Thermodynamics: Enthalpy Change and Calorimetry

Energy and Units

• 1 calorie (cal) is the energy required to raise the temperature of 1g of water by 1°C
• 1 kcal = 4.184 kJ
• 1 kJ = 1000 J = 0.2390 kcal = 239 cal

Internal Energy and Enthalpy

• Internal energy (E) is the total energy of a system
• Enthalpy (H) is the internal energy plus the product of pressure and volume: H = E + PV
• Change in enthalpy (∆H) is the energy plus the product of constant pressure and change in volume: ∆H = ∆E + P∆V

Comparing ∆E and ∆H

• For reactions that do not involve gases, ∆H ≈ ∆E
• For reactions where the amount of gas does not change, ∆H = ∆E
• For reactions where the amount of gas does change, ∆H is very close to ∆E

Exothermic and Endothermic Processes

• Exothermic process: releases heat and results in a decrease in the enthalpy of the system (∆H < 0)
• Endothermic process: absorbs heat and results in an increase in the enthalpy of the system (∆H > 0)

Enthalpy Changes

• Heat of formation (∆Hf): when 1 mol of a compound is produced from its elements
• Heat of fusion (∆Hfus): when 1 mol of a substance melts
• Heat of vaporization (∆Hvap): when 1 mol of a substance vaporizes

Calorimetry

• Heat (q) absorbed by an object is proportional to its temperature change: q = constant × ∆T
• Every object has its own heat capacity: Heat Capacity = q/∆T (J/K)
• Specific heat capacity (S) is the quantity of heat required to change the temperature of 1 gram of a substance by 1K: S = q/mass × ∆T (J/g.K)

Calculating Heat Absorbed or Released

• q = S × mass × ∆T
• Molar heat capacity: quantity of heat required to change the temperature of 1 mole of a substance by 1K: Molar Heat Capacity = q/moles × ∆T (J/mol.K)

Types of Calorimetry

• Constant-pressure calorimetry: measures heat transfer in a process open to the atmosphere
• Constant-volume calorimetry: measures heat transfer in a closed system, such as a bomb calorimeter