Transcript 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