Energetics
Exothermic and Endothermic Reactions
Exothermic reactions release heat energy to the surroundings. The temperature of the surroundings increases during an exothermic reaction.
Examples of exothermic reactions:
- Combustion
- Neutralisation
- Many displacement reactions (e.g. iron + copper sulfate solution)
- Some salts dissolving in water
Endothermic reactions take in heat energy from the surroundings. The temperature of the surroundings decreases during an endothermic reaction.
Examples of endothermic reactions:
- Thermal decomposition (e.g. heating copper carbonate)
- Reaction of citric acid with sodium hydrogencarbonate
- Some salts dissolving in water (e.g. ammonium nitrate)
How to identify from data: If the final temperature is higher than the initial temperature, the reaction is exothermic. If the final temperature is lower, the reaction is endothermic.
Calorimetry
Calorimetry is the experimental measurement of heat energy changes in reactions.
Reactions in Solution (Neutralisation, Displacement, Dissolving)
Apparatus: Polystyrene cup (acts as an insulated calorimeter), thermometer, stirrer.
Method:
- Measure a fixed volume of one reagent into the polystyrene cup and record the initial temperature.
- Add the second reagent and stir continuously.
- Record the maximum (or minimum) temperature reached.
- Calculate the temperature change (DT).
Why use a polystyrene cup? Polystyrene is a poor conductor of heat (an insulator), so less heat is lost to the surroundings, making the measurement more accurate.
Combustion Calorimetry
Method:
- Measure a fixed volume of water into a copper calorimeter (metal can) and record the initial temperature.
- Weigh a spirit burner containing the fuel.
- Burn the fuel beneath the calorimeter and stir the water.
- After a suitable temperature rise (approximately 20 degrees Celsius), extinguish the flame.
- Record the final temperature and re-weigh the spirit burner.
- Calculate the mass of fuel burned and the temperature change.
Sources of error in combustion calorimetry:
- Heat lost to the surroundings (not all heat transferred to water)
- Heat absorbed by the calorimeter itself
- Incomplete combustion of the fuel
To reduce errors: use a lid on the calorimeter, shield from draughts, keep the calorimeter close to the flame.
Calculating Heat Energy Change
The heat energy change, Q, is calculated using:
Q = m x c x DT
Where:
- Q = heat energy change (joules, J)
- m = mass of the substance being heated (grams, g); for solutions, assume the density is 1 g/cm3, so mass = volume in cm3
- c = specific heat capacity of water = 4.18 J/g/degrees Celsius (sometimes given as 4.2 J/g/degrees Celsius)
- DT = temperature change (degrees Celsius)
Worked Example: 100 cm3 of sodium hydroxide solution and 100 cm3 of hydrochloric acid are mixed in a polystyrene cup. The temperature rises from 19.0 degrees Celsius to 26.5 degrees Celsius.
Mass of solution = 200 g (assuming density = 1 g/cm3) DT = 26.5 - 19.0 = 7.5 degrees Celsius Q = 200 x 4.18 x 7.5 = 6270 J = 6.27 kJ
Calculating Molar Enthalpy Change
The molar enthalpy change (DH) is the heat energy change per mole of reactant:
DH = Q / n
Where:
- DH = molar enthalpy change (kJ/mol)
- Q = heat energy change (convert to kJ by dividing by 1000)
- n = number of moles of the substance
For exothermic reactions, DH is negative (energy is released). For endothermic reactions, DH is positive (energy is absorbed).
Worked Example: Using the example above, if 0.1 mol of HCl was used: Q = 6.27 kJ DH = -6.27 / 0.1 = -62.7 kJ/mol (negative because exothermic)
Energy Level Diagrams
Higher Tier
Spec point 3.5C: draw and explain energy level diagrams for exothermic and endothermic reactions.
Energy level diagrams show the relative energies of reactants and products. The y-axis shows energy content; the x-axis shows the progress of the reaction.
