1. Thermochemistry
Chapter 5

2. Thermochemistry Thermochemistry:
Relationships between chemical reactions and energy changes
Thermodynamics:
Therm: heat Dynamics: power

3. Energy Radiant energy Thermal energy Chemical energy Nuclear energy
Energy: the capacity to do work
Radiant energy
Thermal energy
Chemical energy
Nuclear energy
Potential energy- relative to position- stored in bonds
electrostatic potential energy (C) = kQ1Q2 d
Kinetic energy Ek = ½ mv2
Units of Energy: J = 1kg-m2/s2 1 cal= J

4. Energy Changes in Chemical Reactions
Which one has more thermal energy?
900C
900C
Coffee cup?
Bathtub?

5. Energy Changes in Chemical Reactions
Which one has more thermal energy?
400C
900C
What additional information is needed?
How does it relate to heat?

6. Energy Changes in Chemical Reactions
Heat:
the transfer of _______ between two bodies that are at ________temperatures.
Temperature is a measure of the ________.
Measure of the ______________
Temperature = Thermal Energy

7. Energy Changes SURROUNDINGS SYSTEM
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

8. Energy Changes

9. Enthalpy
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 is supplied to the system from the surroundings.
energy + H2O (s) H2O (l)
energy + 2HgO (s) Hg (l) + O2 (g)

10. ΔHreaction = ΔHproducts - ΔHreactants
Enthalpy
Enthalpy: a thermodynamic quantity used to describe heat changes(absorbed/released) at constant pressure (most reactions occur at constant pressure).
H = E + PV
Equation:
ΔHreaction = ΔHproducts - ΔHreactants

11. Enthalpy: Enthalpy Diagrams
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

12. Enthalpy: Enthalpy Diagrams
ΔHrxn= ΔHproducts – ΔHreactants

13. Enthalpy: Examples Indicate the sign of enthalpy change in the
following processes carried under atmospheric
pressure, and indicate whether the process is
exothermic or endothermic:
An ice cube melts
1 g of butane (C4H10) is combusted in sufficient oxygen to give complete combustion to CO2 and H2O
A bowling ball is dropped from a height of 8 ft into a bucket of sand.

14. 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

15. 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

16. Thermochemical Equations: Rules
Stoichiometric coefficients: equal 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 = 12.0 kJ
2 x 6.01kJ/mol =12.0 KJ/2 mol

17. Thermochemical Equations: Rules
When a reaction is reversed, the magnitude of ΔH remains the same, but its sign changes.

18. Thermochemical Equations: Rules
The physical states of all reactants and products must be specified in thermochemical equations.
H2O(s) → H2O(l) ΔH = 6.00 kJ
H2O(l) → H2O(g) ΔH = 44.0 kJ
What will be the ΔH when 1 mole of ice at 0ºC is changed into one mole of steam at the boiling point?

19. Thermochemical Equations: Rules
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
[ kJ]

20. Enthalpy: Examples
4. The complete combustion of liquid octane, C8H18, to produce carbon dioxide and liquid water at 25 ºC and at constant pressure gives off 47.9 kJ per gram of octane. Write a chemical equation to represent this information.
[2C8H18(l) O2(g) → 16 CO2(g) + 9H2O(l)
ΔH = x 104 kJ]

21. Enthalpy: Examples SO2(g) + ½ O2(g) →SO3(g) ΔH = -99.1 kCal
5. Given the thermochemical equation:
SO2(g) + ½ O2(g) →SO3(g) ΔH = kCal
Calculate the heat evolved when 74.6 g of SO2
(MM = g/mol) is converted to SO3.
[ -115 kJ]
6. Hydrogen peroxide can decompose to water and oxygen by the following reaction:
2H2O2(l) → 2H2O(l) + O2(g) ΔH = -196 kJ
Calculate the value of q (heat evolved or absorbed) when 5.00 g of H2O(l) decomposes at constant pressure.
[ kJ]

22. Enthalpy: Examples 7. Given the equation
H2(g) + I2(s) → 2HI(g) ΔH = KJ
Calculate ΔH for the reaction
HI(g) → ½ H2(g) + ½ I2(s)
[ kJ]
8. Given the equation:
3O2(g) → 2O3(g) ΔH = kJ
Calculate ΔH for the following reaction:
3/2 O2(g) → O3(g)
[ kJ]

