Thermochemistry
Energy-capacity
to do work or to produce heat
Law of Conservation of Energy
–1st law of thermodynamics
–Energy can be neither created nor destroyed
–Energy of universe is constant
Potential energy
–in chemistry, this is usually the energy stored in reactants
and products
–Stored energy; sometimes because of position PE = mgh
Kinetic energy
–Energy involving movement
Energy

System
Reactants and products of a reaction

Surroundings
Anything other than reactants and products

Temperature
Measure of motion of particles (average kinetic
energy)
Exothermic reaction
–
–
–
–
–
∆H is negative
energy flows out of the system (heat) (-q )
temp↑
work done by system (on the surroundings ( -w )
ex. compression of gas
reactants
E
products
Reaction proceeds 
Endothermic reaction
–
–
–
–
–
–
∆H is positive
reaction that absorbs energy
energy flow into a system (heat) ( +q )
temp↓
work done by surroundings( + w )
ex. expansion of gas
products
E
reactants
Reaction proceeds 
Exothermic Process
Potential energy
CH4 + 2O2  CO2 + 2H2O + energy
2 mol O2
1 mol CH4
(Reactants)
System
Energy released
to the
surroundings as
heat
2 mol H2O
1 mol CO2
(Products)
Endothermic Process
N2 + O2 + energy (heat)  2NO
Potential Energy
System
1 mol N2
1 mol O2
(Reactants)
2 mol NO
(Product)
Heat absorbed
from the
surroundings
∆E = q + w
Example 1: Calculate the ∆E for a system
undergoing an endothermic process in
which 15.6 kJ of heat flows and where 1.4
kJ of work is done on the system.
∆E = + 15.6 kJ + + 1.4 kJ
= 17.0 kJ
q is positive b/c endothermic
w is positive b/c work is done on the system
Example 2:
q
= + 51 kJ and w = - 15 kJ
∆E = 36 kJ
q
∆E
= +47 kJ and w = 88 kJ
= 41 kJ
Which of the above has work done on
surroundings?
-15 kJ
Enthalpy H
Heat content at constant pressure
∆H = Hproducts - Hreactants
Enthalpy of reaction ∆Hrxn
heat absorbed or released by a chemical reaction
Enthalpy of combustion ∆Hcomb
heat absorbed or released when burning
Enthalpy of formation ∆Hf
heat absorbed or released when ONE mole of compound is
formed from elements in their standard states
Enthalpy of fusion ∆Hfus
heat absorbed to melt 1 mole of solid to liquid @ MP
Enthalpy of vaporization ∆Hvap
heat absorbed to change change 1 mole of liquid to gas @ BP
Endothermic reactions
Exothermic reactions
Enthalpy can be calculated:
Stoichiometrically
From tables
Using Hess’ Law
+H
-H
S(s) + O2 (g)  SO2 (g) ∆H = -296 kJ/mol
a) calculate the heat evolved when 275 g S is burned
275 g
1mol S
32.1 g S
296 kJ
1 mol
=
-2535.8 kJ
= 2.54 X 10-3 kJ
b) calculate the heat evolved when 150 g SO2 is produced
150 g SO2 1 mol SO2 1 mol S
64.1 g SO2 1 mol SO2
296 kJ = 693 kJ
1 mol S
When 1 mole of methane (CH4) is burned at
constant pressure, 890 kJ of energy is released
as heat. Calculate H for a process in which a
5.8 g sample of methane is burned at constant
pressure.
At constant pressure, q = H = -890 kJ/mol
Calorimetry
Science of measuring heat
Calorimeter constant pressure (coffee cup) or
constant volume (bomb)
Heat capacity
C = heat absorbed
increase in temperature
 Specific heat capacity- energy required to raise
the temperature of one gram of a substance by
one degree Celsius
q=mcT

Molar heat capacity-energy required to raise the
temperature of one mole of a substance by one
degree Celsius
Energy released by a reaction =
energy absorbed by the solution

If two reactants at the same temperature are
mixed and the resulting solution gets warmer,
this means the reaction taking place is
exothermic
 H = (-)

If the solution gets cooler, the reaction is
endothermic
 H = (+)
Consider the dissolution of CaCl2:
CaCl2(s)  Ca2+(aq) + 2Cl-(aq) H = -81.5 kJ
An 11.0 g sample of CaCl2 is dissolved in 125 g of
water with both substances at 25.0°C. Calculate
the final temperature of the solution assuming
no heat lost to the surroundings and assuming
the solution has a specific heat capacity of
4.18J/goC

A 46.2 g sample of copper is heated to
95.4 °C and then placed in a calorimeter
containing 75.0 g water at 19.6 °C. The
final temperature of the metal and water
is 21.8°C. Calculate the specific heat
capacity of copper, assuming that all the
heat lost by the copper is gained by the
water.
Camphor (C10H16O) has an energy of
combustion of -5903.6 kJ/mol. When a
sample of camphor with mass 0.1204 g is
burned in a bomb calorimeter, the
temperature increases by 2.28°C.
Calculate the heat capacity of the
calorimeter.