Energy and chemical bonds

Why reactions release or absorb energy: breaking bonds takes energy in and making bonds gives energy out, calculating enthalpy change (ΔH) from bond energies, reading energy profiles, making fire by friction, measuring the energy of neutralisation, thermal decomposition of limestone, and how heating changes yellow ochre into red.

Light a gas stove and its blue flame can boil a litre of water in a few minutes. Yet before methane can burn, every molecule has to be torn apart, and pulling atoms apart always costs energy. So where does all that heat come from? Not from the fuel's bonds breaking, but from the new, stronger bonds that form as the flame makes carbon dioxide and water.

Breaking and making bonds

In Year 8 you learned that exothermic reactions give out energy and endothermic reactions take it in. Now you can explain why, using the bonds that hold atoms together.

Breaking a bond always takes energy in. Atoms are held together by the attraction between their nuclei and shared or transferred electrons. Pulling them apart is work, like pulling two magnets apart.
Making a bond always gives energy out. When atoms snap together, the same amount of energy is released, usually as heat.

Every reaction does both: bonds in the reactants break, then new bonds form in the products. Whether the reaction is exothermic or endothermic depends on which is bigger.

	Exothermic	Endothermic
Bonds made compared with bonds broken	More energy released making bonds than taken in breaking them	More energy taken in breaking bonds than released making them
Energy stored in chemicals	Products store less than reactants	Products store more than reactants
Surroundings	Get hotter	Get colder (or energy must be supplied)
Sign of ΔH	Negative (energy lost by the chemicals)	Positive (energy gained by the chemicals)
Examples	Combustion, respiration, neutralisation, rusting	Photosynthesis, thermal decomposition, electrolysis, cooking an egg

The energy change of a reaction is called the enthalpy change, written ΔH ("delta H"; Δ means "change in"). It is measured in kilojoules (kJ), usually for a fixed amount of reactant called a mole. A mole is just a chemist's counting unit, like a dozen but much bigger: 6.02 × 10²³ particles. You'll use moles properly in Stage 6; for now, read "per mole" as "for the amounts shown in the equation".

Calculating ΔH from bond energies

A bond energy is the energy needed to break one mole of a particular bond. The same amount is released when that bond forms. These values are averages measured across many compounds.

Bond	Energy (kJ/mol)	Bond	Energy (kJ/mol)
H–H	436	C–H	413
O=O	498	C=O (in CO₂)	805
O–H	463	C–C	346
Cl–Cl	242	C–O	358
H–Cl	431	N≡N	945
H–I	299	N–H	391
I–I	151	N=O (in NO)	630

The rule is:

ΔH = energy taken in to break bonds − energy given out making bonds

Worked example: hydrogen and chlorine

H₂ + Cl₂ → 2HCl. Draw the bonds: H–H + Cl–Cl → H–Cl + H–Cl.

The reaction

One H–H (436) and one Cl–Cl (242): 436 + 242 = 678 kJ.

Bonds broken (energy in)

Two H–Cl: 2 × 431 = 862 kJ.

Bonds made (energy out)

ΔH = 678 − 862 = −184 kJ. More energy comes out than goes in, so the reaction is exothermic. The measured value is −185 kJ, very close.

Subtract

Burning methane. Breaking all the bonds in CH₄ and 2O₂ takes in 2648 kJ; forming the bonds in CO₂ and 2H₂O gives out 3462 kJ. The difference, 814 kJ, is released.

For methane, the natural gas in a gas stove: CH₄ + 2O₂ → CO₂ + 2H₂O. Breaking 4 C–H and 2 O=O takes in 4 × 413 + 2 × 498 = 2648 kJ. Making 2 C=O and 4 O–H gives out 2 × 805 + 4 × 463 = 3462 kJ. So ΔH = 2648 − 3462 = −814 kJ per mole of methane. (The measured value with the water as steam is −802 kJ. The small difference is because bond energies are averages.)

