Thermochemistry | General Chemistry 2
Important Terms in Thermochemistry
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
Internal Energy, Heat and Work
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:
wgas = -
Pext = constant external pressure
ΔV =
Internal energy U (in J):
Sum of kinetic and potential energies of the particles of a system:
Usys = Ek + Ep
The energy change of a system ΔU = Uf - Ui
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
ΔUuniv = ΔUsys + ΔUsurr = 0
⇒ ΔUsys = - ΔUsurr = q + w
Enthalpy
Enthalpy H (in J):
Energy transferred between a system and the surroundings under isobaric conditions
Hsys = Usys + PV
H is a state function:
ΔHrxn = Hprod – Hreact
The change in enthalpy for a reaction is:
ΔHrxn = ΔUrxn + Δ(PV)rxn
ΔHrxn = ΔUrxn + PΔVrxn (P = cst)
ΔH < 0 ⇒ exothermic process
ΔH > 0 ⇒ endothermic process
Enthalpy Relationships
Hess’s Law:
Enthalpy changes for chemical equations are additive
Consequences:
- reaction with 2 steps (or more): ΔHrxn = ΔHrxn (step 1) + ΔHrxn (step 2)
- reversing a chemical equation: ΔHrxn = ΔHrxn (forward) + ΔHrxn (reverse) = 0
⇒ ΔHrxn (forward) = - ΔHrxn (reverse)
- multiplication of a chemical equation A by a factor of n: ΔHrxn = n x ΔHrxn (A)
ΔHsublimation = ΔHfusion + ΔHvaporization
Enthalpy of Formation
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 ΔH0rxn (in J):
Enthalpy change when reactants in their standard states change to product in their standard states
Standard enthalpy of formation ΔH0f (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.
ΔH0f of H2 (g) = 0 J
ΔH0f of O2 (g) = 0 J
ΔH0f of C (graphite) = 0 J
ΔH0rxn = Σ n ΔH0f (products) - Σ n ΔH0f (reactants)
with n = stoichiometric coefficients in the reaction
ΔH0rxn of the combustion of propane:
C3H8 (g) + 5 O2 (g) → 3 CO2 (g) + 4 H2O (l)
ΔH0rxn = 3 ΔH0f (CO2) + 4 ΔH0f (H2O) - ΔH0f (C3H8) - 5 ΔH0f (O2)
Bond Enthalpies
Bond enthalpy ΔHbond (in J):
Energy associated with the strength of a chemical bond.
ΔHbond > 0 because heat must be supplied to break a bond.
ΔHrxn = Σ Hbond broken - Σ Hbond formed = Σ Hbond (reactants) - Σ Hbond (products)
Heat Capacity
Heat capacity cP (in J.K-1):
Amount of heat needed to raise the temperature of the sample by one degree.
At constant pressure:
cP =
qP = energy added as heat (in J)
ΔT = Tf – Ti (in K)
Molar heat capacity CP (in J.K-1.mol-1):
Amount of heat needed to raise the temperature of one mole of a sample by one degree.
CP =
cP = heat capacity (in J.K-1)
n = number of mole (in mol)
Specific heat capacity cS (in J.K-1.g-1):
Amount of heat needed to raise the temperature of one gram of a sample by one degree
cS =
cP = heat capacity (in J.K-1)
m = mass (in g)