23. Calorimetry and Heat Exchanges in a Chemical Reaction
Read Section 6.4 (Pages )
Examples: pp. 216 – 220: 6.1 – 6.3
Exercises: pp. 240: 6.19, 6.23, 6.26, 6.28

24. Calorimetry Calorimetry is performed in devices called calorimeters
Calorimetry is based on the Law of Conservation of Energy

25. Calorimetry
9. The specific heat of aluminum is J/g ºC. Calculate the heat necessary to raise the temperature of 40.0 g of aluminum from ºC to 32.3 ºC. Specific heat of aluminum is cal/g ºC.
[1.87 x 103 Joules]

26. Calorimetry: Heat Capacity
Calorimetry: Experimental measurement of heat produced in chemical and physical processes
Heat Capacity, C, of a system:
the quantity of heat needed to raise the temperature of matter 1 K.
Units: J/ºC
Molar Heat Capacity:
The heat required to raise 1 mole of a substance, 1 K.

27. Calorimetry: Specific Heat
Specific heat of a substance
(also represented by cp):
the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius ( or 1 K).

28. Calorimetry: Definitions
Heat (q) absorbed or released:
q = mcpDt
q = CDt
Dt = tfinal - tinitial
When q>0 endothermic reaction
When q<0 exothermic reaction

29. Calorimetry: Specific Heat
10. How much heat, in joules and kilojoules, does it take to raise the temperature of 225 g of water from 25.0 ºC to ºC
[ 7.05 x 104 J; kJ]
11. Calculate the heat capacity, specific heat and the molar heat capacity of an aluminum block that has mass of 5.7 g that must absorb 629 J of heat from its surroundings for its temperature to rise from 22 ºC to 145 ºC.
[C = 5.11 J/ºC; cp = 0.90 J g-1 ºC-1, 24 J mol-1 ºC-1 ]

30. Calorimetry: Problems
12. What will be the final temperature if a 5.00-g silver ring at 37.0 ºC gives off 25.0 J of heat to the surroundings? Specific heat of silver = J g-1 ºC-1
[15.7 ºC]
13. A 15.5 g sample of a metallic alloy is heated to 98.9 ºC and then dropped into 25.0 g of water in a calorimeter. The temperature of the water rises from 22.5 ºC to 25.7 ºC. Calculate the specific heat of the alloy.
[0.29 J g-1 ºC-1]

31. Calorimetry: Problems
14. Suppose a piece of gold with a mass of 21.5 g at temperature of ºC is dropped into an insulated calorimeter containing g of water at ºC. What will be the final temperature of the water?
[22.4 ºC]

32. Measuring Enthalpy Changes for Chemical Reactions: Constant-Volume Calorimetry
Bomb Calorimeter

33. Constant Volume Calorimeter: Isolated System
No exchange of matter and energy
Uses Cp of water

34. Constant-Volume Calorimetry: Isolated system
qsys = qwater + qbomb + qrxn = 0
There is no heat exchange with the surroundings
qrxn = - (qwater + qbomb)
Cbomb= mbombx cbomb
qbomb = CbombDt qwater = mwcpΔt
qrxn = -(qwater + CbombΔt) = -(mwcpΔt + CbombΔt)
qrxn = - (mwcp + Cbomb)Δt
Measured heat is not enthalpy: constant volume

35. Constant Volume Calorimetry: Problems
15. In a preliminary experiment, the heat capacity of a bomb calorimeter assembly is found to be 5.15 kJ/°C. In a second experiment. A g sample of graphite (carbon) is placed in the bomb with an excess of oxygen. The water, bomb, and other contents of the calorimeter are in thermal equilibrium at °C. The graphite is ignited and burned, and the water temperature rises to °C. Calculate ΔH for the reaction.
C(graphite) + O2(g) → CO2(g)
[ΔH = -393 kJ]

36. Measuring Enthalpy Changes for Chemical Reactions
Coffee-Cup Calorimeter: Constant pressure

37. Measured heat is the enthalpy
Constant-Pressure Calorimetry
qsys = qwater + qcal + qrxn
qsys = 0
qrxn = - (qwater + qcalor)
qwater = mcpDt
qcal = CcalDt
Reaction at Constant P
ΔH = qrxn = - (qwater + qcalor)
No heat enters or leaves!
Insulated system!
Pressure is constant.
Measured heat is the enthalpy

