• Thermochemistry:
Study of
that occur during chemical reactions and changes in state

can either occur through heat transfer or work
• heat (q), energy is transferred from a warmer object to a cooler object (always)
• Adding heat increases temperature

17.1: Endothermic & Exothermic processes
•
Kinetic energy vs. potential energy energy due to motion energy due to a substance's chemical composition
• potential energy is determined by the strength of repulsive and attractive forces between atoms
• In a chemical reaction, atoms are recombined into new arrangements that have different potential energies
•change in potential energy: due to absorption and release of energy to and from the surroundings

17.1: Endothermic & Exothermic processes
•
Two parameters crucial in Thermochemistry: a) system-part of the universe attention is focused b) surroundings-everything else in the universe
•
System + surrounding = universe
•
Fundamental goal of Thermochem.: study the heat flow between the system and its surroundings

17.1: Endothermic & Exothermic processes
• If
System (energy) Surrounding
(energy) by the same amount
• the total energy of the universe does not change
Law of conservation of energy

17.1: Endothermic & Exothermic processes
• Direction of heat flow is given from the point of view of the system
• So, endothermic process : system absorbs heat from the surroundings (system heats up)
• Heat flowing into the system = +q
• Exothermic process
: heat is released into the surroundings (system cools down, -q)

17.1: Endothermic & Exothermic processes
Example 1:
A container of melted paraffin wax is allowed to stand at room temperature (r.t.) until the wax solidifies. What is the direction of heat flow as the liquid wax solidifies? Is the process exothermic or endothermic?
Answer: Heat flows from the system (paraffin) to the surroundings (air)
Process: exothermic

17.1: Endothermic & Exothermic processes
Example 2:
When solid Ba(OH) with solid NH
4
2
▪8H
2
O is mixed in a beaker
SCN, a reaction occurs. The beaker quickly becomes very cold. Is the reaction exothermic or endothermic?
Answer: Endothermic surrounding = beaker and air
System = chemicals within beaker

• Two units used: a) calorie (cal)—amount of heat required to raise the temperature of 1g of pure water by
1 o C b) joule (j)—1 joule of heat raises the temperature of 1 g of pure water 0.2390
o C
• Joule = SI unit of energy

• Heat capacity is the quantity of heat needed to raise the temperature of an object exactly 1 o C
• Heat capacity depends on: a) mass b) chemical composition
• So, the greater the mass the greater the heat capacity eg.: cup of water vs. a drop of water
(cup of water = greater heat capacity)

• Specific heat
: amount of heat required to raise the temperature of 1g of a substance by 1 o C
• Table 17.1 (p.508): List of specific heats of substances

C = q = heat (joules/calories) m * ΔT mass (g) * ΔTemp. ( o C)
ΔT = T f
–T i
(T f
= final temperature)
(T i
= initial temperature)
C = j
(g * o C) or cal
(g * o C)

1. The temperature of a 95.4 g piece of copper increases from 25.0 o C to 48.0 o C when the copper absorbs 849 j of heat. What is the specific heat of copper?
unknown: C cu
Know: mass copper = 95.4 g
ΔT = T f
– T i
= 23.0 o C
= (48.0 o C – 25.0 o C) q = 849 j

C = q m * ΔT
C = 849 j
95.4 g * 23.0 o C
C = 0.387 j/g * o C
Sample problem 17.1, page 510

2. How much heat is required to raise the temperature of 250.0 g of mercury (Hg) 52 o C?
unknown: q
Know: mass Hg = 250.0 g
ΔT = 52.0 o C
C
Hg
= 0.14 j/(g * o C)

C = q m * ΔT q = C
Hg
* m * ΔT q = 0.14 (j/g * o C) * 250.0 g * 52 o C q = 1.8 x 10 3 j (1.8 kj)
Problem #4 page 510

Measuring Enthalpy Changes

17.2: Enthalpy (measuring heat flow)
• Calorimetry: accurate measurement of heat flow into or out of a system in chemical and physical processes
• In calorimetry, heat released by a system is equal to the heat absorbed by its surroundings and vice versa
• Instrument used to measure absorption or release of heat is a calorimeter

17.2: Enthalpy (measuring heat flow)
Two types of calorimeters: a) Constant-Pressure calorimeter (eg. foam cups)
• As most reactions occur at constant pressure we can say that:
A change in enthalpy (Δ H ) = heat supplied (q)
• So , a release of heat (exothermic) corresponds to a decrease in enthalpy (at constant pressure)
• An absorption of heat (endothermic) corresponds to an increase in enthalpy (constant pressure)

17.2: Enthalpy (measuring heat flow) b) Constant-Volume Calorimeters (eg. bomb calorimeters)
• Substance is burned (in the presence of O
2
) inside a chamber surrounded by water (high pressure)
• Heat released warms the water
Figure 17.6 (p. 512) Bomb calorimeter.

