fractions, then halve your answer.)
Step 1: Identify All Bonds (for 1 mole of C₂H₆)
Reactants (C₂H₆ + 3½O₂):
C₂H₆ has: 1 × C–C bond + 6 × C–H bonds
3½ O₂ has: 3.5 × O=O bonds
Products (2CO₂ + 3H₂O):
2CO₂ has: 4 × C=O bonds
3H₂O has: 6 × O–H bonds
C₂H₆ has: 1 × C–C bond + 6 × C–H bonds
3½ O₂ has: 3.5 × O=O bonds
Products (2CO₂ + 3H₂O):
2CO₂ has: 4 × C=O bonds
3H₂O has: 6 × O–H bonds
Step 2: Energy to Break Bonds
| Bond | Number | Bond Energy | Total |
|---|---|---|---|
| C–C | 1 | 348 | 348 kJ |
| C–H | 6 | 412 | 6 × 412 = 2472 kJ |
| O=O | 3.5 | 496 | 3.5 × 496 = 1736 kJ |
| Total energy to break bonds | 348 + 2472 + 1736 = 4556 kJ | ||
Step 3: Energy Released Making Bonds
| Bond | Number | Bond Energy | Total |
|---|---|---|---|
| C=O | 4 | 803 | 4 × 803 = 3212 kJ |
| O–H | 6 | 463 | 6 × 463 = 2778 kJ |
| Total energy released making bonds | 3212 + 2778 = 5990 kJ | ||
Step 4: Calculate ΔH
ΔH = 4556 − 5990 = −1434 kJ/mol
ΔH = −1434 kJ/mol. The combustion of ethane is highly exothermic - even more so than methane, because ethane has more C–H bonds and a C–C bond to contribute to bond formation in the products. More bonds formed in the products = more energy released.
Worked Example 5
Calculate ΔH for the Haber process: N₂ + 3H₂ → 2NH₃. Bond energies: N≡N = 944, H–H = 436, N–H = 388 kJ/mol. Is the reaction exothermic or endothermic?
Step 1: Identify All Bonds
Reactants: 1 × N≡N (triple bond!) + 3 × H–H
Products: 2 × NH₃. Each NH₃ has 3 × N–H bonds, so total = 6 × N–H bonds
Products: 2 × NH₃. Each NH₃ has 3 × N–H bonds, so total = 6 × N–H bonds
Step 2: Energy to Break Bonds
N≡N: 1 × 944 = 944 kJ
H–H: 3 × 436 = 1308 kJ
Total = 944 + 1308 = 2252 kJ
H–H: 3 × 436 = 1308 kJ
Total = 944 + 1308 = 2252 kJ
Step 3: Energy Released Making Bonds
N–H: 6 × 388 = 2328 kJ
Total = 2328 kJ
Total = 2328 kJ
Step 4: Calculate ΔH
ΔH = 2252 − 2328 = −76 kJ/mol
ΔH = −76 kJ/mol. The Haber process is slightly exothermic. Despite the huge energy needed to break the N≡N triple bond (944 kJ), forming six strong N–H bonds releases enough energy (2328 kJ) to make the overall reaction exothermic - but only just!
Worked Example 6
Calculate ΔH for the decomposition of hydrogen iodide: 2HI → H₂ + I₂. Bond energies: H–I = 298, H–H = 436, I–I = 151 kJ/mol.
Step 1: Identify All Bonds
Reactants: 2 × H–I
Products: 1 × H–H + 1 × I–I
Products: 1 × H–H + 1 × I–I
Step 2: Energy to Break Bonds
H–I: 2 × 298 = 596 kJ
Total = 596 kJ
Total = 596 kJ
Step 3: Energy Released Making Bonds
H–H: 1 × 436 = 436 kJ
I–I: 1 × 151 = 151 kJ
Total = 436 + 151 = 587 kJ
I–I: 1 × 151 = 151 kJ
Total = 436 + 151 = 587 kJ
Step 4: Calculate ΔH
ΔH = 596 − 587 = +9 kJ/mol
ΔH = +9 kJ/mol. The reaction is endothermic (ΔH is positive). More energy is needed to break the two H–I bonds than is released by forming the H–H and I–I bonds. This makes sense - decomposition reactions are typically endothermic.
Why Bond Energy Calculations Give Approximate Values
Bond energy calculations give approximate values for ΔH. This is because:
- Bond energies given in data tables are average values. The exact energy of a C–H bond depends on what other atoms are bonded to the carbon. For example, the C–H bond in methane (CH₄) is slightly different from the C–H bond in ethanol (C₂H₅OH).
- The values assume that all bonds of the same type have the same energy, which is only approximately true.
- Bond energies apply to gaseous molecules. If reactants or products are liquids or solids, extra energy changes occur (like evaporation or condensation) that are not accounted for.