Notice that energy is not stored "in" the fuel's bonds waiting to escape. Breaking bonds costs energy. The energy is released because the bonds in carbon dioxide and water are stronger than the bonds in methane and oxygen. A good fuel is one whose products have much stronger bonds than the fuel and oxygen did.

Reading energy profiles

Energy profiles. The height of the hump is the activation energy (Eₐ). The difference between the reactant and product levels is ΔH. A catalyst lowers the hump but not the levels.

An energy profile (energy level diagram) plots the energy stored in the chemicals as the reaction goes from reactants to products.

The activation energy (Eₐ) is the energy needed to start breaking bonds so the reaction can get going. It is measured from the reactants up to the top of the hump. Even very exothermic reactions, like the hydrogen balloon, need it, which is why fuels don't burst into flame on their own.
The enthalpy change (ΔH) is measured from the reactant level to the product level. Downwards is negative (exothermic); upwards is positive (endothermic).
A catalyst provides a different pathway with a lower hump. It changes Eₐ but never ΔH, because the reactants and products are the same.

Going further: Fire from friction

For tens of thousands of years, Aboriginal people across Australia have made fire by friction. In the fire-drill method, a straight, light stick, often the dried flower spike of a grass tree (Xanthorrhoea), is spun quickly between the palms against a notch in a flat piece of soft, dry wood. In the fire-saw method, the edge of a hard wooden tool such as a spear-thrower is rubbed rapidly across a split piece of softwood. Friction converts kinetic energy into heat until fine wood dust in the notch reaches its ignition temperature, supplying the activation energy. A glowing ember forms, and is tipped into dry grass and blown on to provide more oxygen. From then on the exothermic combustion supplies its own activation energy, particle by particle. Fire was also carried from place to place as slowly smouldering fire sticks, a controlled, slow combustion that could be fanned back into flame when needed.

InteractiveBreak and make the bonds

Try this

Choose H₂ + Cl₂. Break the bonds, then make the new ones. Check that ΔH is −184 kJ, as in the worked example.
Try every reaction. Which three are endothermic? What do they have in common?
Compare "Water split" with "H₂ + O₂". What do you notice about the two values of ΔH?
Methane releases 814 kJ per mole and ethanol 1272 kJ. Now look at the energy per gram shown for each. Which is the better fuel for the same mass, and why is it different from the per-mole answer?
N₂ + O₂ → 2NO is endothermic. Use the bond energies to explain why this reaction only happens at the very high temperatures inside a car engine or a lightning bolt.

Experiment: measure the energy of neutralisation

Aim: to measure the energy released when an acid neutralises a base, and compare it with the accepted value.

Method: Nest two polystyrene cups inside each other and stand them in a beaker so they can't tip. Measure 25 mL of 1.0 mol/L hydrochloric acid into the cup and record its temperature every 30 seconds for 2 minutes. Measure 25 mL of 1.0 mol/L sodium hydroxide solution at the same starting temperature. Add it all at once, put a cardboard lid with a hole for the thermometer on top, stir gently and record the temperature every 15 seconds until it has clearly started to fall. Wear safety glasses: both solutions irritate the eyes and skin.

Results: plot temperature against time and find the temperature rise, ΔT (typically about 6 to 7 °C).

Calculation: the energy given to the solution is Q = m × c × ΔT, where m is the mass of solution (about 50 g, since 1 mL of a dilute solution has a mass of about 1 g) and c = 4.18 J/g/°C, the specific heat capacity of water. For ΔT = 6.7 °C, Q = 50 × 4.18 × 6.7 = 1400 J. The reaction made 0.025 mol of water (25 mL of 1.0 mol/L acid), so ΔH ≈ −1400 J ÷ 0.025 mol = −56 000 J per mole, or about −56 kJ per mole of water formed. (The accepted value is −57 kJ/mol.)

Discussion: why is ΔH negative when the temperature went up? Where could heat be lost, and would that make your value too big or too small? Why use polystyrene and a lid rather than a glass beaker? Repeat with ethanoic acid (a weak acid) instead of hydrochloric acid and compare.

Endothermic reactions at work

Thermal decomposition breaks a compound into simpler substances using heat. It is always endothermic, because bonds are broken and the products are less stable on their own.

