Chapter 5
Thermochemistry
S
System and Surroundings
 System
 In which the chemical
reaction is taking place
 Surroundings
 All else (container, air
outside of the container,
etc.)
Universe= System +
Surroundings
Open system and closed system
 Open System: the transfer of matter and
energy is possible across its boundary
 Closed system: allow no transfer of matter,
though energy may be transferred across the
boundary.
 Isolated system: No matter and Energy change.

Label each system
as open or closed,
Explain

Which system
would you use to
demonstrate
conservation of
mass? Why?

Why do you think
the fish bowl
above is
considered a
system, but an
empty fish bowl is
not?
The First Law of Thermodynamics
Law of Conservation of Energy
First Law of Thermodynamics: Energy
cannot be created or destroyed  energy
is conserved.
Important!!
 Energy lost by a system is gained by the
surroundings (and vice versa)
 Energy can be transferred but can not be
created or destroyed
The First Law of Thermodynamics
Law of Conservation of Energy
First Law of Thermodynamics: Energy
cannot be created or destroyed  energy is
Important!!
conserved.
 The overall change in energy of a system
will be equal to:
The final energy of the system minus the
initial energy of the system
E = Efinal − Einitial
Enthalpy Change(system energy)
Enthalpy change – The energy difference between the final state
and initial state. It is a state function,any change in value is independent of the
pathway between the inial and fianl state instead of the passway.
H = Efinal − Einitial
What are the
units of
Enthalpy?
The process is
exothermic
The process is
endothermic
Question
 Are the two following processes endothermic or
exothermic?
 An ice cube melts in a reaction flask (ice cube is
the system)
 Combustion of propane in a reaction flask
(propane and oxygen are the system, flask and
all else is surrounding)
1.
2.
3.
4.
Exothermic, Endothermic
Exothermic, Exothermic
Endothermic, Exothermic
Endothermic, Endothermic
Fillling the Blank
 During an exothermic reaction the chemical reactants
are _____(losing/gaining) energy. This energy is used
to heat the surroundings – e.g. the _______________
,The products end up with _______(more/less) energy
than the reactants had – but the surroundings end up
with ______(more/less), and get
______(hotter/cooler).
 We measure the energy transferred to and from the
surroundings as enthalpy change, H.
State Functions
 State Function: property of a system that
 is defined by specifying the system’s
condition or state (temperature, pressure,
etc.)
 is not defined by the path by which the
system achieved that state
Enthaphy is an example of a State Function
Question
 An unknown substance that was initially at 25 °C and had