38. Calorimetry

39. Constant Pressure Calorimetry: Problems
(a) A 50.0 mL sample of M HCl at °C is added to 50.0 mL of M NaOH, also at °C, in a calorimeter. After mixing, the solution temperature rises to °C. Calculate the heat of this reaction. Assume te volumes of the solutions are additive. Ignore the heat absorbed by the calorimeter.
[-715 J]
(b) Determine the value of ΔH that should be written in the equation for the neutralization reaction:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) ΔH = ?
[-57.2 kJ]

40. Calorimetry: Problem
17. A quantity of 1.00 x 102 mL of 0.500M HCl is mixed with 1.00 x 102 mL of M NaOH in a constant pressure calorimeter having a heat capacity of 335 J/ ºC. The initial temperature of the HCl and NaOH solutions is the same, ºC, and the final temperature of the mixed solution is ºC. Calculate the heat change for the neutralization reaction and the molar heat of the neutralization reaction. Assume s = 4.18 J/gºC for the solution.
[-2.81 kJ; kJ/mol]

41. Fuel Values of Foods and Other Substances
Chemistry in Action:
Fuel Values of Foods and Other Substances
C6H12O6 (s) + 6O2 (g) CO2 (g) + 6H2O (l) DH = kJ/mol
1 cal = J
One food Calorie:
1 Cal = 1000 cal = 4184 J

42. Standard Enthalpy of Formation
Absolute value of the enthalpy of a substance cannot be measured. We need a frame of reference.
Standard states for elements and compounds has to be established.
Standard state for a solid or liquid substance is the pure element or compound at 1 atm at a given temperature.
Standard state for gas: pure gas behaving as an ideal gas at 1 atm and temperature of interest.
Standard enthalpy of formation: the enthalpy change for a reaction in which the reactants in their standard state yield products in their standard states. Standard state denoted with a superscript (º)

43. Standard Enthalpy of Formation
Standard enthalpy of formation ΔHºf (also called heat of formation
enthalpy change that occurs in the formation of 1 mole of the substance from its elements when both products and reactants are in their standard states (f stands for formation).
Reference forms:
standard enthalpy of formation of a pure element in its reference form is 0.
standard enthalpy of formation (ΔH0) as a reference point for all enthalpy expressions is the formation of H2(g) = 0 kJ/mol

44. Standard Enthalpy of Formation
Standard enthalpy of formation (DH0f) is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm and 25 ºC .
The standard enthalpy of formation of any element in its most stable form is zero.
DH0 (O2) = 0 f
DH0 (C, graphite) = 0 f
DH0 (C, diamond) = 1.90 kJ/mol f
DH0 (O3) = 142 kJ/mol f

45. standard

46. Determination of Heat of Formation
1. Direct method: From elements when it can be easily measured.
C(graphite) + O2(g) →CO2(g) ΔHrxn = kJ
ΔHrxn= ΔHoCO2 – ΔH0O2 – ΔHgraphite = kJ
ΔH0CO2 = kJ
2. Indirect method: Using the heat of reaction and heats of formations
Given CH4(g) +2O2(g) → CO2(g) + 2H2O(g) ΔHrxn= kJ
Calculate the heat of formation of CH4
[-74.8kJ/mol]

47. Standard Enthalpy of Reaction
The standard enthalpy of reaction (DH0 ) is the enthalpy of a reaction carried out at 1 atm and 25ºC.
rxn
aA + bB cC + dD
DH0
rxn
dDH0 (D) f
cDH0 (C) = [ + ] -
bDH0 (B)
aDH0 (A)
DH0
rxn
nDH0 (products) f = S
mDH0 (reactants) -

48. 2C6H6 (l) + 15O2 (g) 12CO2 (g) + 6H2O (l)
Benzene (C6H6) burns in air to produce carbon dioxide and liquid water. How much heat is released per mole of benzene combusted? The standard enthalpy of formation of benzene is kJ/mol.
2C6H6 (l) + 15O2 (g) CO2 (g) + 6H2O (l)
DH0
rxn
nDH0 (products) f = S
mDH0 (reactants) -
DH0
rxn
6DH0 (H2O) f
12DH0 (CO2) = [ + ] -
2DH0 (C6H6)
DH0
rxn =
[ 12x– x–187.6 ] – [ 2x49.04 ] = kJ
-5946 kJ
2 mol
= kJ/mol C6H6

