1. A First Look at Thermodynamics…
Energy
A First Look at Thermodynamics…

2. Definition… Ability to do “work”; capacity of a system to do “work”
“Work” is defined for macroscopic systems as moving an object against a force. But chemistry deals with many microscopic systems so the physics definition of “work” does not always seem to apply.

3. Forms… Potential – energy due to position
Kinetic – energy due to motion
These forms of energy are interconvertible.

5. Energy is the capacity to do work.
less stable
change in potential energy EQUALS
kinetic energy
more stable
A gravitational system. The potential energy gained when a lifted weight is converted to kinetic energy as the weight falls.

6. Energy is the capacity to do work.
less stable
change in potential energy EQUALS
kinetic energy
more stable
A system of two balls attached by a spring. The potential energy gained by a stretched spring is converted to kinetic energy when the moving balls are released.

7. Energy is the capacity to do work.
less stable
change in potential energy EQUALS
kinetic energy
more stable
A system of oppositely charged particles. The potential energy gained when the charges are separated is converted to kinetic energy as the attraction pulls these charges together.

8. Energy is the capacity to do work.
less stable
change in potential energy EQUALS
kinetic energy
more stable
A system of fuel and exhaust. A fuel is higher in chemical potential energy than the exhaust. As the fuel burns, some of its potential energy is converted to the kinetic energy of the moving car.

9. Types of Energy There are also different types of energy: chemical
electrical
gravitational
heat
light (electromagnetic radiation)
magnetic
mechanical
nuclear
These are also interconvertible.

11. Laws!
No matter what the conversion the system must obey the Law of Conservation of Energy.
Energy cannot be created or destroyed.
(but may be converted to mass under certain conditions)

12. Thermal Energy
In chemistry we deal primarily with thermal energy. We also encounter light and electrical energies.
Thermal energy = energy due to molecular or atomic motion
Heat (q) = transfer of chemical energy due to a temperature difference
Thermodynamics = study of heat and its transformations.
Thermochemistry = branch of thermodynamics that deals with the heat involved with chemical and physical changes.

13. Temperature Measure of the “hotness” or “coldness” of something
Measure of the effect of heat on an object or system
Measure of particle motion
Heat always flows from an area of higher temperature to areas of lower temperatures.

14. Temperature Measurement
Typically measured using a thermometer which has an expandable liquid (mercury or alcohol) trapped in a sealed cylinder
US meteorological and heating unit temperatures are expressed using the Fahrenheit scale.
Water freezes at 32oF and boils at 212oF.

15. Figure 1.8
The freezing and boiling points of water.

16. Temperature Measurement
Other countries and the scientific community use the Celsius scale.
Water freezes at 0oC and boils at 100oC.

17. The freezing and boiling points of water.
Figure 1.8
The freezing and boiling points of water.
Silberberg, Principles of Chemistry

18. Temperature Measurement
There is a problem using either of these scales in certain systems. Because there are negative temperatures certain mathematical relationships generate impossible results.
Lord Kelvin generated a new scale with the coldest temperature as 0 K (absolute zero). The unit of temperature is a kelvin (K). The term degree is not used.
Water freezes at K and boils at K.

19. The freezing and boiling points of water.
Figure 1.8
The freezing and boiling points of water.
Silberberg, Principles of Chemistry

20. Temperature Conversions
oF  oC
oC  K

21. Systems…
Changes in thermal energy of systems are determined by transfer of heat (q) observed by changes in temperature.
First you need to define what your system and surroundings are…
When heat moves out of a system to the surroundings the process is exothermic for the system.
When heat moves into a system from the surroundings the process is endothermic for the system.

22. A system transferring energy as heat only.
Exothermic
Endothermic

23. Heat Units Joule (J) 1 J = 1 kg*m2/s2 Calorie (cal) 1 cal = 4.18J
British Thermal Unit
1 Btu = 1055 J

24. Specific Heat Capacity
Substances respond differently to the same quantity of energy.
Consider a wooden spoon and a metal spoon in a pot of boiling water. Which spoon gets hotter?
Specific heat capacity (c) – energy required to raise 1 gram of substance by 1oC
J/g. oC (J g-1 oC-1) or kcal/g. oC (kcal g-1 oC-1)

26. Heat Calculations PROBLEM:
A layer of copper welded to the bottom of a skillet weighs 125 g. How much heat is needed to raise the temperature of the copper layer from 250C to 300.0C? The specific heat capacity (c) of Cu is J/g*K.
PLAN:
Given the mass, specific heat capacity and change in temperature, we can use q = c x mass x DT to find the answer. DT in 0C is the same as for K.
SOLUTION:
q =
0.387 J
g*K
125 g x
(300-25)0C
= 1.33x104 J x

27. Enthalpy The Meaning of Enthalpy w = - PDV DH ≈ DE in H = E + PV
Silberberg, Principles of Chemistry
Enthalpy
The Meaning of Enthalpy
w = - PDV
DH ≈ DE in
H = E + PV
1. Reactions that do not involve gases.
where H is enthalpy
2. Reactions in which the number of moles of gas does not change.
DH = DE + PDV
3. Reactions in which the number of moles of gas does change but q is >>> PDV.
qp = DE + PDV = DH
For most of the processes we will encounter this semester
we can use ΔH as heat exchanged by a system.

28. Enthalpy
ΔH for a reaction is calculated as: ΔHproducts – ΔHreactants If the ΔH is positive, energy of the products is higher than reactants and the process is endothermic. If the ΔH is negative, energy of the reactants is higher than products and the process is exothermic.

29. Enthalpy diagrams for exothermic and endothermic processes.
CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
H2O(l) H2O(g)
CH4 + 2O2
H2O(g)
Hfinal
Enthalpy, H
Hinitial
Enthalpy, H
heat out
heat in
DH < 0
DH > 0
CO2 + 2H2O
H2O(l)
Hfinal
Hinitial
A Exothermic process
B Endothermic process

30. Determining the heat exchange
in a chemical change

31. A bomb calorimeter

32. molar ratio from balanced equation
AMOUNT (mol)
of compound A
Summary of the relationship between amount (mol) of substance and the heat (kJ) transferred during a reaction.
AMOUNT (mol)
of compound B
molar ratio from balanced equation
HEAT (kJ)
gained or lost
DHrxn (kJ/mol)

33. Using the Heat of Reaction (DHrxn) to Find Amounts
PROBLEM:
The major source of aluminum in the world is bauxite (mostly aluminum oxide). Its thermal decomposition can be represented by
If aluminum is produced this way, how many grams of aluminum can form when 1.000x103 kJ of heat is transferred?
Al2O3(s) Al(s) + 3/2O2(g) DHrxn = 1676 kJ
2 mol Al
1676 kJ
26.98 g Al
1 mol Al
SOLUTION:
1.000x103 kJ x
= g Al

34. Reaction Progress
All reactions have an energy barrier that must be overcome.
Activation energy (Ea)
This is an exothermic
reaction.
Energy of reactants
Energy of products

35. Reaction Progress This diagram illustrates an endothermic reaction.
Activation energy (Ea)
This diagram
illustrates an
endothermic
reaction.
Energy of reactants
Energy of products

36. Reaction Enthalpy The difference between the starting
and ending energies
is ΔH. ΔH

37. Catalysts reduce the activation energy of a reaction.
Reaction energy diagram of a catalyzed and an uncatalyzed process.

38. Disorder aka Entropy
Systems tend to greater disorder or greater energy dispersal.
Disorder is the preferred state.
Energy dispersal or system disorder is called entropy.
All systems work to achieve lowest energy and greatest disorder or a balance between these goals.