1. Chemical Thermodynamics

2. Introduction
Matter
Any object that has mass and volume. Things that you interact with in the physical world
Energy
The potential to do work
Work
Change in self or surroundings
Is there only one form of energy?
Okay smarty-pants, what are they?

3. Introduction Key Concept!
Two things can happen to a form of energy as it moves along. What are they?
(think about what happens when energy comes in contact with your body)
Energy can change forms or do work (cause change in matter).
What do we call that principle (law)?

4. Introduction The Law of Conservation of Energy
Energy does not just vanish; it either changes forms or is converted into work.
Galileo: pendulum (conversion of potential and kinetic energy)
Leibniz: mathematical analysis of the conservation of kinetic energy

5. Chemical Thermodynamics
Energy is exchanged or transformed in all chemical reactions and physical changes of matter.
What is Thermal energy?
Internal kinetic energy, the motion of the particles (e.g. molecules, atoms) of an object

6. Chemical Thermodynamics
How is thermal energy measured?
Temperature
Students know how to describe temperature and heat flow in terms of the motion of molecules (or atoms).

7. Chemical Thermodynamics
What is Heat?
Transfer of thermal energy.
How does heat flow?
(from high temperature to low temperature)

8. Chemical Thermodynamics
Often, heat is described as flowing from the system to the surroundings or from the surroundings to the system.
The system is defined by its boundaries, and the surroundings are outside the boundaries, with “the universe” frequently considered as the surroundings.

9. Chemical Thermodynamics

11. Chemical Thermodynamics
Energy and Chemical Bonds
breaking a bond always requires energy
making a bond almost always releases energy.
The amount of energy per bond depends on the strength of the bond.
Because of this chemical processes can release or absorb thermal energy.

12. Chemical Thermodynamics
Endothermic processes absorb heat
Exothermic processes release heat
The net heat released to or absorbed from the surroundings comes from the making and breaking of chemical bonds during a reaction.

13. Endothermic Processes
Endothermic processes absorb heat, and their equations can be written with heat as a reactant.
A + B + nrg → C + D
e.g.
6CO2(g) + H2O(l) + nrg → C6H12O6(aq) + 6O2(g)
The equation above represents what process?
Photosynthesis

14. Exothermic Processes
Exothermic processes release heat, and their equations can be written with heat as a product.
e.g.
C6H12O6(aq) + 6O2(g) → 6CO2(g) + H2O(l) + nrg
The equation above represents what process?
Respiration
Combustion

15. Exothermic Processes
The net heat released to or absorbed from the surroundings comes from the making and breaking of chemical bonds during a reaction.

16. Graph of potential energy over time for a chemical reaction
Exothermic Reaction
Delta H: Enthalpy, internal energy
An exothermic rxn releases nrg into the surroundings
The nrg of products is lower than the nrg of reactants

17. Graph of potential energy over time for a chemical reaction
Endothermic Reaction
Delta H: Enthalpy, internal energy
An exothermic rxn releases nrg into the surroundings
The nrg of products is higher than the nrg of reactants

18. Phase Changes What physical states can matter be found in? (4)
solid, liquid, gas, plasma
What do you call changing from liquid to gas?
evaporation
Changing from solid to liquid?
melting
Changing from liquid to solid?
freezing
Changing from gas to liquid?
condensation

19. Phase Changes Physical changes are accompanied by changes in
internal energy.
Evaporation and melting require energy to overcome the bonds of attractions in the corresponding liquid or solid state.

20. Phase Changes Physical changes are accompanied by changes in
internal energy.
Condensation and freezing
release heat to the surroundings
as internal energy is reduced and
bonds of attraction are formed.

21. Problem Solving
Specific heat is the energy needed to change the temperature of one gram of substance by one degree Celsius.
The unit of specific heat is joule/gram-degree (J/g oC)
Q = m ▪ Cp ▪ ΔT
Q = Heat Energy (in Joules)
m = mass (grams)
Cp or s  = specific heat (J/g oC)
ΔT = change in temperature (Celsius)

22. Problem Solving ΔT = change in temperature (Celsius)
ΔT = Tfinal – Tinitial
ΔT = Tf – Ti
ΔT = T2 – T1

23. Problem Solving
Example: How much heat energy is absorbed when 88.0 g of water is heated from 5.00o C to 37.0o C?  ( CP of H2O = J/ g xoC)
What do we know?
m =
CP =
ΔT =
Q =
Q = m ▪ Cp ▪ ΔT

24. Problem Solving
Example: How much heat energy is absorbed when 88.0 g of water is heated from 5.00o C to 37.0o C?  ( CP-H2O = J/ g xoC)
What do we know?
m = 88.0 g
CP = J / g x C
ΔT = 37o C - 5o C = 32o C
Q = ?
Q = (88.0g)( 4.184J/go C)(32oC)
Q = J
Q = m ▪ Cp ▪ ΔT
or 1.18 x 104 J
or 11.8 x KJ

25. Enthalpy
Enthalpy is a measure of the total energy of a thermodynamic system.
What do you expect to see as you add energy to a substance?

27. Latent Heat
During phase changes, energy is added or removed without a corresponding temperature change.
This phenomenon is called latent (hidden) heat.
There is a latent heat of fusion and a latent heat of vaporization.
The unit of latent heat is joule/gram or kilojoule/mole.

28. Latent Heat
Phase change graph, temp vs. energy

29. Latent Heat
Look at the graph and think. What is happening?

30. Heat of Vaporization (ΔHvap)
Problem Solving
Specific Heat (Cp)
Latent
Heat of Fusion (ΔHfus)
Heat of Vaporization (ΔHvap)
Water (H2O)
4.18 J/g °C
334 J/g
2257 J/g
How much heat energy is required to melt 100. g of H20? 
Q = m x ΔH
Q = 100g x 334 J/g
= J
= 33.4 kJ

31. Questions?