1. Energy and Chemical Change
Thermodynamics
Energy and Chemical Change

2. Energy The ability to do work or produce heat.
Two basic forms of energy: potential and kinetic energy.
Potential energy is due to the composition or position of an object.
Kinetic energy is the energy of motion.

3. Law of Conservation of Energy
Energy can be converted from one form to another,
but can neither be created nor destroyed.
Chemical potential energy is the energy stored in
a substance due to the arrangement of atoms and
the strength of the chemical bonds.

4. Heat q Energy that flows from a warmer object to a cooler object.
As a warmer object loses heat, the temperature decreases.
SI unit of heat and energy is the Joule (J).

5. Calories are also often used as heat units. A
calorie is defined as the amount of energy required
to raise the temperature of one gram of pure water
by one degree Celsius.
Food calories (C) are the equivalent of 1000 calories.
1 cal = J
Convert 86.5 J to cal; convert 142 Cal to Joules.

6. Specific Heat
Amount of energy required to raise one gram of a substance one degree Celsius.
Water has an extremely high specific heat (4.18 J) which means that a large amount of energy must be used to raise the temperature of a quantity of water.

7. In order to calculate the heat energy that must be
absorbed or released by a substance, you need to
know the amount of the substance (mass in grams),
the specific heat capacity (Cp) and the temperature
change that the substance will undergo.
q = m Cp T

8. What is the specific heat of iron if a 10.0 g sample
changes from 50.4oC to 25.0oC with the release of
114 J of heat?

9. If the temperature of 34.4 g of ethanol increases
from 25.0oC to 78.8oC, how much heat is absorbed
by the ethanol? The specific heat of ethanol is
2.42 J/goC.

10. What temperature change will 100.0 mL of water
undergo when it absorbs 1360 Joules of heat?

11. Thermochemistry: the study of heat changes that
accompany chemical reactions and phase changes.
System: the specific part of the universe under study.
Surroundings: everything in the universe other than
the system.
Universe = System + Surroundings

12. The heat required for one mole of a substance to
change phase is called the molar heat of:
fusion when s l or
vaporization when l  g
Energy is released when l  s or g  l and
Energy is absorbed when s  l or l  g

14. Calculate the heat required to melt 25.7 g of solid
methanol at its melting point. The molar heat of
fusion for methanol is 3.22 kJ/mol.

15. How much heat is evolved when 275 g of ammonia
gas condenses to a liquid at its boiling point? The
molar heat of vaporization for ammonia is
23.3 kJ/mol.

16. A thermochemical equation is a balanced chemical
equation that includes the physical states of all
reactants and products and the energy change.
C6H12O6 (s) + 6O2(g)  6CO2(g) + 6H2O(l)
H = kJ

17. Enthalpy (H) is the heat content of a system at
constant pressure.
The change in enthalpy for a reaction is called
the Heat of reaction (Hrxn = Hfinal – Hinitial ) or
Hrxn = Hproducts – Hreactants

18. The sign for the enthalpy change depends
on whether the products or reactants have
more energy.

19. Endothermic Reactions:
Reactants + heat  Products Hrxn = +
Hproducts > Hreactants

20. Exothermic Reactions:
Reactants  Products + heat Hrxn = -
Hproducts < Hreactants

23. If A  B and B  C what can you conclude?

24. Hess’s Law:
In going from a particular set of reactants
to a particular set of products, the change in
enthalpy is the same whether the reaction takes
place in one step or a series of steps.
A  C = A  B
B  C
A  C

25. Given:
2CO(g) + O2 (g)  2CO2(g) H = ?
Use these:
C(s) + O2(g)  CO2(g) H = kJ
2C(s) + O2(g)  2CO(g) H = kJ

26. Calculate Hrxn for the reaction:
HCl(g) + NH3(g)  NH4Cl(s)
given the following:
H2(g) + Cl2(g) 2HCl(g) H=-184kJ
N2(g) + 3H2(g)  2NH3(g) kJ
N2(g) + 4H2(g) + Cl2(g)  2NH4Cl(s) kJ

