Bond energies and calorimetry
Bond energies and calorimetry
- During a reaction, old bonds break and new bonds form.
- Breaking takes energy in; making gives energy out.
- We can also measure $\Delta H$ in the lab.
Practice
Which statement is correct?
Energy is needed to break bonds and released when bonds form; ΔH is the difference.
Bond energy calculations
- Breaking a bond is endothermic; making a bond is exothermic.
$$\Delta H_r = \sum(\text{bonds broken}) - \sum(\text{bonds made})$$
- The bond energy is the energy to break one mole of a bond (gas state) — always positive.
- Some values are exact; others are averages over many molecules, so the answer is approximate.
Practice
Using bond energies, ΔH of reaction equals:
ΔH = Σ(bonds broken) − Σ(bonds made); energy in to break minus energy out to make.
Practice
Why are calculations using average bond energies only approximate?
Average bond energies are means over many molecules, so they don't match any single molecule exactly.
Measuring in the lab
- When a reaction heats a known mass of solution, the heat transferred is:
$$q = mc\Delta T$$
- ($m$ = mass, $c$ = specific heat capacity, $\Delta T$ = temperature change.) Then per mole:
$$\Delta H = -\frac{mc\Delta T}{n}$$
- The minus sign makes $\Delta H$ negative when the temperature rises (exothermic).
Practice
The heat transferred to a solution is calculated with:
q = mcΔT (mass × specific heat capacity × temperature change); ΔH per mole = −mcΔT/n.
You've got it
Key idea
- breaking bonds is endothermic; making bonds is exothermic
- $\Delta H_r = \sum(\text{bonds broken}) - \sum(\text{bonds made})$
- average bond energies give only an approximate answer
- calorimetry: $q = mc\Delta T$, then $\Delta H = -\dfrac{mc\Delta T}{n}$