1. Unit 1: Thermochemistry
Introduction
The First Law of Thermodynamics
Enthalpy
Enthalpy of Reaction
Calorimetry
Hess’s Law
Enthalpy of Formation

2. Introduction
Most daily activities involve processes that either use or produce energy:
Activities that produce energy
Metabolism of food
Burning fossil fuels
Activities that use energy:
Photosynthesis
Pushing a bike up a hill
Baking bread

3. Introduction Thermodynamics
The study of energy and its transformations
Thermochemistry:
A branch of thermodynamics
The study of the energy (heat) absorbed or released during chemical reactions

4. Introduction Objects can have two types of energy: Kinetic energy
Energy of motion
Thermal energy
The type of kinetic energy a substance possesses because of its temperature
Potential energy
Energy of position
“stored” energy resulting from the attractions and repulsions an object experiences relative to other objects

5. Introduction Attractive and repulsive forces include: Gravity
Electrostatic forces between charged particles
e- has potential energy due to its position near the positively charged nucleus
Most important attractive/replusive forces in chemistry

6. Introduction
Attractive and repulsive forces within a substance lead to a type of potential energy called chemical energy
The potential energy stored in substances resulting from the arrangements of the atoms in the substance

7. Introduction Units of Energy SI unit = joule (J)
1 J = the kinetic energy of a 2 kg mass moving at a speed of 1 m/s
A very small quantity
Kilojoule (kJ)
1 kJ = 1000 J

8. Introduction Units of Energy (cont) Calorie (cal)
Originally defined as the amount of energy needed to raise the temperature of 1g of water from 14.5oC to 15.5oC.
1 cal = J (exactly)
Kilocalorie (kcal)
1 kcal = 1000 cal

9. Introduction
On the exam, you must be able to convert from one set of energy units to the other using dimensional analysis.
You must know the conversion factors given on the previous slides!!!

10. Introduction
Example: Convert 725 cal to kJ.

11. Introduction Example: A particular furnace produces
9.0 x 104 BTU/hr of heat. If 1.00 BTU = cal, use dimensional analysis to calculate the number of kJ of heat delivered by the furnace after running for 2.50 hours.

12. Introduction
When using thermodynamics to study energy changes, we generally focus on a limited, well-defined part of the universe.
System:
The portion of the universe singled out for study
Surroundings:
Everything else

13. Introduction The system is usually the chemicals in the flask/reactor.
The flask and everything else belong to the surroundings.

14. Introduction Open system:
A system that can exchange both matter and energy with the surroundings
Closed system:
A system that can exchange energy with the surroundings but not matter
A cylinder with a piston is one example of a closed system.

15. Introduction
In a closed system energy can be gained from or lost to the surroundings as:
Work
Heat
Work:
Energy used to cause an object to move against a force
Lifting an object
Hitting a baseball

16. Introduction
Heat:
The energy used to cause the temperature of an object to increase
The energy transferred from a hotter object to a cooler one
Energy:
The capacity to do work or to transfer heat

17. Potential energy Kinetic energy
Introduction
The potential energy of a system can be converted into kinetic energy and vice versa.
Energy can be transferred
back and forth between the system and the surroundings
as work and/or heat.
Potential energy Kinetic energy
work

18. The First Law of Thermodynamics
Although energy can be converted from one form to another and can be transferred between the system and the surroundings:
Energy cannot be created or destroyed.
(First Law of Thermodynamics)
Any energy lost by the system must be gained by the surroundings and vice versa.

19. The First Law of Thermodynamics
The First Law of Thermodynamics can be used to analyze changes in the Internal Energy (E) of a system.
The sum of all kinetic and potential energy of all components of a system
For molecules in a chemical system, the internal energy would include:
the motion and interactions of the molecules
the motion and interactions of the nuclei and electrons found in the molecules

20. The First Law of Thermodynamics
Internal Energy:
Extensive property
depends on mass of system
Influenced by temperature and pressure
Has a fixed value for a given set of conditions
State function

21. The First Law of Thermodynamics
The internal energy of a system is a state function.
A property of the system that is determined by specifying its condition or its state in terms of T, P, location, etc
Depends only on its present condition
Does not depend on how the system got to that state/condition

22. The First Law of Thermodynamics
The internal energy of a system can change when:
heat is gained from or lost to the surroundings
work is done on or by the system.
The change in the internal energy
D E = Efinal - Einitial
DE = change in internal energy
Efinal = final energy of system
Einitial = initial energy of system

23. The First Law of Thermodynamics
If Efinal > Einitial,
DE >0 (positive)
the system has gained energy from the surroundings.
endergonic

24. The First Law of Thermodynamics
The decomposition of water is endergonic (DE > 0):
2 H2O (l) H2 (g) + O2 (g)
H2 (g), O2 (g)
final
Energy must be gained from the surroundings. E
H2O (l)
initial

25. The First Law of Thermodynamics
If Efinal < Einitial,
DE < 0 (negative)
the system has lost energy to the surroundings.
exergonic

26. The First Law of Thermodynamics
The synthesis of water is exergonic (DE < 0)
2 H2 (g) + O2 (g) H2O (l)
H2 (g), O2 (g)
H2O (l)
Energy is lost to the surroundings in this reaction.
initial E
final

27. The First Law of Thermodynamics
The internal energy of a system can change when energy is exchanged between the system and the surroundings
Heat
Work
The change in internal energy that occurs can be found:
D E = q + w
Where q = heat
w = work

28. The First Law of Thermodynamics
By convention:
q = positive
Heat added to the system
w = positive
Work done on the system by the surroundings
q = negative
Heat lost by the system
w = negative
Work done by the system on the surroundings

29. The First Law of Thermodynamics
Example: Calculate the change in internal energy of the system for a process in which the system absorbs 240. J of heat from the surroundings and does 85 J of work on the surroundings.

30. The First Law of Thermodynamics