1. Thermochemistry

2. Specific Heat Formula Q = Energy (heat) lost or gained
cp = Specific Heat
T = Temperature change
m = Mass

3. Two Types of Thermal Reactions
Exothermic: Releases Thermal Energy (heat) sign is –
Endothermic: Absorbs Thermal Energy (heat)
sign is +

4. Exothermic Processes
Processes in which energy is released as it proceeds, and surroundings become warmer
Reactants  Products + energy

5. Endothermic Processes
Processes in which energy is absorbed as it proceeds, and surroundings become colder
Reactants + energy  Products

6. Enthalpy = H
In a chemical reaction, Enthalpy (H) is equal to the energy that flows as heat.
(at a constant pressure)
Example:
When 1 mol of methane gas is burned it releases 890 kJ of energy
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + heat
H = -890kj = exothermic reaction

7. Solution: Molar mass of CH4 = 16.04 g / mol
Calculate H for a process in which 5.8 g of methane are burned. H = -890kJ per mol CH4
Rxn: CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) +heat
Solution:
Molar mass of CH4 = g / mol
5.8 g CH4 1 mol CH4 = 0.36 mol CH4
16.04 g CH4
0.36 mol CH kJ = -320 kJ
1 mol CH4

8. Solution: Molar mass of SO2 = 64.07 g / mol
Calculate H when 12.8 g of sulphur dioxide reacts with excess oxygen to form sulphur trioxide. H = kJ per mol SO2
Solution:
Molar mass of SO2 = g / mol
12.8 g SO2 1 mol SO2 = mol SO2
64.07 g SO2
mol SO kj = kJ
1 mol SO2

9. Hydrogen peroxide decomposes according to the following thermochemical reaction:
H2O2(l) → H2O(l) + 1/2 O2(g)
ΔH = kJ
Calculate the change in enthalpy (ΔH) when 4.00 grams of hydrogen peroxide decomposes.
Calculate the change in enthalpy (ΔH) when grams of water are created…

10. The reaction that occurs in hand warmers is:
4Fe(s) + 3O2(g)  2Fe2O3(s) + heat
H of reaction = -1652kJ
How much heat is released when 2.5 grams of Fe(s) is reacted with excess oxygen?

11. Law of Concert Tickets!
Miller Stitts  LadyGaga $600
2 Winters + Miller  Justin Bieber + $300
Stitt  2 Elvis Winter $500
Lady Gaga  Justin Bieber Elvis $_____ ???

12. Law of Concert Tickets = Hess’s Law
You can add KNOWN equations to solve UNKOWN equations
Two Rules:
FLIP (reverse) the equation – you must FLIP the sign (- +)
MULTIPY / DIVIDE equation – must Multiply / Divide the ΔH!

13. Example #1: Calculate the enthalpy for this reaction:
2C(s) + H2(g) ---> C2H2(g)
ΔH° = ??? kJ
Given the following thermo chemical equations:
C2H2(g) + (5/2)O2(g) ---> 2CO2(g) + H2O(l)
ΔH° = kJ
C(s) + O2(g) ---> CO2(g)
ΔH° = kJ
H2(g) + (1/2)O2(g) ---> H2O(l)
ΔH° = kJ
a) first eq: flip it so as to put C2H2 on the product side  b) second eq: multiply it by two to get 2C  c) third eq: do nothing. 
kJ + (-787 kJ) + ( kJ) = kJ

14. 2C(s) + H2(g) ---> C2H2(g)
ΔH° = ??? kJ
C2H2(g) + (5/2)O2(g) ---> 2CO2(g) + H2O(l)
ΔH° = kJ
C(s) + O2(g) ---> CO2(g)
ΔH° = kJ
H2(g) + (1/2)O2(g) ---> H2O(l)
ΔH° = kJ

15. +234 + (+592) + (-1220) = -394 Example #2: Given the following data:
Find the ΔH of the following reaction:
C(s) + O2(g) ---> CO2(g)
SrO(s) + CO2(g) ---> SrCO3(s)
ΔH = -234 kJ
2SrO(s) ---> 2Sr(s) + O2(g)
ΔH = kJ
2SrCO3(s) ---> 2Sr(s) + 2C(s) + 3O2(g)
ΔH = kJ
a) first equation - flip it
b) second equation - divide by two
c) third equation - flip it, divide by two
(+592) + (-1220) = -394

16. +234 + (+592) + (-1220) = -394 Example #2: Given the following data:
Find the ΔH of the following reaction:
C(s) + O2(g) ---> CO2(g)
SrO(s) + CO2(g) ---> SrCO3(s)
ΔH = -234 kJ
2SrO(s) ---> 2Sr(s) + O2(g)
ΔH = kJ
2SrCO3(s) ---> 2Sr(s) + 2C(s) + 3O2(g)
ΔH = kJ
a) first equation - flip it
b) second equation - divide by two
c) third equation - flip it, divide by two
(+592) + (-1220) = -394

17. Hess's Law
In a reaction, the change in enthalpy (ΔH)
is the same - regardless if the reaction occurs in a single step or in several steps.
If a series of reactions are added together, the net change in ΔH is the sum of the enthalpy changes for each step.