Interpreting Results: What Does the Sign Tell You?
| Calculation Result | Meaning | Explanation |
|---|---|---|
| ΔH is negative | Reaction is exothermic | More energy is released making new bonds in the products than is absorbed breaking bonds in the reactants. The extra energy goes to the surroundings as heat. |
| ΔH is positive | Reaction is endothermic | More energy is absorbed breaking bonds in the reactants than is released making new bonds in the products. The shortfall is taken from the surroundings. |
Linking Bond Energies to Energy Level Diagrams
The bond energy calculation connects directly to the energy level diagram:
Bond energy calculation visualised: bonds break first (energy absorbed), then new bonds form (energy released). The difference is ΔH.
The diagram shows that:
- Breaking all bonds takes the system to the highest energy point (separate atoms)
- Forming new bonds brings the energy back down
- If the products end up lower than the reactants, the reaction is exothermic (energy out > energy in)
- If the products end up higher than the reactants, the reaction is endothermic (energy in > energy out)
Exam Tips for 5.2
1. Count bonds carefully: The most common mistake is miscounting bonds. Draw out the structural formulae and count EVERY bond. For CH₄, there are 4 C–H bonds, not 1. For CO₂, there are 2 C=O bonds.
2. Remember: it is BREAK minus MAKE: ΔH = bonds broken − bonds made. If you do it the wrong way round, you will get the right number but the wrong sign, which means the wrong answer.
3. Use a table: Set up a clear table for bonds broken and bonds made. This avoids confusion and makes it easy for the examiner to follow your working (and award method marks).
4. Watch out for double and triple bonds: O=O is a double bond (496 kJ/mol). N≡N is a triple bond (944 kJ/mol). These are much stronger than single bonds. Do not confuse O=O with O–O.
5. C=O in CO₂ is special: The C=O bond in carbon dioxide (803 kJ/mol) is different from a C=O in other molecules. Always use the value given in the question.
6. Show all working: Even if you get the final answer wrong, you can still earn method marks for showing the correct process. Write out: bonds broken, bonds made, and the subtraction clearly.
7. Fractional moles are OK: If the balanced equation gives 3½ O₂, you can use 3.5 × 496 in your calculation. No need to double everything.
8. State the type: Always finish by stating whether the reaction is exothermic or endothermic based on the sign of ΔH.
Apply It: Real-World Chemistry
Bond energy calculations are not just exam exercises - they help engineers, chemists, and scientists predict the energy output of fuels, design safer processes, and develop new materials.
1
A chemical engineer at Indian Oil Corporation's R&D centre in Faridabad is comparing the energy output of different fuels for a new refinery process. She needs to decide between using methane (CH₄) and hydrogen (H₂) as fuel, comparing the energy released per gram.
Using bond energies, how would she calculate which fuel releases more energy per gram when burned?
The Calculation
From our worked examples: methane combustion gives ΔH ≈ −818 kJ/mol (Mᵣ = 16 g/mol) = 51.1 kJ/g. Hydrogen combustion: ΔH ≈ −484 kJ for 2 mol H₂ (Mᵣ = 2 g/mol), so per mole = −242 kJ/mol = 121 kJ/g. Hydrogen releases about 2.4 times more energy per gram!
Why It Matters
This is exactly why hydrogen is considered a superior fuel for rockets and is being explored for fuel cells. Per gram, it delivers far more energy than any hydrocarbon fuel. The engineer would factor in storage costs and safety too, but the bond energy calculation gives the first critical comparison.
Chemistry Connection
Bond energy calculations let engineers predict energy output BEFORE building expensive equipment. The engineer uses the exact same formula you learn at IGCSE: ΔH = bonds broken − bonds made.
2
A materials scientist at the University of Cambridge is developing a new type of single-use hand warmer for hikers. She is testing different metal powder combinations. She knows that the oxidation of iron releases energy, but she wants to calculate exactly how much energy the iron-oxygen reaction produces using bond energies of the metallic and oxide bonds.
Why do bond energy calculations help her predict the temperature rise in the hand warmer before building prototypes?
The Approach
By calculating ΔH for the oxidation reaction, the scientist knows how many kilojoules of heat will be released per gram of iron powder. She can then use the equation Q = mcΔT (from physics) to predict how hot the pack will get and for how long.
Practical Application
If the calculation shows 100 kJ per hand warmer, and the pack has a mass of 40 g with a specific heat capacity similar to water, she can predict a temperature rise. This saves weeks of trial-and-error prototyping.
Chemistry Connection
Bond energy calculations are a predictive tool. Engineers calculate energy changes on paper (or computer) before spending money on materials and testing. The sign and magnitude of ΔH tell you if a reaction is suitable for your application.
