AP Chemistry
Unit #5
Chapter 6 – Zumdahl & Zumdahl
Thermochemistry
Thermodynamics
I.
The Nature of Energy
A. Definitions of basic concepts:
1. Thermodynamics –
2. Energy –
3. Work –
4. Heat –
5. Newton (N) –
6. Joule (J) –
7. calorie (cal) –
B. Kinetic vs. Potential Energy
1. Kinetic Energy –
2. Potential Energy –
C. Systems and Surroundings
1. Systems –
a) Open Systems –
b) Closed Systems –
2. Surroundings –
II.
The First Law of Thermodynamics:
A. Internal Energy –
B. Energy Diagrams –
C. Relating ∆E to Heat and Work –
Heat added to the system
Work done on the system
Sample Problem 6.1: 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.
D. State Functions –
State functions depend only on the present state of the system, not the past
history of how it got to that state.
A change in state is independent of the particular pathway taken between the two
states
E. PV Work –
The magnitude of work required to expand a gas (∆V) against a pressure (P) is
defined by the following equation:
For an expanding gas, ∆V is a positive quantity because the volume is
increasing. Thus ∆V and w must have opposite signs, which leads to the
equation:
Note that for a gas expanding against an external pressure P, w is negative
quantity as required, since work flows out of the system. When a gas is
compressed, ∆V is a negative quantity (the volume decreases), which makes w
a positive quantity (work flows into the system)
Sample Problem 6.2: Calculate the work associated with the expansion of a gas from 46 L
to 64 L at a constant external pressure of 15 atm.
In dealing with “PV work,” keep in mind that the P in P∆V always refers to the
external pressure—the pressure that causes a compression or that resists an
expansion
Sample Problem 6.3: A balloon is being inflated to its full extent by heating the air inside it.
In the final stages of this process, the volume of the balloon changes from 4.00 x 106 L to
4.50 x 106 L by the addition of 1.3 x 108 J of energy as heat. Assuming that the balloon
expands against a constant pressure of 1.0 atm, calculate ∆E for the process. (To convert
between L atm and J, use 1 L atm = 101.3 J)
III.
Heat and Enthalpy Changes
A. Endothermic vs. Exothermic
1. Endothermic – reactions in which the system absorbs energy from the
surroundings
2. Exothermic – reactions in which energy flows out of the system into the
surroundings
B. Enthalpy (H) – The heat energy absorbed or released by the system under constant
pressure
∆H =
∆H = Hfinal – Hinitial
IV.
Enthalpies of Reaction (Hrxn)
A. Defn –
Example)
2H2 (g) + O2 (g) 2H2O (g)
∆H = - 483.6 kJ
B. Things to Remember about Enthalpy:
1. Enthalpy is an extensive property –
Sample Problem 5.2) How much heat is released when 4.50g of methane gas is burned in
a constant pressure system?
CH4 (g) + 2O2 (g) CO2 (g) + H2O (g) ∆H = -802 kJ
2. The ∆Hrxn is equal in magnitude but opposite in sign to the ∆Hrxn of the reverse rxn-
3. The ∆Hrxn depends on the state of the reactants and products-
V.
Calorimetry
Defn –
A. Heat Capacity and Specific Heat –
B. The Specific Heat Equation
q = m C ∆T
Sample Problem 5.3) a) How much heat is required to raise the temperature of 250g of water from 22°C to near
its boiling point, 98°C? The specific heat of water is 4.18 J/g-K
b) What is the molar heat capacity of water?
C. Constant Pressure Calorimetry
Sample Problem 5.4) A student mixes 50mL of 1.0 M HCl and 50mL of 1.0 M NaOH in a coffee cup calorimeter,
the temperature of the resultant solution increases from 21.0°C to 27.5°C. Calculate the enthalpy change for the
reaction, assuming that no energy is lost to the surroundings, and that the volume of the solution is 100 mL and
with a density of 1.0 g/mL, and its specific heat is 4.18 J/g-K.
D. Bomb Calorimetry
Sample Problem 5.5) N2H4(l) + O2(g) N2(g) + 2H2O(g)
When 1.00 g of N2H4 is burned in a bomb calorimeter, the temperature increases by 3.51°C. If the calorimeter
has a heat capacity of 5.510 kJ/°C, what is the quantity of heat evolved? What amount of heat is evolved upon
the combustion of a mole of N2H4?
VI.
Hess’s Law
Defn-
A. Heat of Combustion
•
o
Standard Enthalpy of Combustion (∆
∆ Hc )
Is defined as the enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions.
o
Since energy is usually released in such a reaction ∆Hc will usually be negative.
o
-1
e.g. ∆Hc [C2H6(g)] = - 1560 kJmol means
C2H6(g) + 3½O2(g) 2CO2(g) + 3H2O(l)
o
-1
∆H = - 1560 kJmol
It is useful to remember that compounds containing some combination of carbon, hydrogen and oxygen, when completely
burned in air (O2), produce carbon dioxide and water only. Other compounds may require other knowledge or intelligent
guesswork to determine the products of combustion.
Task 5.1
Write equations that represent the following processes.
a) The standard enthalpy of combustion of CO(g)
b) The standard enthalpy of combustion of C6H14(l)
c) The standard enthalpy of combustion of C6H5CO2H(s)
d) The standard enthalpy of combustion of H2(g)
e) The standard enthalpy of combustion of Zn(s)
Sample Problem 5.6) The enthalpy of the combustion of C to CO2 is – 393.5 kJ/mol C, and the enthalpy of
combustion of CO to CO2 is –283.0 kJ/mol CO.
1
C(s) + O2(g) CO2(g) ∆H = - 393.5 kJ
2
CO(g) + ½ O2(g) CO2(g)
∆H = - 283.0 kJ
Calculate the Hcombustion of C to CO.
VII.
Enthalpies of Formation (Hf)
A. Enthalpy of Formation (∆H) –
B. Standard State –
C. Standard Enthalpy of Formation (Hf°) –
D. Using Hf to calculate Hrxn –
∆H°rxn = ∑ n∆H°f(products) - ∑ m∆H°f(reactants)
°
Sample Problem 5.8) Compare the quantity of heat produced by combustion of 1.00g of C3H8 (∆H f = -103.85
°
kJ/mol) to that produced by 1.00 g of C6H6 (∆H f = 49.04 kJ/mol)
Sample Problem 5.9) The Standard Enthalpy change for the reaction:
°
CaCO3 (s) CaO (s) + CO2 (g) ∆H rxn= 178.1 kJ
°
°
Calculate the standard enthalpy of formation of CaCO3(s) if CaO (s) ∆H f = -1207.1 kJ/mol and CO2 (g) ∆H f = 393.5 kJ/mol