Chapter Six—Principles of Reactivity:
Energy and Chemical Reactions
VOCABULARY--Define and review the following terms prior to our lecture:
 Thermodynamics
 Energy
 Potential Energy
 First Law of Thermodynamics
 Heat
 Temperature
 System
 Surroundings
 Endothermic
 Exothermic
 Heat Capacity
 Specific Heat Capacity
 Internal Energy
 Work
 Enthalpy
 State Function
Calorimetry:
E = q + w
w = - P V
E = internal energy
H = E at constant V
E = q - PV
q = H at constant P
Temperature Units: Kelvin (K) is the SI unit for temperature
o
K = 273.15 + oC
C = (5/9) (oF – 32)
Heat Units:
Joule (J) is the SI unit for thermal energy.
1 calorie = 4.184 J
1 J = 1 kg m2 / s2
1 Cal = 1000 cal = 1 kcal
1000 J = 1 kJ
Example 1: How much energy is needed to heat a 10.0 g piece of copper metal
from 25oC to 325oC? The specific heat of copper is 0.385 J/goC.
Example 2: A 55.0 g piece of metal is heated in boiling water to 99.8oC and then
dropped into a calorimeter of 225 mL of water at 21.0oC. The final
temperature of the water and metal is 23.1oC. What is the specific
heat of the metal?
Example 3: You mix 200. mL of 0.400 M HCl with 200. mL of 0.400 M NaOH in a
coffee-cup calorimeter. The temperature of the solutions before
mixing was 25.10oC; after mixing the temperature is 27.78oC. What
is the molar enthalpy of neutralization of the acid?
(let the densities of all solutions be 1.00 g/mL and their specific
heat capacities be 4.20 J/goC)
Example 4: A 1.00 g sample of sucrose (C12H22O11) is burned in a bomb
calorimeter. The temperature of the 1500. g of water in the
calorimeter rises from 25.00oC to 27.32oC. The heat capacity of the
bomb is 837 J/K.
a) calculate the heat evolved per gram of sucrose.
b) calculate the heat of combustion for sucrose.
HEATING CURVES:
Changes of State (Phase Changes)
vaporization
condensation
sublimation
deposition
solidification
fusion
Example 5: How much heat is needed to convert 500.0 g of ice at –25oC to steam
at 200oC?
Specific heat of ice = 2.06 J/g K
Specific heat of water = 4.184 J/g K
Specific heat of steam = 1.92 J/g K
Heat of fusion of ice = 333 J/g
Heat of vaporization of water = 2256 J/g
Energy Diagrams:
H
rxn
=
ENDOTHERMIC
H =
EXOTHERMIC
H =
Standard Enthalpy of Formation:
Standard State:
Example 6:
a) Using standard heats of formation, calculate the molar
enthalpy change for the combustion of methane, CH4.
b) How much energy would be released if you combusted 200.0 g
of CH4?
Example 7: What quantity of energy must be supplied to sublime 10.0 g I2?
I2 (s)  I2 (g)
H = 62.4 kJ
Hess’ Law of Heat Summation: If a reaction can be written as the _____ of 2
or more other reactions, the ____ for the
overall reaction is the sum of the _____ of
those other reactions.
Example 8: Use the following to calculate the standard heat of formation of
methane:
C (s) + O2 (g)  CO2 (g)
H = -393.5 kJ
H2 (g) + ½ O2 (g)  H2O (l)
H = -285.8 kJ
CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (l)
H = -890.3 kJ
Example 9: Calculate the standard heat of formation of B2H6 (g):
4B (s) + 3O2 (g)  2B2O3 (s)
H = -2543.8 kJ
2H2 (g) + O2 (g)  2H2O (g)
H = -484 kJ
B2H6 (g) + 3O2 (g)  B2O3 (s) + 3H2O (g)
H = -2032.9 kJ
HW Chapter 6: 14,22,26,30,32,34,38,42,44,60,74,92