Thermochemistry | General Chemistry 2

Thermochemistry: study of the energy change associated with chemical reactions.

System: object being investigated (= reaction mixture in thermochemistry)
Surroundings: everything else
Surroundings can exchange energy and matter if the system is open or only energy if the system is closed. There is no exchange with an isolated system.

Process: path of a system from an initial state to a final state
The final state is generally an equilibrium ⇒ all the properties of the system do not change with time

Isobaric = constant pressure
Isochoric = constant volume
Isothermal process = constant temperature
Adiabatic = no matter or heat transfer from the system to its surroundings

Heat q (in J): 

Transfer of energy based on a temperature difference (from high temperature region to low temperature region)
q < 0: reaction releases energy as heat ⇒ exothermic process
q > 0: reaction absorbs heat energy ⇒ endothermic process

Work w (in J):

Force x distance
w < 0: work is done by the system ⇒ expansion
w > 0: work is done on the system ⇒ compression

Gases can do work:

Internal energy U (in J):

Sum of kinetic and potential energies of the particles of a system:
U_sys = E_k + E_p

The energy change of a system ΔU = U_f - U_i
ΔU < 0: energy flows from the system into the surroundings
ΔU > 0: energy flows from the surroundings into the system

First Law of Thermodynamics:

The energy of the universe is constant ⇒ energy can neither be created nor destroyed

ΔU_univ = ΔU_sys + ΔU_surr = 0
⇒ ΔU_sys = - ΔU_surr = q + w

Enthalpy H (in J):

Energy transferred between a system and the surroundings under isobaric conditions
H_sys = U_sys + PV

H is a state function:
ΔH_rxn = H_prod – H_react 

The change in enthalpy for a reaction is:
ΔH_rxn = ΔU_rxn + Δ(PV)_rxn 
ΔH_rxn = ΔU_rxn + PΔV_rxn (P = cst)

ΔH < 0 ⇒ exothermic process
ΔH > 0 ⇒ endothermic process

Hess’s Law:

Enthalpy changes for chemical equations are additive

Consequences:

- reaction with 2 steps (or more): ΔH_rxn = ΔH_rxn (step 1) + ΔH_rxn (step 2)
- reversing a chemical equation: ΔH_rxn = ΔH_rxn (forward) + ΔH_rxn (reverse) = 0
⇒ ΔH_rxn (forward) = - ΔH_rxn (reverse)
- multiplication of a chemical equation A by a factor of n: ΔH_rxn = n x ΔH_rxn (A)

ΔH_sublimation = ΔH_fusion + ΔH_vaporization

Standard states: reactant and products in their pure forms at T = 298 K and P = 1 bar
(or at a concentration of 1 mol.L^-1 for a liquid solution)

Standard reaction enthalpy ΔH⁰_rxn (in J):
Enthalpy change when reactants in their standard states change to product in their standard states

Standard enthalpy of formation ΔH⁰_f (in J): 
Standard reaction enthalpy for the formation of the compound from its elements in their standard states. The standard enthalpy of formation of the most stable form of an element is 0.
 

ΔH⁰_f of H₂ (g) = 0 J
ΔH⁰_f of O₂ (g) = 0 J
ΔH⁰_f of C (graphite) = 0 J

ΔH⁰_rxn = Σ n ΔH⁰_f (products) - Σ n ΔH⁰_f (reactants)
with n = stoichiometric coefficients in the reaction
 

ΔH⁰_rxn of the combustion of propane:
C₃H₈ (g) + 5 O₂ (g) → 3 CO₂ (g) + 4 H₂O (l)
ΔH⁰_rxn = 3 ΔH⁰_f (CO₂) + 4 ΔH⁰_f (H₂O) - ΔH⁰_f (C₃H₈) - 5 ΔH⁰_f (O₂)

Bond enthalpy ΔH_bond (in J):

Energy associated with the strength of a chemical bond.
ΔH_bond > 0 because heat must be supplied to break a bond.

ΔH_rxn = Σ H_{bond broken} - Σ H_{bond formed} = Σ H_bond (reactants) - Σ H_bond (products)

Heat capacity c_P (in J.K^-1):

Amount of heat needed to raise the temperature of the sample by one degree.
At constant pressure:

Molar heat capacity C_P (in J.K^-1.mol^-1):

Amount of heat needed to raise the temperature of one mole of a sample by one degree.

Specific heat capacity c_S (in J.K^-1.g^-1):

Amount of heat needed to raise the temperature of one gram of a sample by one degree