5. Chemical Energetics
Candidates should be able to:
- understand that chemical reactions are accompanied by enthalpy changes and these changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive)
- construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy
- define and use the terms:
- standard conditions (this syllabus assumes that these are 298 K and 101 kPa) shown by ⦵
- enthalpy change with particular reference to: reaction, ΔHr, formation, ΔHf, combustion, ΔHc, neutralisation, ΔHneut
- understand that energy transfers occur during chemical reactions because of the breaking and making of chemical bonds
- use bond energies (ΔH positive, i.e. bond breaking) to calculate the enthalpy change of reaction, ΔHr
- understand that some bond energies are exact and some bond energies are averages
- calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mcΔT and ΔH = −mcΔT/n
Think of two reactions or processes from everyday life that give out heat, and two that take in heat. Keep them in mind — you’ll be able to explain exactly why by the end of this topic.
1 Exothermic and Endothermic Reactions
Almost all chemical reactions are accompanied by energy changes, because bonds are broken in the reactants and new bonds are formed in the products. Breaking bonds needs energy; forming bonds releases energy. It’s very unlikely that these two amounts of energy will be exactly equal, so overall, energy is either absorbed or released during a reaction.
Usually this energy change involves heat, but it can also appear as sound, light, or electrical energy.
A reaction that gives out heat to the surroundings. Heat is lost from the reactants, so the value of ΔH is negative.
A reaction that takes in heat from the surroundings. Heat is gained by the reactants, so the value of ΔH is positive.
The heat energy transferred during a chemical reaction, measured in kJ mol−1.
Reaction pathway diagrams
A reaction pathway diagram plots enthalpy against the progress of the reaction. It shows the enthalpy of the reactants, the enthalpy of the products, the overall enthalpy change ΔH, and the activation energy EA.
The minimum energy that colliding particles must possess to break bonds and start a chemical reaction.
Rules for drawing reaction pathway diagrams
The EA arrow starts at the level of the reactants and ends at the highest point of the energy hump (the transition state).
The ΔH arrow points downwards for an exothermic reaction and upwards for an endothermic reaction.
2 Standard Conditions and Types of Enthalpy Change
To make comparisons of enthalpy changes fair, chemists always quote them under the same set of conditions, called standard conditions:
- a pressure of 101 kPa (1.01 × 105 Pa — approximately normal atmospheric pressure)
- a temperature of 298 K (25 °C)
- each substance in the reaction is in its normal physical state (solid, liquid or gas) at 101 kPa and 298 K
The symbol ⦵ (a superscript Plimsoll symbol) shows that an enthalpy change refers to a reaction carried out under standard conditions, e.g. ΔH⦵.
A standard enthalpy change refers to the energy transferred at 298 K and standard pressure (101 kPa), with reactants and products in their standard states. There are several named standard enthalpy changes you need to know precisely:
The enthalpy change when the amounts of reactants shown in the stoichiometric equation react to give products, under standard conditions, with all reactants and products in their standard states.
The enthalpy change when one mole of a compound is formed from its elements, under standard conditions, with all reactants and products in their standard states. By definition, ΔHf⦵ of any element in its standard state is zero.
The enthalpy change when one mole of a substance is burnt completely in excess oxygen, under standard conditions, with all reactants and products in their standard states.
The enthalpy change when one mole of water is formed by the reaction of an acid with an alkali, under standard conditions.
3 Bond Energies and Calculating ΔH
Enthalpy changes happen because bonds are broken and new bonds are formed:
- Bond breaking requires energy to overcome the attractive forces between atoms — it is endothermic (ΔH positive).
- Bond forming releases energy — it is exothermic (ΔH negative).
In a chemical reaction: if the energy needed to break bonds is less than the energy released forming new bonds, the reaction releases energy overall and is exothermic. If the energy needed to break bonds is more than the energy released, the reaction absorbs energy overall and is endothermic.
Bond energy data
Bond energies in diatomic molecules are known exactly, because every molecule of that type is identical. Bond energies in polyatomic molecules are average values, because the exact energy of, say, a C−H bond varies slightly depending on which molecule it is in.
| Bond | Energy / kJ mol−1 |
|---|---|
| H−H | 436 |
| N≡N | 944 |
| O=O | 496 |
| P≡P | 485 |
| S=S | 425 |
| F−F | 158 |
| Cl−Cl | 242 |
| Br−Br | 193 |
| I−I | 151 |
| H−F | 562 |
| H−Cl | 431 |
| H−Br | 366 |
| H−I | 299 |
| C≡O | 1077 |
| Bond | Energy / kJ mol−1 |
|---|---|
| C−C | 350 |
| C=C | 610 |
| C≡C | 840 |
| C−H | 410 |
| C−Cl | 340 |
| C−Br | 280 |
| C−I | 240 |
| C−N | 305 |
| C=N | 610 |
| C≡N | 890 |
| C−O | 360 |
| C=O | 740 |
| C=O (in CO2) | 805 |
| N−H | 390 |
| O−H | 460 |
| N−N | 160 |
| O−O | 150 |
| S−H | 340 |
A fuller data set (Si, P and S environments) is provided in your data booklet — the values above are the ones most commonly used in calculations.