27. Calculate ΔH for the reaction:
2C(s) + H2(g)  C2H2(g)
Given:
ΔH=
C2H2(g) + 5/2 O2(g)  2CO2(g) + H2O(l)
C(s) + O2(g)  CO2(g)
H2(g) + ½ O2(g)  H2O(l)

28. Given the following reaction:
Ca(s) + H2O(l) + CO2(g)  H2(g) + CaCO3(s)
Use the following reactions:
CaCO3(s) CaO(s) + CO2(g) ΔH = kJ
Ca(s) + 2 H2O(l)  Ca(OH)2(s) + H2(g) kJ
Ca(OH)2(s)  CaO(s) + H2O(l) kJ

29. Hess’s Law Quiz Tomorrow
Find ΔHorxn for
C2H2(g) + 2 H2(g)  C2H6(g)
Using the standard enthalpies of reaction below:
2 H2(g) + O2(g)  2 H2O(l) kJ
2 C2H2(g) O2(g)  4 CO2(g) + 2 H2O(l)
-2598 kJ
2 C2H6(g) O2(g)  4 CO2(g) + 6 H2O(l)
-3122 kJ

30. Enthalpies of Formation
The standard enthalpy of formation of a compound is called ΔHof
ΔHof is the change in enthalpy for the reaction that forms 1 mole of the compound from its elements.
All elements must be in their standard states at 298 K.

31. The enthalpy of formation for ethanol is
kJ. Write the equation for the formation
of ethanol.

32. Standard enthalpies of formation can be used to
determine the enthalpy of reaction using:
Hrxn =  Hproducts –  Hreactants
H2S(g) + 4F2(g)  2HF(g) + SF6(g)

33. Write the thermochemical equation for the
formation of zinc nitrate.
Determine ΔH for the following reaction:
NH4NO3(s)  N2O(g) + 2H2O(l)

34. Laws of Thermodynamics
1st Law- the energy of the universe is constant.
Keeps track of thermodynamics doesn’t correctly predict spontaneity.
Entropy (S) is disorder or randomness
2nd Law: the entropy of the universe increases.

36. Entropy S
Measure of the randomness or disorder of the particles in a system.
Gas > Liquid > Solid
Substances at higher temperatures have a higher entropy than substances at lower temperatures.

37. Determine whether the following results in an
increase or decrease in the entropy:
CO2(s)  CO2(l)
CaO(s) + CO2(g)  CaCO3(s)
Ag+(aq) + Cl-(aq)  AgCl(s)

38. For each of the following pairs, choose the
substance with the higher entropy:
Ar(l) or Ar(g)
1 mol As at 400K or 1 mol As at 298K
100g Na2SO4(s) at 30oC or 100g Na2SO4(l)
at 30oC or 100g

39. Spontaneous
A reaction that will occur without outside intervention.
We can’t determine how fast.
We need both thermodynamics and kinetics to describe a reaction completely.
Thermodynamics compares initial and final states.
Kinetics describes pathway between.

40. Spontaneous Reactions
Most spontaneous reactions are exothermic.
However, some endothermic reactions occur spontaneously due to an increase in entropy of the system.
Two factors need to be considered, enthalpy and entropy in deciding spontaneity

41. Gibbs Free Energy (G) is the energy available
to do work.
G = H - T S
G = - reaction is spontaneous
G = + reaction is not spontaneous

42. H S G
Charac-teristics - +
Spon.
Non-
Low T
High T

43. The enthalpy of formation of hydrogen chloride is
kJ/mol.
1. Write the thermochemical equation for the
formation of hydrogen chloride.
2. Determine the heat of reaction when hydrogen
chloride reacts with ammonia to form
ammonium chloride with an enthalpy of formation of kJ/mol.