Categories of Chemical Reactions
Senior 2 · Chemistry
Categories of Chemical Reactions
Lesson 3 of 4
22 min
Deep in Lake Kivu lie an estimated 65 cubic kilometres of dissolved methane. The KivuWatt project pumps gas-rich water up to a floating platform, separates the methane and pipes it ashore, where it is burned in engines to generate electricity; its first stage was designed to supply 26 megawatts to the national grid. The whole project depends on one fact: burning methane releases energy.
Other reactions do the opposite. A lime kiln must be heated for hours to turn limestone into quicklime, and a tea bush must absorb sunlight to make sugar. By the end of this lesson you will be able to classify reactions as exothermic or endothermic, interpret temperature changes, and explain energy changes in terms of bond breaking and bond making.
The system is the reacting chemicals. The surroundings are everything else: the water they are dissolved in, the test tube, the air, your hand. We detect energy changes by measuring the temperature of the surroundings.
An exothermic reaction transfers heat energy from the system to the surroundings, so the surroundings get hotter. The energy change, written ΔH ("delta H"), is negative because the chemicals have lost energy.
An endothermic reaction takes in heat energy from the surroundings, so the surroundings get colder (or the reaction needs continuous heating). ΔH is positive.
Ishimwe does two experiments. (a) He mixes 50 cm³ of dilute hydrochloric acid with 50 cm³ of dilute sodium hydroxide; the temperature rises from 22.0 °C to 28.5 °C. (b) He dissolves ammonium nitrate in water; the temperature falls from 23.0 °C to 14.0 °C. Classify each change.
Solution. (a) The temperature rises by 28.5 − 22.0 = 6.5 °C. Heat was released into the solution: exothermic, ΔH negative. (b) The temperature falls by 23.0 − 14.0 = 9.0 °C. Heat was taken from the water: endothermic, ΔH positive.
During a reaction, bonds in the reactants break and new bonds form in the products.
If more energy is released by making bonds than is absorbed by breaking them, the reaction is exothermic overall. If less is released, it is endothermic.
Use these bond energies to find ΔH for H2(g) + Cl2(g) → 2HCl(g): H–H 436 kJ/mol, Cl–Cl 243 kJ/mol, H–Cl 432 kJ/mol.
Solution.
ΔH is negative, so the reaction is exothermic.
An energy level diagram shows the energy of the reactants and products as two horizontal lines. For an exothermic reaction the products line is below the reactants line (energy has been given out); for an endothermic reaction the products line is above it. Even an exothermic reaction usually needs a small push to start: the activation energy, the minimum energy needed for a reaction to begin. That is why charcoal needs a lit match before it starts to burn, and then keeps itself going.
Burning 1 mole of methane releases 890 kJ. How much energy is released by burning 3 moles? And by 8 g of methane (CH4, molar mass 16 g/mol)?
Solution. 3 × 890 = 2670 kJ. For 8 g: 8 ÷ 16 = 0.5 mol, and 0.5 × 890 = 445 kJ. (You will learn more about moles in the Mole Concept unit.)
Tip
EXO = EXit, ENDO = ENter. In an exothermic reaction, heat exits into the surroundings, which get hot. In an endothermic reaction, heat enters the chemicals, and the surroundings get cold.
Signs: exothermic ΔH is negative (−); endothermic ΔH is positive (+). The sign is from the point of view of the chemicals, not the thermometer.
Try each item before you read the solution.
Classify each as exothermic or endothermic: (a) burning biogas, (b) photosynthesis, (c) respiration, (d) heating limestone in a kiln, (e) dissolving ammonium nitrate in water.
Solution: (a) exothermic, (b) endothermic, (c) exothermic, (d) endothermic, (e) endothermic.
The decomposition of calcium carbonate has ΔH = +178 kJ per mole. How much energy is needed to decompose 2 moles?
Solution: 2 × 178 = 356 kJ must be supplied, because the reaction is endothermic.
Use the bond energies H–H 436, O=O 498 and O–H 464 kJ/mol to find ΔH for 2H2(g) + O2(g) → 2H2O(g).
Solution: Bonds broken: 2 H–H + 1 O=O = 2 × 436 + 498 = 1370 kJ absorbed. Bonds made: each water molecule has 2 O–H bonds, so 4 O–H = 4 × 464 = 1856 kJ released. ΔH = 1370 − 1856 = −486 kJ: exothermic.
Charcoal burns in a highly exothermic reaction. Why does it not catch fire by itself in the store?
Solution: Every reaction needs its activation energy before it can start. At room temperature the charcoal and oxygen do not have enough energy; a match supplies it, and after that the heat released keeps the reaction going.
Summary
Sources and credits
This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.
