# Energy Part III: Calculation of ?H from a) Thermochemical Equations b) Heat of Formation Chapter 7 Sec 6 - PowerPoint PPT Presentation

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## Energy Part III: Calculation of ?H from a) Thermochemical Equations b) Heat of Formation Chapter 7 Sec 6

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### Energy Part III: Calculation of H from a) Thermochemical Equations b) Heat of Formation Chapter 7 Sec 6 Sec 8 of Jespersen 6TH ed Dr. C. Yau Spring 2013 – PowerPoint PPT presentation

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Title: Energy Part III: Calculation of ?H from a) Thermochemical Equations b) Heat of Formation Chapter 7 Sec 6

1
Energy Part III Calculation of ?H from a)
Thermochemical Equations b) Heat of
Formation Chapter 7 Sec 6 Sec 8 of Jespersen
6TH ed
• Dr. C. Yau
• Spring 2013

1
2
Standard Heat of Reaction
?H enthalpy of reaction or heat of
reaction heat transferred in a rxn (usually
in kJ, not kJ/mol) ?Ho standard heat of
reaction ?H at standard conditions
o means standard conditions (1 atm, 1M if
aq soln, usually 25oC) Remember these conditions!
2
3
Thermochemical Equations
N2 (g) 3H2 (g) 2NH3 (g) ?Ho
-92.39 kJ This tells us that 1 mole of N2 would
produce 92.39 kJ of heat, that 3 moles of H2
would produce 92.39 kJ of heat, that production
of 2 moles of NH3 would be accompanied by a
release of 92.39 kJ of heat. NOTE that ?Ho is in
kJ and not kJ/mol. The amt of heat transferred is
directly proportional to the moles in
thermochemical eqn shown.
3
4
Thermochemical Equations
• Example Magnesium burns in air to produce a
bright light and is often used in fireworks
displays.
• 2 Mg (s) O2 (g) 2MgO (s) ?Ho -1203 kJ
• How many grams of Mg is needed to produce 400.
kJ of heat?
• p.297 7.62
• How much heat (in kJ) is liberated by the
combustion of 6.54 g of Mg?
• Set these problems up in dimensional analysis.

5
Thermochemical Equations
N2 (g) 3H2 (g) 2NH3 (g) ?Ho
-92.39 kJ The value of ?Ho depends on the
coefficients in the equation. If coefficients are
doubled, ?Ho would be doubled 2N2 (g) 6H2
(g) 4NH3 (g) ?Ho -92.39x2 kJ Note
also that ?Ho is dependent on the physical states
as stated in the equation. CH4 (g) 2O2 (g)
CO2(g) 2H2O(l) ?Ho - 890.5kJ CH4 (g) 2O2
(g) CO2(g) 2H2O(g) ?Ho - 802.3kJ
5
6
Example 7.7 p.276
• The following thermochemical equation is for the
exothermic reaction of hydrogen and oxygen that
produces water.
• 2H2 (g) O2 (g) 2H2O (l) ?Ho
-571.8kJ
• What is the thermochemical equation for this rxn
when it is conducted to produce 1.000 mol H2O?
• Do Pract Exer 9 10 p.277

7
Thermochemical Equations
• If we reverse a reaction, the magnitude of ?Ho is
the same but the sign is changed
• C (s) O2 (g) ?? CO2 (g) ?H ? 393.5 kJ
• CO2 (g) ?? C (s) O2 (g) ?H 393.5 kJ

8
• Determination of ?H by manipulation of Eqns.
• Use the two equations below to determine the
standard enthalpy change for the reaction
• H2O2 (l) ?? H2O (l) ½ O2 (g)
• (1) H2 (g) O2 (g) ?? H2O2 (l) ?H ?188 kJ
• (2) H2 (g) ½ O2 (g) ?? H2O (l) ?H ?286 kJ
• Strategy
• We need H2O2 on the left side, so Eqn 1 must be
reversed.
• Eqn 2 probably can stay as is. WE MUST CHECK.
How?
• On my exams you are expected to show your work
as we are doing here in class. Take notes!

9
Determination of ?H by manipulation of Eqns.
• Ethylene glycol, HOCH2CH2OH, is used as
antifreeze. It is produced from ethylene oxide,
C2H4O, by the reaction
• (1) C2H4O (g) H2O (l) ?? HOCH2CH2OH (l)
• What is the heat of reaction of this reaction...
• Given
• (2) 2C2H4O (g) 5O2 (g) ?? 4CO2 (g) 4H2O (l)
• ?H ?2612.2 kJ
• (3) HOCH2CH2OH(l) 5/2 O2(g) ??2CO2(g) 3H2O(l)
• ?H ?1189.8 Kj
• What is the strategy?
• Do Pract Exer 13, 14, 15 p.283

10
• Determination of ?H by manipulation of Eqns.
• 2Cu (s) O2 (g) ?? 2CuO (s) ?H ?310 kJ
• 2Cu (s) ½ O2 (g) ?? Cu2O (s) ?H ?169 kJ
• Use the two equations above to determine the ?H
of this reaction
• Cu2O (s) ½ O2 (g) ?? 2CuO (s)
• Is this exothermic or endothermic?
• Solve the problem by manipulating the given eqns.
• ANS -141 kJ

11
Determination of ?H by manipulation of Eqns.
• Example 7.9 p.282
• Fe2O3 (s) 3CO (g) ?? 2Fe(s) 3CO2(g) ?H ?
26.7 kJ
• CO(g) ½ O2 (g) ?? CO2 (g)
?H ? 283.0 kJ
• Calculate the value of ?H for the following
reaction
• 2 Fe (s) ?O2 (g) ?? Fe2O3 (s)

12
Enthalpy Diagrams
• C (s) ½ O2 (g) ?? CO (g) ?H ?110.5 kJ
• Construct an enthalpy diagram for the reaction.
• Learn the terminology, "enthalpy diagram."
• Know what is asked for on an exam.

