Exothermic Reactions:
- Reactants → Products + Heat energy
- The temperature of the surroundings increases because chemicals are converted to heat energy; the surroundings get warmer.
- Examples: Reactions of metals with acids, neutralisation of an acid with an alkali, combustion of fuels and respiration in body cells.
- In an exothermic reaction, the reactants have more energy than the products. The amount of heat energy released = ΔH.
- Exothermic reaction = -ΔH because reactants use energy in the form of heat.
- An exothermic reaction can be shown on an energy diagram:
Endothermic Reactions:
- Reactants + Heat energy → Products
- The temperature of the surroundings decreases because chemicals gain stored chemical potential energy from the surroundings, which is converted to CPE.
- Examples: Photosynthesis, thermal decomposition of CaCO3 (s), and cooking (e.g., frying an egg or baking a cake).
- In an endothermic reaction, reactants have less energy than the products. ΔH is positive (+ΔH) because the reactants are gaining energy in the form of heat.
With a catalyst, the Ea bump/height would be lower for exo and endo.
Exothermic and Endothermic Reactions:
Exothermic: A process that gives out heat energy (-ΔH).
Endothermic: A Process that takes in heat energy (+ΔH).
Calorimetry:
- The experimental process of measuring the thermal energy change in a chemical/physical change.
- A calorimeter is a device used to measure thermal energy changes in a chemical/physical change.
1) For combustion of fuels 2) For reactions involving solutions:
- To simplify calculations, we assume:
- The calorimeter transfers negligible thermal energy to the outside environment.
- The calorimeter itself absorbs negligible thermal energy.
- All dilute aq solutions have the same density and specific heat capacity.
ΔH = molar enthalpy change (kJ/mol)
→To find the change Q kJ and divide by moles.
Worked Examples:
1) 25cm3 of 0.5 mol/dm3 NaOH mixed with 25cm3 0.5 mol/dm3 HCl. The temperature increased from 25oC to 28oC. Find ΔH.
a) Balanced symbol eq. = NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)
b) Q = m c ΔT
→ Q = (25 + 25) × 4.2 × (28 - 25) → 50 × 4.2 × 3 = 630 J
c) Convert to kJ→ 0.63kJ
d) Find moles of limiting reactant → Here, there's no limiting reactant.
e) 
f) Exo or endo? → is exothermic, so add a sign! = 50.4 kJ/mol
2. When 25cm3 of 0.2 mol/dm3 silver nitrate was mixed with 0.1 g of zinc powder, the temperature rose by 8.6oC. Calculate ΔH.
2AgNO3 (aq) + Zn(s) → Zn(NO3)2 (aq) + 2Ag (s)
Q = 4.2 x 25 x 8.6 = 903 J → 0.903 kJ
Enthalpy of Combustion:
- Burning a fuel will always be an exothermic reaction.
- We can measure how much energy is released when fuel is burnt per mole or gram.
→ It allows us to compare different fuels.
- Measuring heat transfers of fuel = fuel calorimetry.
Ref to diagram:
- The metal used is often made of copper, as it's a powerful conductor of heat.
- Measurements taken:
→ Thermometer readings, before and after.
→ Amount of alcohol, before and after.
→ Amount of water in the calorimeter.
- You can find out how much energy water has gained by Q = mc ΔT, and then the molar enthalpy of the ethanol using ΔH.
- The calculated value may differ from the data book value for the following reasons:
→ Heat loss to surroundings.
→ We assume complete combustion occurs.
- When comparing different fuels, the type of can, the amount of fuel and water, and the tripod's height should be the same.
Measuring Heat Energy from a Fuel:
Ref to diagram
① Set up equipment and place the calorimeter so the base is ≈ 8 cm above the spirit lamp.
② Use a measuring cylinder to transfer 100cm3 water to the calorimeter.
③ Weigh the spirit burner with alcohol and cap over the burner (prevents evaporation) and record mass.
④ Place the spirit burner under the calorimeter, stir the water in the calorimeter, and record the initial temperature.
⑤ Remove the spirit burner cap and light the wick.
⑥ Keep stirring water and when the temperature has risen by 35oC record the final temp put the cap back on to extinguish the flame.
⑦ Reweigh the spirit burner with the lid on to find the mass of fuel burnt (before - after).
⑧ Calculate ΔH for fuel (i.e., moles for fuel and mass of fuel, not water).
- The main source of error = heat lost to surroundings → ΔT becomes lower and ΔH is less exothermic/negative.
- Improve by adding a lid on the calorimeter and insulating the sides to prevent heat loss.
- Soot at the bottom of the can could mean not as much heat would get to the can, leading to a lower ΔT and a less negative ΔH.
Bond Breaking and Making:
- Chemical processes have two steps: bond breaking and making energy.
- Bond breaking (in reactants) → endothermic—takes energy.
- Bond making (in products) → exothermic—gives out energy.
- If the energy used to break bonds in reactants is greater than the energy released when new bonds are formed, then the overall process is endothermic.
- If the energy used to break bonds in reactants is less than the energy released when new bonds are made, then the overall process is exothermic.
Energy profile diagram for exothermic and endothermic reactions:
For an exothermic reaction, ΔH will be negative, and products will have less CPE stored than reactants.
For an endothermic reaction, ΔH will be positive, and the products will have more CPE stored than reactants.
Energy/enthalpy changes (ΔH) are defined in terms of mole quantities, so they have units of kJ/mol.
H stands for energy/enthalpy change, which leads to a difference in the stored CPE between the products and reactants.
Bond Energies and calculating ΔH:
- The value of ΔH can be estimated using mean molar bond energies for bonds broken and created.
- Mean bond energy = Average amount of energy to break 1 mol of a particular type of covalent bond.
E.g., Calculate ΔH for H2 + Cl2 2HCl
→ Can also unite out the displayed formula to help count the number of bonds.
| Bond | Bond Energy (kJ/mol) |
|---|---|
| H-H | 436 |
| CI-CI | 242 |
| H-CI | 431 |
Break bonds = 436 + 242 = 678 kJ/mol
Make bonds = 431 × 2 = 862 kJ/mol
Bonds broken - Bonds that were made = 678 - 862 = 184 kJ/mol
→ Exothermic means that bond formation releases more energy than bond breaking does.
- Calculated values often differ from the true value because bond energies are mean and aren't the same in every bond.
Reversible Reactions (Start of Equilibria):
- During a chemical reaction, reactants are turned into products. Certain reactions allow for the reversibility of products back into reactants.
1. Heating hydrated copper (II) sulfate
- This is an endothermic reaction because we externally provide the heat so it can thermally decompose and dehydrate.
→ However, the reverse process is exothermic: temperature↑.
- CuSO4. 5H2O is a blue solid before heating.
- During heating, it turns white and water vapour starts forming (prove using cobalt chloride paper).
→ CuSO4. 5H2O → CuSO4 + 5H2O
- After heating, it’s anhydrous copper sulfate as the blue solid loses its water of crystallisation and becomes a white solid.
- After leaving CuSO4 to cool, add a few drops of water, and it becomes warmer and blue as it gains its water of crystallisation.
→ CuSO4 + 5H2O → CuSO4. 5H2O
- Other hydrated salt crystals can also be de- and re-hydrated in a similar manner, e.g., CoCl2.H2O 
CoCl2 + 2H2O.