18. Rules for using Hess's Law
If the reaction is multiplied (or divided) by some factor, Δ H must also be multiplied (or divided) by that same factor.
If the reaction is reversed (flipped), the sign of Δ H must also be reversed.

19. Water phase changes & Energy
constant
Temperature remains __________ during a phase change.

20. Phase Change Diagram Processes occur by addition of energy 
 Processes occur by removal of energy

21. Phase Diagram
Represents phases as a function of temperature and pressure.
Critical temperature: temperature above which the vapor can not be liquefied.
Critical pressure: pressure required to liquefy AT the critical temperature.
Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

22. Phase Changes

23. Effect of Pressure on Boiling Point
Boiling Point of Water at Various Locations
Location
Feet above sea level
Patm (kPa)
Boiling Point (C)
Top of Mt. Everest, Tibet
29,028 32 70
Top of Mt. Denali, Alaska
20,320
45.3 79
Top of Mt. Whitney, California
14,494
57.3 85
Leadville, Colorado
10,150 68 89
Top of Mt. Washington, N.H.
6,293
78.6 93
Boulder, Colorado
5,430
81.3 94
Madison, Wisconsin
900
97.3 99
New York City, New York 10
101.3
100
Death Valley, California
-282
102.6
100.3

24. Phase Diagram
Represents phases as a function of temperature and pressure.
Critical temperature: temperature above which the vapor can not be liquefied.
Critical pressure: pressure required to liquefy AT the critical temperature.
Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

25. Phase changes by Name

26. Water

27. Phase Diagram for Carbon
Carbon dioxide
Phase Diagram for Carbon
dioxide

28. Phase Diagram for Carbon

29. Phase Diagram for Sulfur

30. Reaction Pathway
Shows the change in energy during a chemical reaction

31. Exothermic Reaction 2H2(l) + O2(l)  2H2O(g) + energy
reaction that releases energy
products have lower PE than reactants
energy released
2H2(l) + O2(l)  2H2O(g) + energy

32. Endothermic Reaction 2Al2O3 + energy  4Al + 3O2
reaction that absorbs energy
reactants have lower PE than products
energy absorbed
2Al2O3 + energy  4Al + 3O2

33. Latent Heat of Phase Change
Molar Heat of Fusion
The energy that must be absorbed in order to convert one mole of solid to liquid at its melting point.
Molar Heat of Solidification
The energy that must be removed in order to convert one mole of liquid to solid at its freezing point.

34. Latent Heat of Phase Change #2
Molar Heat of Vaporization
The energy that must be absorbed in order to convert one mole of liquid to gas at its boiling point.
Molar Heat of Condensation
The energy that must be removed in order to convert one mole of gas to liquid at its condensation point.

35. Latent Heat – Sample Problem
Problem: The molar heat of fusion of water is kJ/mol. How much energy is needed to convert 60 grams of ice at 0C to liquid water at 0C?
Mass
of ice
Molar
Mass of
water
Heat of
fusion

36. Heat of Solution
The Heat of Solution is the amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent.
Substance
Heat of Solution
(kJ/mol)
NaOH
-44.51
NH4NO3
+25.69
KNO3
+34.89
HCl
-74.84

37. Energy is the capacity to do work
Thermal energy is the energy associated with the random motion of atoms and molecules
Chemical energy is the energy stored within the bonds of chemical substances
Nuclear energy is the energy stored within the collection of neutrons and protons in the atom
Electrical energy is the energy associated with the flow of electrons
Potential energy is the energy available by virtue of an object’s position
6.1

38. Thermochemistry is the study of heat change in chemical reactions.
The system is the specific part of the universe that is of interest in the study.
SURROUNDINGS
SYSTEM
open
closed
isolated
Exchange:
mass & energy
energy
nothing
6.2

39. 2H2 (g) + O2 (g) 2H2O (l) + energy energy + 2HgO (s) 2Hg (l) + O2 (g)
Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings.
2H2 (g) + O2 (g) H2O (l) + energy
H2O (g) H2O (l) + energy
Endothermic process is any process in which heat has to be supplied to the system from the surroundings.
energy + H2O (s) H2O (l)
energy + 2HgO (s) Hg (l) + O2 (g)
6.2

40. DH = H (products) – H (reactants)
Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure.
DH = H (products) – H (reactants)
DH = heat given off or absorbed during a reaction at constant pressure
Hproducts < Hreactants
Hproducts > Hreactants
DH < 0
DH > 0
6.3