49. Standard Enthalpy of Reaction
Synthesis gas is a mixture of carbon monoxide and hydrogen that is used to synthesize a variety of organic compounds, such as methanol. One reaction for producing synthesis gas is shown here:
3CH4(g) + 2H2O(l) + CO2(g) → 4CO(g) + 8H2(g) ΔHº = ?
[747.5 kJ]
The combustion of isopropyl alcohol, common rubbing alcohol, is represented by this equation:
2(CH3)2CHOH(l) + 9O2(g) → 6CO2(g) + 8H2O(l)
ΔHº = kJ
use this equation and data from Thermodynamic tables to establish the standard enthalpy of formation for isopropyl alcohol.
[ -318 kJ/mol]

50. Heat of Formation: Examples
20. Pentaborane-9, B5H9, is a colorless liquid. It is highly reactive substance that will burst into flame or even explode when exposed to oxygen:
2B5H9(l) + 12O2 → 5B2O3(s) + 9H2O(l)
It was considered a potential rocket fuel as it release large amount of heat per gram. Calculate the kJ of heat released per gram of the compound reacted with oxygen. The standard enthalpy of formation of B5H9 is 73.2 kJ/mol
[ kJ/gB5H9]

51. Heat of Reactions (Formations): Hess’s Law of Constant heat of Summations
Hess’s Law: When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.
(Enthalpy is a state function. It doesn’t matter how you get there, only where you start and end.)

52. Hess’ Law: Example Given the following reactions:
CO(g) + ½ O2(g) → CO2(g) ΔH = kJ
C(graphite) + O2(g) → CO2(g) ΔH = kJ
Calculate the enthalpy for the formation of CO from its elements
C(graphite) + ½ O2(g) → CO(g) ΔH = ?

53. Hess’s Law: Example

54. Calculate the standard enthalpy of formation of CS2 (l) given that:
C(graphite) + O2 (g) CO2 (g) DH0 = kJ
rxn
S(rhombic) + O2 (g) SO2 (g) DH0 = kJ
rxn
CS2(l) + 3O2 (g) CO2 (g) + 2SO2 (g) DH0 = kJ
rxn
1. Write the enthalpy of formation reaction for CS2
C(graphite) + 2S(rhombic) CS2 (l)
2. Add the given rxns so that the result is the desired rxn.
rxn
C(graphite) + O2 (g) CO2 (g) DH0 = kJ
2S(rhombic) + 2O2 (g) SO2 (g) DH0 = x2 kJ
rxn +
CO2(g) + 2SO2 (g) CS2 (l) + 3O2 (g) DH0 = kJ
rxn
C(graphite) + 2S(rhombic) CS2 (l)
DH0 = (-296.1) = 86.3 kJ
rxn
6.6

55. Hess’ Law: Examples 21. Given
C(graphite) + O2 (g)→ CO2(g) ΔHºrxn= kJ
H2(g) + ½ O2(g) → H2O(l) ΔHºrxn = kJ
2C2H2(g) + 5O2(g) → CO2(g) + 2H2O(l) ΔHºrxn = kJ
Calculate the heat of formation of C2H2(g)
2C(graphite) + H2(g) → C2H2(g)
[52.3 kJ]

56. Hess’ Law: Conceptual Example
Without performing a
calculation, determine
which of the two
substances should
yield the greater
quantity of heat upon
Complete combustion,
on a per mole basis:
ethane, C2H6(g), or
ethanol, CH3CH2OH(l)

57. The Solution Process for NaCl
DHsoln = Step 1 + Step 2 = 788kJ – 784kJ = 4 kJ/mol

58. The enthalpy of solution (DHsoln) is the heat generated or absorbed when a certain amount of solute dissolves in a certain amount of solvent.
DHsoln = Hfinal soln - Hcomponents
Which substance(s) could be used for melting ice?
Which substance(s) could be used for a cold pack?
Which substances could be used for hot pack?

59. Heat of Solution and Dilution
The process: NaCl(s) → Na+(aq) + Cl-(aq)
The steps:
1. Separate ions in the crystal into ions in gaseous state
2. Hydration of the gaseous ions
Energy involved:
Lattice Energy, U
Hydration energy, ΔH hydration
ΔH solution= U + ΔH hydration

60. Heat of Solution and Dilution
Lattice energy: energy required to completely separate one mole of solid ionic compound into gaseous ions (endothermic process)
U + NaCl(s) →Na+(g) + Cl-(g)
Hydration of the gaseous ions
Heat of hydration: energy released when ions interact with water molecules.
Na+(g) + Cl-(g) → Na+(aq) + Cl-(aq) + Energy

61. Heat of Dilution The heat change associated with the
dilution of a solution.
Can be exothermic or endothermic

62. The End