Terms
• Energy
• Heat
• Calorie
• Joule
• Specific heat
• Calorimeter
• Thermochemistry
Chapter 16
Thermochemistry
Energy
Energy is the capacity to do work, and can
take many forms
• Potential energy - stored energy or
energy of position
• Kinetic energy - energy of motion
• Thermal energy - (heat) movement at
the atomic level
Heat (Enthalpy) Change, ΔH
• The amount of heat energy released or
absorbed during a process.
Law of conservation of energy
•Energy can be converted
from one form to another, but
neither created nor
destroyed.
Heat vs. Temperature
• Heat – q, the energy that flows from
hot to cold;
• Temperature – measure of average
kinetic energy.
Calorimetry
Measures the amount of heat absorbed or released
during a physical/chemical change usually by the
change in temperature of a known quantity of water
in a calorimeter.
Measuring heat
• Metric system: calorie (cal)
– Heat required to raise the temperature of
1 gram of pure water 1ºC
• Food Calories differ from heat calories
– 1 Calorie = 1000 cal
• SI unit: joule (J)
– 1 J = 0.2390 cal
– 1 cal = 4.184 J
Converting Energy Units
• Use a conversion to convert nutritional
Calories to calories.
• Use a conversion factor to convert
calories to joules.
Practice
• Convert the following
–
–
–
–
–
–
934 Calories to calories
7.42x106 calories to Calories
236 Calories to joules
2.59x106 calories to joules
84 joules to calories
7.19x106 joules to Calories
Changes in State
Evaporation
Melting
Condensation
Freezing
Changes in State
• Heat of Fusion (Hf)
– Conversion from Solid
Liquid
• Heat of Vaporization (Hv)
– Conversion from Liquid
Gas
Formula q = m Hf or m Hv
Hf water = 80 cal/g or 334 J/g
Hv water = 540 cal/g or 2260 J/g
Change in State Graph
Change in State
250
Temperature
200
150
Temperature
100
50
0
0
10
20
30
40
Time ( Min )
1. Copy this graph
2. Identify specific heat, heat of fusion, and heat of
vaporization.
Change in State Areas
Sample Heat of Fusion
• How much heat energy is needed to melt
5.0 kilograms of water at its melting point?
Q = m Hf
= 5 kg x 334 J/kg
= 1,670J
Q = m Hf
= 5000 g x 80 cal/g
= 400,000 cal
Collision Model
• Collisions must have enough
energy to produce the reaction
(must equal or exceed the
activation energy).
• Reactants must have proper
orientation to allow the formation of
new bonds.
Activation Energy
• The minimum energy required to
transform reactants into the activated
complex
• Flame, spark, high temperature, radiation
are all sources of activation energy
Endo. Vs. Exo-thermic Rxn
• When the reaction releases heat it is
exothermic, feels hot. -ΔH
– Energy is given off
• When the reaction absorbs heat it is
endothermic, feels cold. +ΔH
– Energy is taken in
Exothermic Processes
Processes in which energy is released as it
proceeds, and surroundings become warmer
Reactants  Products + energy
Endothermic Processes
Processes in which energy is absorbed as it
proceeds, and surroundings become colder
Reactants + energy  Products
Endothermic Reaction w/Catalyst
Exothermic Reaction w/Catalyst
Exo vs. Endo
Exothermic
1. Which has more energy, reactants or products?
2. If the reaction ran backwards, which would have more energy,
reactants or products?
Endothermic
1. Which has more energy, reactants or products?
2. If the reaction ran backwards, which would have more energy,
reactants or products?
3. Looking at your two diagrams, which has larger activation
energy?
4. In general does an exothermic reaction have a larger activation
energy running forwards or backwards?
5. In general does an endothermic reaction have a larger activation
energy running forwards or backwards?
Practice: State whether each of the following
are exothermic or endothermic.
1. H + Cl  HCl + 432 kJ
2. 12 CO2 + 11 H2O  C12H22O11 + 12 O2
ΔH=5638 kJ
3. ice  water
4. C + D  CD ΔH=-65.8 kJ
5. E + F + 437 kJ  G + H
6. H2O vapor  H2O liquid
Practice: State whether each of the following
are exothermic or endothermic.
7.
8. KOH  K+ + OHΔH=-57.8 kJ
9. C3H8 + 5 O2  3 CO2 + 4 H2O
ΔH=-2221 kJ
10. Ca(OH)2  Ca + O2 + H2
ΔH=+986 kJ
11. Fe2O3 + 2 Al  Al2O3 + 2 Fe ΔH=-852 kJ
12. 2 H2O  2 H2 + O2
ΔH=+572 kJ
Specific Heat
• Know: amount of heat required to
raise the temperature of 1 gram by
1°C.
• That quantity is defined as the
specific heat (c).
• Each substance has its own specific
heat.
Calculating heat
• Heat absorbed/released by a
substance depends on the specific
heat, mass, and the temperature
change.
Calculating Heat
• q = the heat absorbed or released
• c = the specific heat of the substance
• m = the mass of the sample in grams
• ∆T is the change in temperature in °C.
• ∆T final temperature - the initial
temperature: Tfinal – Tinitial.
Converting
• 256 cal into Jolues
• 987 cal into Joules
4
• 5.9 x10 J into cal
• 3.7 x104 J into cal
Practice Problems: q = m c ∆T
1. How much heat is lost when a piece of
aluminum with a mass of 4110 g cools from
660oC to 25oC? (c for Al = 0.9025 J/goC)
2. If 7.5 x 104J of energy are released when 500
g of a metal cools from 100.0oC to 20.0oC, what
is the specific heat of the metal?
3. If the temperature of 34.4 g of ethanol
increases from 25.0ºC to 78.8 ºC how much heat
has been absorbed? (c=.00244 J/goC)
Specific Heat Practice Problems
4. A 4.50 g nugget of gold absorbed 276 J of heat.
What was the final temperature of the gold if the
initial temperature was 25.0 ºC and the specific
heat is 0.129 J/g • ºC
5. A 155 g sample of unknown substances was
heated from 25.0 ºC to 40.0 ºC and it absorbed
5696 J of energy. What is the specific heat of the
substance?
Create your own Graphs
Specific Heat
SUBSTANCE
SPECIFIC HEAT
SUBSTANCE
CAPACITY (J/KG ºC)
SPECIFIC HEAT
CAPACITY (J/KG ºC)
Aluminum
Brass
9.0 x 102
3.8 x 102
Alcohol (ethyl)
Alcohol (methyl)
2.3 x 102
2.5 x 102
Copper
Glass (crown)
3.9 x 102
6.7 x 102
Glycerine
Mercury
2.4 x 102
1.4 x 102
Glass (pyrex)
7.8 x 102
Nitrogen (liquid)
1.1 x 102
Gold
Iron
Lead
Sand
Silver
1.3 x 102
4.5 x 102
1.3 x 102
8.0 x 102
2.3 x 102
Water (liquid)
Water (ice)
Water (steam)
air
4.2 x 103
2.1 x 103
2.0 x 103
1.0 x 103