3
A rocket propulsion engineer at ISRO's Liquid Propulsion Systems Centre in Thiruvananthapuram is evaluating liquid hydrogen (H₂) and liquid oxygen (O₂) as rocket propellants. The GSLV Mk III rocket uses this fuel combination for its upper stage.
Using bond energy data, explain why the H₂/O₂ fuel combination is so effective for rockets, and calculate the energy released per mole.
The Calculation
2H₂ + O₂ → 2H₂O. Bonds broken: 2(H–H) = 872 kJ + 1(O=O) = 496 kJ = 1368 kJ. Bonds made: 4(O–H) = 1852 kJ. ΔH = 1368 − 1852 = −484 kJ. This is −242 kJ per mole of H₂.
Why H₂/O₂ is Ideal for Rockets
The H₂/O₂ combination is ideal because: (1) hydrogen has the highest energy per gram of any fuel (121 kJ/g); (2) the only product is water - no harmful emissions; (3) the O–H bonds formed are very strong (463 kJ/mol each), releasing a large amount of energy.
Chemistry Connection
ISRO uses the same bond energy principles you study at IGCSE! The reason hydrogen-oxygen is the best rocket fuel comes down to bond energies: the O–H bonds in water are extremely strong, so forming them releases enormous energy.
4
A food scientist at Nestle's research facility in Lausanne, Switzerland, is calculating the calorific value (energy content) of a new protein bar. She needs to estimate how much energy the body can extract from the fats and carbohydrates in the bar by considering the combustion reactions of these nutrients.
How do bond energy calculations relate to the Calories listed on food packaging?
The Connection
When your body metabolises food (through respiration), it breaks the C–H, C–C, and C–O bonds in nutrients and forms C=O bonds (in CO₂) and O–H bonds (in H₂O). The ΔH of these reactions tells you how much energy is available. Fats have more C–H bonds per gram than carbohydrates, which is why fats contain more energy (9 kcal/g vs 4 kcal/g).
From Bond Energies to Food Labels
The "Calories" on a food label are essentially the ΔH of the combustion reaction of that food. Scientists measure this using a bomb calorimeter, but they could estimate it using bond energies from the molecular structures of the fats, proteins, and carbohydrates.
Chemistry Connection
Every time you read a nutrition label, you are looking at the result of a ΔH calculation! The energy your body extracts from food is the difference between the energy needed to break bonds in nutrients and the energy released forming bonds in CO₂ and H₂O.
5
An environmental chemist at the National Oceanic and Atmospheric Administration (NOAA) in Boulder, Colorado, is studying ozone depletion. She is calculating the bond energy changes when CFCs (chlorofluorocarbons) break apart in the upper atmosphere and the chlorine atoms attack ozone (O₃) molecules.
How do bond energies explain why the C–Cl bond in CFCs breaks more easily than other bonds, leading to ozone destruction?
The Bond Energy Analysis
The C–Cl bond energy is 338 kJ/mol, which is relatively weak compared to C–F (484 kJ/mol) or C–H (412 kJ/mol). UV radiation in the upper atmosphere provides enough energy to break the C–Cl bond, releasing free chlorine atoms (Cl radicals).
Why Ozone is Destroyed
The free Cl atom then reacts with ozone: Cl + O₃ → ClO + O₂. The bond energies show this reaction is exothermic - the new Cl–O bond (218 kJ/mol) plus O=O bond (496 kJ/mol) release more energy than is needed to break the bonds in ozone. One Cl atom can destroy thousands of O₃ molecules.
Chemistry Connection
Bond energy data explains environmental chemistry! The relatively weak C–Cl bond is the reason CFCs are so damaging - UV light can break it. This is why the Montreal Protocol banned CFCs. Understanding bond energies helps us understand and solve environmental problems.
Practice Questions: 5.2
20 multiple choice questions. Click an option to check your answer.
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Question 1
What is bond energy?
Bond energy is defined as the energy needed to break one mole of a particular covalent bond. Breaking bonds requires energy input (endothermic process).
Question 2
Breaking bonds is:
Breaking bonds ALWAYS requires energy input - it is ALWAYS endothermic. Think of it like pulling apart two magnets - you need to put effort (energy) in.
Question 3
Making bonds is:
Making (forming) bonds ALWAYS releases energy - it is ALWAYS exothermic. When atoms come together to form a bond, they become more stable and release excess energy.
Question 4
The formula for calculating ΔH using bond energies is:
ΔH = total energy to break bonds (in reactants) − total energy released making bonds (in products). This gives a negative value for exothermic and positive for endothermic reactions.
Question 5
For the reaction H₂ + Cl₂ → 2HCl, given H–H = 436, Cl–Cl = 242, H–Cl = 431 kJ/mol, what is ΔH?
