Download Thermochem

Thermochemistry
“The Quick and Dirty”
 Energy changes accompany every chemical and
physical change.
 In chemistry heat energy is the form of energy that we
are most often interested in.
 Kinetic energy (energy of motion)
 Potential energy (stored energy)
 Chemical bond energy is the major form of potential
energy we are concerned about in chemistry.
 Heat transfer is always from the warmer object to the
colder object.
 The standard unit of heat energy is joule (J).
 Kelvin to 0C + 273
 Open/closed systems: Closed systems can exchange
energy but not matter.
 Law of conservation of Energy = 1st Law of
Thermodynamics
Phase Change Graph For Water
 Exothermic reaction – molar enthalpy (ΔH) is lost by
conversion to heat or light.
- energy is lost to the
surroundings
- energy of the system decreases
 Endothermic reaction – energy in the surroundings is
absorbed and converted to molar enthalpy.
 Calorimeter – thermally insulated container in which
the exchange between the system and its surroundings
can be measured
 q = cmΔT
 m= q
cΔT
ΔT = q
cm
 Energy changes during a state change:
q = nΔH phase
n- number of moles
Hf – heat of fusion
Hv – heat of vapourization
 Exothermic Reactions - net release of energy
 Energy term is on the product side
 Fe2O3 + 2Al  2Al2O3 + 2Fe + 847.6 KJ
 Endothermic Reactions – net input of energy
 Energy term is on the reactants side.
 2 SO3 + 198 KJ  2 SO2 + O2
 Chemical Reactions occur spontaneously for two
reasons:
1. The products of the reaction have less energy than the
reactants (burning a match) always exothermic.
2. Products are more random than the reactants. This is
the entropy (S).
Changes that produce substances with greater randomness
(+ΔS) are favoured in nature and drive the reaction to
occur.
More Randomness:
1. Solid state  liquid
2. Liquid state  gas
3. Solid  gas
4. Formation of a mixture
5. Increase in volume of a gas

Gas
Highest S
Aqueous
Liquid
Solid
Lowest S
Enthalpy:
Δ H = H final – H initial OR
Δ H = H products – H reactants
Exothermic:
-ΔH
Endothermic:
+ΔH
The reactants have less potential energy than do the products.
Energy must be input in order to raise the particles up to the higher
energy level.
Energy + A + B --> AB
The reactants have more potential energy than the products have.
The extra energy is released to the surroundings.
A + B --> AB + Energy
 Writing Equations:
1. ΔH notation for 1 mol of CO
exothermic
Fe2O3(s) + 3CO(g)  3CO2(g) + 2 Fe (s) + 25 kJ
1/3 Fe2O3(s) + CO(g)  CO2 (g) + 2/3 Fe(s)
2. Using Energy as a term:
3FeCl3(s)  3FeCl2(s) + 3/2 Cl2
ΔH = -8.3 kJ
endothermic
ΔH = + 173 kJ
6 FeCl3(s) + 346 kJ  6 FeCl2(s) + 3Cl2(g)
 Calculating Heat:
 How much heat is produced when 95 g of methane is
burned in oxygen?
CH4 (g) + 2 O2(g)  CO2(g) + 2H2O(g)
95g x 1mol = 5.9 mol
16g
ΔH = 891 kJ/mol
1 mol so 891 kJ
q= nΔH
= 5.9 mol x 891 kJ/mol
= 5300 kJ
ΔH = -891kJ
 Calorimetry

ΔH substance = mcΔT
n
m -mass of water (usually taking in heat)
C -specific heat of water 4.18 J/goC
ΔT – change in temperature
n – moles of the substance that you are calculating the ΔH
of
 Hess’s Law
 Based on 1 mole.
 Reaction occurs in a series of steps in which the
intermediates are cancelled.
OR
ΔH reaction = ΣHf products –ΣHf products
Standard Heat of Formations from elements( you will need to
write a balanced equation for the formation of the
substance).
Bond Energies – uses Lewis Structures to draw structural
formulas (energy values come from table)
ΔH = Σ Reactants – Σproducts
ΔG:
- ΔG reaction is spontaneous