Use bond energies to calculate ΔH for the complete combustion of methane:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
This is close to, but not exactly, the experimental value of −890.3 kJ mol−1 shown in Section 1 — the difference arises because C−H and O−H bond energies used here are averages, not exact values for this specific molecule.
4 Measuring Enthalpy Changes Experimentally
Enthalpy changes can be found experimentally using calorimetry — typically by using the reaction to heat a known mass of water and measuring the temperature rise.
The energy needed to raise the temperature of 1 g of a substance by 1 °C (1 K).
To express this per mole of a stated reactant or product, divide by the number of moles, n, reacted:
The negative sign shows that heat released by the reaction (a temperature rise in the water) corresponds to an exothermic, negative ΔH.
0.50 g of ethanol, C2H5OH (M = 46 g mol−1), is burned and used to heat 100 g of water, raising its temperature from 21.0 °C to 43.4 °C. Specific heat capacity of water, c = 4.18 J g−1 °C−1. Calculate the enthalpy change of combustion of ethanol.
The negative sign is added because combustion is exothermic — heat is released by the reaction and absorbed by the water.
Simple calorimetry experiments like this almost always give a ΔHc that is less exothermic (less negative) than the true, data-book value — mainly because of heat loss to the surroundings and the container, and because combustion may be incomplete.
Practice questions
State the meaning of the term “standard enthalpy change of formation”.
Write your definition, then reveal the mark scheme.
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The enthalpy change when one mole of a compound is formed from its elements (1), under standard conditions (298 K, 101 kPa) with all substances in their standard states (1).
Sketch a reaction pathway diagram for an endothermic reaction and label ΔH and EA.
Sketch your diagram, then check the key features below.
Show mark scheme
Products drawn higher than reactants (1); EA arrow starting at reactants and ending at the top of the hump (1); ΔH arrow pointing upwards from reactants level to products level (1).
Use the bond energies provided to calculate ΔH for: H2(g) + Cl2(g) → 2HCl(g). (H−H = 436, Cl−Cl = 242, H−Cl = 431 kJ mol−1)
Show your working, then reveal the mark scheme.
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Bonds broken: (436 + 242) = 678 kJ mol−1 (1)
Bonds made: 2 × 431 = 862 kJ mol−1 (1)
ΔH = 678 − 862 = −184 kJ mol−1 (1)
0.020 mol of a fuel raises the temperature of 150 g of water by 18.5 °C. Calculate ΔH for this reaction in kJ mol−1. (c = 4.18 J g−1 °C−1)
Show your working, then reveal the mark scheme.
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q = mcΔT = 150 × 4.18 × 18.5 = 11 600 J (3 s.f.) (1)
ΔH = −q / n = −11 600 / 0.020 = −579 000 J mol−1 (1)
ΔH = −579 kJ mol−1 (1)
Explain, in terms of bond breaking and bond making, why some reactions are exothermic and others endothermic.
Write your explanation, then reveal the mark scheme.
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Breaking bonds requires (absorbs) energy; this is endothermic (1). Forming bonds releases energy; this is exothermic (1). If more energy is released forming bonds than is needed to break bonds, the reaction is exothermic overall (and vice versa for endothermic) (1).
Exam Tips & Quick Summary
Do remember
- Exothermic = ΔH negative, heat given out; endothermic = ΔH positive, heat taken in.
- Bond breaking is endothermic; bond forming is exothermic.
- ΔH arrow: down for exothermic, up for endothermic. EA arrow always starts at reactants and ends at the peak.
- Standard conditions: 298 K, 101 kPa, substances in their normal physical states — shown by the symbol ⦵.
- ΔHf of an element in its standard state is always zero.
Do not confuse
- ΔHf (formed from elements) with ΔHr (from any stated reactants).
- ΔHc needs excess oxygen and one mole of the substance burnt — not one mole of oxygen.
- ΔHneut is defined per mole of water formed, not per mole of acid or alkali.
- Bond energy values in the table are always positive (energy to break that bond) — the sign of ΔH comes from subtracting bonds made from bonds broken.
- Diatomic bond energies are exact; bond energies in bigger molecules are only averages.
Recap in 20 seconds
Reactions absorb or release energy because bond breaking needs energy (endothermic) while bond making releases it (exothermic) — the balance between the two gives the overall ΔH. Standard enthalpy changes (ΔHr, ΔHf, ΔHc, ΔHneut) are measured at 298 K and 101 kPa, and can be calculated from bond energies or measured experimentally using q = mcΔT and ΔH = −mcΔT/n.