Exothermic reaction:
- Reactants have higher energy than products
- DH is negative
- The arrow on the diagram points downward (from reactants to products)
Endothermic reaction:
- Products have higher energy than reactants
- DH is positive
- The arrow on the diagram points upward (from reactants to products)
Bond Breaking and Bond Making
Higher Tier
Spec points 3.6C and 3.7C.
During chemical reactions:
- Bond breaking requires energy input. It is an endothermic process.
- Bond making releases energy. It is an exothermic process.
Whether a reaction is overall exothermic or endothermic depends on the balance between energy taken in (to break bonds) and energy released (when new bonds form):
- If energy released (bond making) > energy taken in (bond breaking): the reaction is exothermic (DH is negative)
- If energy released (bond making) < energy taken in (bond breaking): the reaction is endothermic (DH is positive)
Enthalpy change from bond energies:
DH = total energy taken in (to break bonds in reactants) - total energy released (when bonds form in products)
Worked Example: H-H + Cl-Cl → 2 H-Cl
Bond energies: H-H = 436 kJ/mol; Cl-Cl = 242 kJ/mol; H-Cl = 431 kJ/mol
Energy in (bonds broken in reactants) = 436 + 242 = 678 kJ Energy out (bonds formed in products) = 2 x 431 = 862 kJ DH = 678 - 862 = -184 kJ/mol
The negative sign confirms the reaction is exothermic.
Exam Tip
Always check the balancing numbers in the equation when counting bonds. For example, “2 H-Cl” means two H-Cl bonds are formed.
Practical: Investigating Temperature Changes
Aim (spec 3.8): Investigate temperature changes in:
- Salts dissolving in water
- Neutralisation reactions
- Displacement reactions
- Combustion reactions
Method for neutralisation (HCl + NaOH):
- Measure 25 cm3 of NaOH solution into a polystyrene cup and record the temperature.
- Add 5 cm3 of dilute HCl and stir. Record the highest temperature reached.
- Repeat, adding further 5 cm3 portions of HCl (up to a total of about 40 cm3), recording the temperature after each addition.
- Plot temperature against total volume of acid added.
- Draw two lines of best fit; they cross at the point of exact neutralisation (where the temperature is maximum).
Exam Questions
Q: What is the difference between an exothermic and an endothermic reaction?
In an exothermic reaction, heat energy is released to the surroundings, causing the temperature to increase. In an endothermic reaction, heat energy is taken in from the surroundings, causing the temperature to decrease.
Q: In a calorimetry experiment, 50 cm3 of copper sulfate solution was mixed with excess iron powder. The temperature rose from 20.0 degrees Celsius to 28.6 degrees Celsius. Calculate the heat energy released. (c = 4.18 J/g/degrees Celsius)
Mass of solution = 50 g; DT = 28.6 - 20.0 = 8.6 degrees Celsius; Q = 50 x 4.18 x 8.6 = 1797 J (approximately 1.80 kJ). The reaction is exothermic (temperature increased).
Q: (Higher) Explain why bond breaking is endothermic and bond making is exothermic.
Energy must be supplied to overcome the electrostatic attraction between bonded atoms, so bond breaking takes in energy and is endothermic. When bonds form, energy is released as the atoms become more stable at a lower energy level, so bond making is exothermic.
Q: (Higher) Calculate the enthalpy change for: N2(g) + 3H2(g) → 2NH3(g). Bond energies: N=N 945 kJ/mol; H-H 436 kJ/mol; N-H 391 kJ/mol.
Energy in: 1 x 945 (N=N) + 3 x 436 (H-H) = 945 + 1308 = 2253 kJ. Energy out: 2 x 3 x 391 (6 N-H bonds) = 2346 kJ. DH = 2253 - 2346 = -93 kJ/mol. The reaction is exothermic.
Exam Tip
When asked for the enthalpy change, always give the sign: negative for exothermic, positive for endothermic. When using Q = mcDT, make sure the mass is in grams and the temperature change is in degrees Celsius. Convert Q from joules to kilojoules (divide by 1000) before calculating DH in kJ/mol.