13
Enthalpy Diagrams
• N2 (g) O2 (g) ?? 2 NO (g) ?H 181 kJ
• Draw the enthalpy diagram for this reaction.
• We are skipping Example 7.8, Pract Exer 11 12.
You do not need to know how to draw enthalpy
diagrams of that sort. However, you should know
how to add equations together to determine the
enthlapy change for the overall reaction.

14
Hess' Law
• The value of ?H for any reaction that can be
written in steps equals the sum of the values of
?H of each of the individual steps.
• This is based on the fact that enthalpy is a
state function.
• The implication is that regardless of how many
steps are taken the overall enthalpy change is
the same.

15
Enthalpy as a State Function
• Enthalpy (H) depends only on its current state
and not the path taken to get there.
• It is not affected by how many steps are used to
get there.
• A ????? B
• C F
• D E

?H1 ?H2 ?H3 ?H4?H5 ?H6
16
Hess Law is extremely useful because if ?H1
cannot be measured, it can be calculated from ?H2
and ?H3.
• For example
• A B C
• D E

?H1?
?H3
?H2
?H1 ?H2 ?H3
17
Hess' Law
• Watch out for the directions of the arrows.
• What is ?H1 in terms of ?H2 and ?H3?
• F G J
• K L

?H1
?H2
?H3
?H1 ?H2 - ?H3
18
Standard Heat of Combustion
• ?Hoc standard heat of combustion
• It is the amount of heat released when one mole
of a fuel substance is completely burned in pure
oxygen gas with all reactants and products
brought to 25oC and 1 bar pressure (1 atm).
• Combustion reactions are always exothermic.
Therefore, ?Hoc is always negative.

19
• Example 7.10 p.283
• How many moles of carbon dioxide gas are produced
by a gas-fired power plant for every 1.00 MJ
(megajoule) of energy it produces? The plant
burns methane, CH4 (g), for which ?Hoc is -890
kJ/mol
• Do Pract Exer 16 17 p.235

20
Standard Enthalpy of Formation
• LEARN THIS DEFINITION!
• ?Hfo standard enthalpy of formation
• It is the amount of heat absorbed or evolved when
specifically one mole of substance is formed at
25oC at 1 atm from its elements in their standard
states.

21
• ?Hfo for solid potassium sulfate is -1433.7 kJ.
Write the thermochemical equation corresponding
to this value.
• ?Hfo for solid ammonium chloride is 315.4 kJ.
Write the thermochemical equation corresponding
to this value.
• ?Hfo for solid calcium hydroxide is -986.6 kJ.
Write the thermochemical equation corresponding
to this value.
• Do Example 7.11 and Pract Exer 18 19 on p.286

22
?Hfo of Elements
• Write the thermochemical equation corresponding
to the ?Hfo of chlorine gas.
• What do you think the value of ?Hfo would be for
chlorine gas?
• What about the value of ?Hfo of solid silver? of
liquid mercury?
• Remember this! You will not be provided ?Hfo
elements in their standard states. They are
always ZERO kJ.

23
Applying Hess' Law to Heats of Formation
• A B C
D

All elements in their standard states.
What is heat of reaction of AB CD?
24
Hess' Law of Summation
• ?Hfo values are given in Table 7.2 p. 285 and
will be provided on exams.

25
Example 7.12 p.287
• Some chefs keep baking soda, NaHCO3, handy to put
out grease fires. When thrown on the fire, baking
soda partly smothers the fire and the heat
decomposes it to give CO2, which further smothers
the flame. The eqn is
• 2NaHCO3 (s) ?? Na2CO3 (s) H2O (l) CO2 (g)
• Use the data in Table 7.2 (p.285) to calc the ?Ho
for this reaction in kilojoules.
• Do not put out a kitchen fire with water!
related

26
Calculating ?H For Reactions Using ?Hf
• 2Fe(s) 6H2O(l) ? 2Fe(OH)3(s)
3H2(g)
• ?Hf -285.8 -696.5
• kJ mol-1
• CO2(g) 2H2O(l) ? 2O2(g) CH4(g)
• ?Hf -393.5 -285.8 -74.8
• kJ mol-1
• Do Pract Exer 20, 21 22 p. 288

27
SUMMARY
• What are the different ways you can determine the
?H of a reaction?
• Measure q from calorimetry experiment. In open
containers, q ?H
• Calculate by manipulating given thermochemical
equations.
• Calculate from ?Hfo
• Calculate from bond energies (handout)
• Also, remember how to calculate amt of heat from
stoichiometry thermochemical eqn.