41. Thermochemical Equations
Is DH negative or positive?
System absorbs heat
Endothermic
DH > 0
6.01 kJ are absorbed for every 1 mole of ice that melts at 00C and 1 atm.
H2O (s) H2O (l)
DH = 6.01 kJ
6.3

42. Thermochemical Equations
Is DH negative or positive?
System gives off heat
Exothermic
DH < 0
890.4 kJ are released for every 1 mole of methane that is combusted at 250C and 1 atm.
CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l)
DH = kJ
6.3

43. Thermochemical Equations
The stoichiometric coefficients always refer to the number of moles of a substance
H2O (s) H2O (l)
DH = 6.01 kJ
If you reverse a reaction, the sign of DH changes
H2O (l) H2O (s)
DH = kJ
If you multiply both sides of the equation by a factor n, then DH must change by the same factor n.
2H2O (s) H2O (l)
DH = 2 x 6.01 = 12.0 kJ
6.3

44. Thermochemical Equations
The physical states of all reactants and products must be specified in thermochemical equations.
H2O (s) H2O (l)
DH = 6.01 kJ
H2O (l) H2O (g)
DH = 44.0 kJ
How much heat is evolved when 266 g of white phosphorus (P4) burn in air?
P4 (s) + 5O2 (g) P4O10 (s) DH = kJ
1 mol P4
123.9 g P4 x
3013 kJ
1 mol P4 x
266 g P4
= 6470 kJ
6.3

45. Heat (q) absorbed or released:
The specific heat (s) of a substance is the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius.
The heat capacity (C) of a substance is the amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius.
C = ms
Heat (q) absorbed or released:
q = msDt
q = CDt
Dt = tfinal - tinitial
6.4

46. Dt = tfinal – tinitial = 50C – 940C = -890C
How much heat is given off when an 869 g iron bar cools from 940C to 50C?
s of Fe = J/g • 0C
Dt = tfinal – tinitial = 50C – 940C = -890C
q = msDt
= 869 g x J/g • 0C x –890C
= -34,000 J
6.4

47. Constant-Volume Calorimetry
qsys = qwater + qbomb + qrxn
qsys = 0
qrxn = - (qwater + qbomb)
qwater = msDt
qbomb = CbombDt
Reaction at Constant V
DH = qrxn
DH ~ qrxn
No heat enters or leaves!
6.4

48. Constant-Pressure Calorimetry
qsys = qwater + qcal + qrxn
qsys = 0
qrxn = - (qwater + qcal)
qwater = msDt
qcal = CcalDt
Reaction at Constant P
DH = qrxn
No heat enters or leaves!
6.4

49. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Terminology
Energy, U
The capacity to do work.
Work
Force acting through a distance.
Kinetic Energy
The energy of motion.
Energy is from the Greek “work within”.
Moving objects do work when they slow down or are stopped.
Kinetic means “motion” in greek.
Prentice-Hall © 2002
General Chemistry: Chapter 7

50. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Energy
Kinetic Energy 1
kg m2
ek =
mv2
[ek ] =
= J 2 s2
w = Fd
[w ] =
kg m s2
= J m
Work
[w] means UNITS OF WORK, not concentration in this case.
Prentice-Hall © 2002
General Chemistry: Chapter 7

51. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Energy
Potential Energy
Energy due to condition, position, or composition.
Associated with forces of attraction or repulsion between objects.
Energy can change from potential to kinetic.
Energy changes continuously from potential to kinetic.
Energy is lost to the surroundings.
Prentice-Hall © 2002
General Chemistry: Chapter 7

52. Energy and Temperature
Thermal Energy
Kinetic energy associated with random molecular motion.
In general proportional to temperature.
An intensive property.
Heat and Work
q and w.
Energy changes.
Prentice-Hall © 2002
General Chemistry: Chapter 7

53. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Heat
Energy transferred between a system and its surroundings as a result of a temperature difference.
Heat flows from hotter to colder.
Temperature may change.
Phase may change (an isothermal process).
Heat is transfer of energy. Bodies do NOT contain heat.
Prentice-Hall © 2002
General Chemistry: Chapter 7

54. General Chemistry: Chapter 7
Units of Heat
Calorie (cal)
The quantity of heat required to change the temperature of one gram of water by one degree Celsius.
Joule (J)
SI unit for heat
1 cal = J
Prentice-Hall © 2002
General Chemistry: Chapter 7

55. General Chemistry: Chapter 7
Heat Capacity
The quantity of heat required to change the temperature of a system by one degree.
Molar heat capacity.
System is one mole of substance.
Specific heat capacity, c.
System is one gram of substance
Heat capacity
Mass  specific heat.
q =
mcT
q =
CT
Prentice-Hall © 2002
General Chemistry: Chapter 7