Bonds broken: 436 + 242 = 678 kJ. Bonds made: 2 × 431 = 862 kJ. ΔH = 678 − 862 = −184 kJ/mol. Negative, so exothermic.
Question 6
How many C–H bonds need to be broken in the combustion of methane (CH₄)?
Methane (CH₄) has one carbon atom bonded to four hydrogen atoms. Each C–H bond must be broken, so there are 4 C–H bonds to break.
Question 7
If ΔH for a reaction is −350 kJ/mol, it means:
A negative ΔH means exothermic. Since ΔH = bonds broken − bonds made, a negative result means the energy released making bonds (bonds made) exceeds the energy absorbed breaking bonds (bonds broken).
Question 8
In the combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O), how many O–H bonds are formed in the products?
The products include 2H₂O. Each water molecule has 2 O–H bonds. So 2 × 2 = 4 O–H bonds are formed.
Question 9
Why do bond energy calculations give approximate values for ΔH?
Bond energies listed in data tables are average values. The exact energy of a C–H bond varies slightly depending on the molecule it is in. This means calculations using average bond energies give approximate results.
Question 10
For the reaction N₂ + 3H₂ → 2NH₃, how many N–H bonds are formed?
2NH₃ molecules are produced. Each NH₃ has 3 N–H bonds, so the total number of N–H bonds formed = 2 × 3 = 6.
Question 11
Given: A–B bond energy = 300 kJ/mol, C–D bond energy = 400 kJ/mol, A–D bond energy = 500 kJ/mol, C–B bond energy = 350 kJ/mol. For the reaction AB + CD → AD + CB, what is ΔH?
Bonds broken: A–B (300) + C–D (400) = 700 kJ. Bonds made: A–D (500) + C–B (350) = 850 kJ. ΔH = 700 − 850 = −150 kJ/mol. Exothermic.
Question 12
In the reaction 2H₂ + O₂ → 2H₂O, if H–H = 436, O=O = 496, O–H = 463 kJ/mol, the total energy needed to break all bonds in the reactants is:
Bonds broken: 2(H–H) = 2 × 436 = 872 kJ, plus 1(O=O) = 496 kJ. Total = 872 + 496 = 1368 kJ.
Question 13
Which bond requires the most energy to break?
The N≡N triple bond has the highest bond energy at 944 kJ/mol. Triple bonds are the strongest because three shared pairs of electrons hold the atoms together very tightly. This is why nitrogen gas is so unreactive.
Question 14
A reaction has: bonds broken = 1200 kJ, bonds made = 1500 kJ. The reaction is:
ΔH = 1200 − 1500 = −300 kJ/mol. The negative sign means it is exothermic. More energy is released making bonds (1500 kJ) than is absorbed breaking bonds (1200 kJ).
Question 15
In CO₂ (O=C=O), how many C=O bonds are there?
Carbon dioxide (CO₂) has the structure O=C=O. There are 2 C=O double bonds. This is important in combustion calculations where CO₂ is a product.
Question 16
For the reaction CH₄ + 2O₂ → CO₂ + 2H₂O, the total energy released making bonds in the products is (C=O = 803, O–H = 463 kJ/mol):
Products: CO₂ has 2 C=O bonds (2 × 803 = 1606 kJ). 2H₂O has 4 O–H bonds (4 × 463 = 1852 kJ). Total = 1606 + 1852 = 3458 kJ.
Question 17
A reaction has: bonds broken = 800 kJ, bonds made = 650 kJ. The reaction is:
ΔH = 800 − 650 = +150 kJ/mol. The positive sign means it is endothermic. More energy is absorbed breaking bonds (800 kJ) than is released making bonds (650 kJ).
Question 18
Which statement about the N≡N bond in the Haber process (N₂ + 3H₂ → 2NH₃) is correct?
The N≡N triple bond is very strong (944 kJ/mol). It must be broken for the Haber process to occur, which is why the process requires high temperatures (450 °C), high pressure (200 atm), and an iron catalyst.
Question 19
Given: H–H = 436, F–F = 158, H–F = 567 kJ/mol. Calculate ΔH for: H₂ + F₂ → 2HF.
Bonds broken: H–H (436) + F–F (158) = 594 kJ. Bonds made: 2 × H–F (2 × 567) = 1134 kJ. ΔH = 594 − 1134 = −540 kJ/mol. Highly exothermic because the H–F bond is very strong.
Question 20
A reaction is exothermic. Which statement must be true?
For a reaction to be exothermic, more energy must be released when forming bonds in the products than is absorbed when breaking bonds in the reactants. This gives a negative ΔH. All reactions still need activation energy - even exothermic ones.
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Cambridge IGCSE Chemistry (0620) | 2026-2028 Syllabus | Extended Tier
Topic 5: Chemical Energetics - Complete Study Guide