17.2: Thermochemical equations
CaO
(s)
+ H
2
O
(l)
Ca(OH)
2
+ 65.2 kJ
Heat released
•
A chemical equation that includes the enthalpy change is called a thermochemical equation
• Reactants and products at their usual physical state (at 25 o C) given at standard pressure (101.3 kPa)
• So, the heat of reaction (or Δ
H ) for the above equation is -65.2 kJ

So, rewrite the equation as follows:
CaO
(s)
+ H
2
O
(l)
Ca(OH)
2(s)
  
65.2 kJ
• Other reactions absorb heat from the surroundings, eg.:
2 NaHCO
3(s)
+ 129 kJ
Na
2
CO
3(s)
+ H
2
O
(g)
+ CO
2(g)
2
Rewrite to show heat of reaction
NaHCO
3(s)
Na
2
CO
3(s)
+ H
2
O
(g)
+ CO
2(g)
  
kJ

• Amount of heat released/absorbed during a reaction depends on the number of moles of reactants involved eg.: 2 NaHCO
3(s)
4 NaHCO
3(s)
Na
2
CO
3(s)
+ H
2
O
(g)
+ CO
2(g)
  
kJ
Na
2
CO
3(s)
+ H
2
O
(g)
+ CO
2(g)
  
258 kJ

CaO
(s)
+ H
2
O
(l)
Na
2
CO
3(s)
+ H
2
O
(l)
+ CO
2(g)

H = 129 kJ

H = -65.2 kJ
Ca(OH)
2(s)
Exothermic Reaction
2 NaHCO
3(s)
Endothermic Reaction
Diagram A:
Enthalpy of reactants greater than of products
Diagram B:
Enthalpy of reactant less than of products

17.2: Thermochemical equations
Physical states of reactants and products must be stated:
H
2
O
(l)
H
2(g)
+ 1 O
2(g)
2
  
285.8 kJ
H
2
O
(g)
H
2(g)
+ 1 O
2(g)
2
  
241.8 kJ
Difference = 44.0 kJ
Vaporization of H
2
O
(l) requires more heat (44.0 kJ)

1. Calculate the amount of heat in (kJ) required to decompose 2.24 mol
NaHCO
3(S)
2 NaHCO
3(s)
Na
2
CO
3(s)
+ H
2
O
(g)
+ CO
2(g)
  
kJ
Known:
•2.24 mol NaHCO
3 decomposes
•Δ H = 129 kJ (2 mol NaHCO
3
)
Solve:
Unknown:
•Δ
H = ?
129 kJ = ΔH
2 mol NaHCO
3(s)
2.24 mol NaHCO
3(s)
Δ H = (129 kJ) * 2.24 mol NaHCO
3(s)
2 mol NaHCO
3(s)
Δ H = 144 kJ
Sample problem 17.3; p. 516

2. When carbon disulfide is formed from its elements, heat is absorbed. Calculate the amount of heat in (kJ) absorbed when 5.66 g of carbon disulfide is formed.

H = 89.3 kJ
C
(s)
+ 2 S
(s)
CS
2(l)
Known:
•5.66 g CS
2 is formed
•Δ H = 89.3 kJ (1 mol CS
2(l)
)
•Molar mass: CS
2(l)
: C = 12.0 g/mol
2 *S = 32.1 g/mol = 64.2 g/mol
76.2 g/mol
Unknown:
•Δ
H = ?

Example 2
Solve:
1. Moles CS
2(l)
= 5.66g CS
2
76.2 g/mol CS
2(l)
=
0.0743 mol CS
2(l)
2. 89.3 kJ
1 mol CS
2(l)
=
ΔH
0.074 mol CS
2(l)
Δ H = (89.3 kJ) * 0.074 mol CS
2(l)
1 mol CS
2(l)
Δ
H = 6.63 kJ

Objective:
-Heats of Fusion and Solidification
-Heats of Vaporization and Condensation
-Heat of solution

17.3: Heat of fusion and solidification
• The temperature remains constant when a change of state occurs via a gain/loss of energy
• Heat absorbed by 1 mole of a solid during melting (constant temperature) is the molar heat of fusion (Δ H fus
)
• Molar heat of solidification
(Δ H solid
) is the heat lost by 1 mole of liquid as it solidifies
(constant temperature)
• So, Δ
H fus
= Δ H solid

17.3: Heat of fusion and solidification
• Figure 17.9: Enthalpy changes and changes of state

17.3: Heat of fusion and solidification
H
2
O
(s)
H
2
O
(l)
H
2
O
(l)
H
2
O
(s)
  fus
 
.01 kJ/mol
  solid
 
6.01 kJ/mol

17.3: Heats of Vaporization and Condensation
• Molar heat of vaporization
(Δ H vap
): Amount of heat required to vaporize one mole of a liquid at the liquid’s normal boiling point
H
2
O
(l)
H
2
O
(g)
  vap
 
kJ/mol
•
Molar heat of condensation (Δ H cond
): heat released when 1 mole of vapor condenses
  cond
 
kJ/mol H
2
O
(g)
•So, Δ
H vap
= -Δ H cond
H
2
O
(l)

17.3: Heat of vaporization and condensation
• Figure 17.10: Heating curve of water

• There is heat released/gained when a solute dissolves in a solvent
• The enthalpy change due to 1 mole of a substance dissolving: molar heat of solution
(Δ H soln
)
H
2
O
(l)
NaOH
(s)
Na
+
(aq)
+ OH
-
(aq)
  soln
 
5.1 kJ/mol

• Applications: hot/cold packs
• Hot pack:
H
2
O
(l)
CaCl
2(s)
•Cold pack:
H
2
O
(l)
NH
4
NO
3(s)
Ca
2+
(aq)
+ 2Cl
-
(aq)
  soln
 
82.8 kJ/mol
NH
4
+
(aq)
+ NO
3
-
(aq)
  soln
 
25.7 kJ/mol