The most important example is turning limestone into quicklime (calcium oxide) in a lime kiln at about 900 °C:

CaCO₃ → CaO + CO₂ ΔH = +178 kJ

Quicklime is used to make cement and steel, to treat drinking water and to neutralise acidic soils and mine waste. Adding water to quicklime is strongly exothermic (it makes slaked lime, Ca(OH)₂, and can boil the water). Cement making releases carbon dioxide twice over: from the fuel burned to heat the kiln, and from the limestone itself. Together these make cement responsible for about 8% of the world's carbon dioxide emissions, so chemists are developing lower-carbon cements and ways to capture the CO₂.

The stone walls and arched opening of an old lime kiln, with a timber fence in front — Ruins of a 19th-century lime kiln at Walkerville, Victoria. Limestone was burned here to make quicklime for building mortar.Photo: Damien Frawley, CC BY-SA 3.0

Ochre: chemistry on Country

Heating yellow ochre drives water out of the mineral goethite and turns it into red hematite. The colour change is a chemical change that needs energy.

Aboriginal and Torres Strait Islander peoples have used ochre for tens of thousands of years, for painting on rock, bark and bodies, in ceremony, as medicine, to protect skin and to tan hides. Ochres are natural earths coloured by iron compounds. Yellow ochre contains the mineral goethite, FeO(OH); red ochre contains hematite, Fe₂O₃.

Heating yellow ochre in a fire turns it red. The heat supplies the energy to break bonds and drive water out of the goethite, an endothermic decomposition: 2FeO(OH) → Fe₂O₃ + H₂O. Making red pigment this way is a practice known from many cultures, and archaeologists can sometimes tell whether ancient ochre was heated by looking at its crystals.

Good ochre was valuable and was traded across the continent. Red ochre from Pukardu (near Parachilna) in Adnyamathanha Country in the Flinders Ranges, and from Wilgie Mia on Wajarri Yamaji Country in Western Australia, one of the oldest mines in the world, was carried hundreds of kilometres. Many ochre sources are sacred sites, and who may collect and use ochre is governed by cultural law.

A person standing beside a cliff of yellow, cream and red ochre layers — The Ochre Pits in Tjoritja (West MacDonnell Ranges), on Western Arrernte Country, are still an important ochre source for Western Arrernte people. Visitors must not remove any ochre.Photo: Ian Cochrane, CC BY 2.0

Mistakes we see in tests

"Energy is stored in bonds and released when they break." Breaking a bond always takes energy. Energy is released when new, stronger bonds form.
"Exothermic reactions have a positive ΔH because energy comes out." ΔH describes the chemicals, which lose energy, so it is negative.
"Ionic bonds are weak and release little energy when they form." Some older books say this. Ionic bonds are strong; making sodium chloride from sodium and chlorine releases about 411 kJ per mole, which is why the reaction is so violent.
"A catalyst makes a reaction release more energy." A catalyst only lowers the activation energy so the reaction goes faster. ΔH stays the same.
"Exothermic reactions don't need any energy to start." Nearly all do: the activation energy. A match, a spark, friction or warmth provides it.

Test yourself

When a chemical bond forms, energy is
A reaction takes in 1200 kJ breaking bonds and gives out 1500 kJ making bonds. ΔH is
On an energy profile, the activation energy is measured from
Adding a catalyst to a reaction
Which reaction is endothermic?

Label the energy profiles.

Is each change exothermic or endothermic?

Breaking chemical bonds energy. The energy change of a reaction is called the enthalpy change, symbol . For an exothermic reaction ΔH is . The minimum energy needed to start a reaction is the energy. Heating a compound to break it down is called thermal .

5 marks

Hydrogen burns in oxygen: 2H₂ + O₂ → 2H₂O. Using bond energies H–H 436 kJ/mol, O=O 498 kJ/mol and O–H 463 kJ/mol, calculate ΔH for this reaction. State whether it is exothermic or endothermic and explain your answer in terms of bonds.

Your answer