1.
2.
3.
4.
an internal energy of E1 = 1 kJ
It was heated to 50 °C, where its internal energy was E2 =
2 kJ
It was then cooled down in a refrigerator to 4 °C where its
internal energy was E3 = 0.5 kJ.
Determine the Enthalphy change from 4 °C to 50
°C(ΔH).
1 kJ
1.5 kJ
0.5 kJ
- 0.5 kJ
ENERGY
 Energy is the capacity to do work or transfer heat
 SI Unit:
Joule (J)
 The Joule is the unit for all types of energy, so it applies
to:
■
■
■
■
■
■
Kinetic Energy
Potential Energy
Heat
Work
molecular Energy
Enthalpy
WORK Example
Work: Energy used to cause
an object to move against a
force
HEAT Example
Heat: The energy used to
cause the temperature of an
object to increase
 Heat flows from warmer
objects to cooler objects
Enthalpy of Reaction
S
Enthalpy of Reaction
 The change in heat energy of a system
undergoing a chemical reaction at constant
pressure is defined as the enthalpy change, ΔH,
of the process
 Also referred to as “heat change” or “heat
of reaction”, ΔH
 Unit: KJ
Enthalpy of Reaction
Reactants
 For a chemical reaction the
enthalpy of reaction is:
ΔH = Hproducts – Hreactants
 ΔH can be measured
experimentally (Going to
introduce later)
Products
Exothermic
Endothermic
Exothermic reaction: A chemical
reaction that give out heat energy to the
surrounding
ΔH<0
Endothermic reaction: A chemical
reaction that absorb heat energy to the
surrounding
ΔH>0
Compare the stability of reactants and products
 Substance with low energy is more stable.
Write down a Thermal Equation
 State
 Coefficient Number
 Enthalpy ΔH
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
890 kJ
H =
Will the ΔH change with the change of
coefficient number?
Will the ΔH change with the change of state of
reactants or products?
Enthalpy of Reaction---Q1
1.
Enthalpy is an extensive property ( proportional to
the amount of reactants consumed and the amount of
products generated.)
 Value of ΔH will given for a balanced reaction
(thermochemical equation)
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
H = 890 kJ
What would be the enthalpy associated with the combustion of 2
moles of CH4?
What would be the enthalpy associated with the production of 3
moles of 3 moles of CO2?
Enthalpy of Reaction---Q2
2. H for a reaction in the forward direction is
equal in value, but opposite in sign, to H for
the reverse reaction.
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
H = 890 kJ
CO2(g) + 2 H2O(l)  CH4(g) + 2 O2(g)
H = ()(890) kJ
H = + 890 kJ
Enthalpy of Reaction---Q3
3.
The enthalpy change for a reaction depends on the state of the reactants
and the products.
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
H = 890 kJ
If the product was H2O(g) instead of H2O(l), would H be the same?
H = 802 kJ because:
2 H2O(l)  2 H2O(g), H = +88 kJ
Question 4
Calculate the quantity of heat produced when
0.345 mol of O2 is formed at constant pressure.
2 KClO3(s)  2 KCl(s) + 3 O2(g)
H = 89.4 kJ
Question 5
Calculate the quantity of heat produced when
2.50 g of NH4NO3 decomposes at constant
pressure.
NH4NO3(s)  N2O(g) + 2 H2O(g)
H = 37.0 kJ
Some special Enthalpy
 Standard enthalpy change of a reaction:
determined at a stardard condition
( 25℃/298K & 100KPa)
Symbol:
Unit: KJ
 Enthalpy change of formation:the energy
change upon the formation of 1 mol of a
substance from its constituent elements
in their standard state.
Symbol :
Unit: KJ/mol
Enthalpy of Formation, ΔHf
 Also known as “heat of formation”  ΔHf
 Definition: Enthalpy change associated with the
formation of a compound from its constituents
 Example: “formation” of methane (CH4)
 Constituents = C and H
 C + 2H2  CH4 ---------------Like any reaction, this
“formation” reaction involves energy (enthalpy of
reaction)
 The enthalpy of the “formation reaction” is the
“enthalpy of formation.”
Standard Enthalpy of Formation, ΔHf°
Balanced thermochemical equations are written to
depict the formation of one mole of compound from
elements (in their most stable form):
2C(graphite) + 3 H2(g) + ½ O2(g)  C2H5OH(l)
ΔHf° = -277.7 kJ
6C(graphite) + 6H2(g) + 3O2(g)  C6H12O6(s)
ΔHf° = -1273.02 kJ
Standard Enthalpy of Formation, ΔHf°
 Standard enthalpies of formation
 Given for 1 mole of product from its
constituents at a “standard state”
(atmospheric pressure, 1 atm) and 298 K (25 °C)
Standard Enthalpies of Formation, Hf°
Standard enthalpies of formation, Hf°, are measured under
standard conditions (25 °C and 1.00 atm pressure).
Practice
 Write the equation describing the formation of
propane (C3H8) gas and indicate its heat of
formation
Using
to find H
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
H =  nHf°products –  mHf° reactants
Hf° [CO2, g] =
Hf° [H2O, l] =
Hf° [C3H8] =
Hf° [O2] =
H =
QUESTION
 Calculate the enthalpy change of reaction
using the data of

Enthalpy Change of Combustion
 is the heat evolved upon the complete
combustion of 1 mol of substance (Datatable
13)
 Symbol:
Unit: KJ/mol
 The negative enthalpy change indicates an
exothermic reaction so the value would be
included on the product side
Using
to find H
Can we test the enthalpy
change of combustion?
S
Questions
 Why do you think the lab can be used to test
the energy of comubstion?
 Which evidence can indicate the
Change of temperature
Temperature Changes
 Why is a
temperature
change of 1 °C
equal to that of
1 degree K?
 Is a ΔT of 1 °C
equal to a
change of 1 °F?
 Any other factors will influence the change of
the tempertature?
The amount of water
The type of solution
Specific Heat Capacity
 The quantity of heat required to raise the
temperature of a substance by one degree
Celsius is called the specific heat capacity of 1
gram of the substance.
 Unit: J/g K
Specific Heat Capacity
1. Which of the substances in the table changes temperature the most
when the same amount of heat is applied to an equal mass of each
substance?
2. Which substance changes temperature the least with the same heat?
Conclusion
is responding to
The change of the temperature
The type of solution(Water)
The amount of water(mass)
Formula
Q  C  m  T
units of each variable?
C: J/g/K
m: g
T : K
Q: ?
Question
 What is the specific heat of water if it takes
4.184 J to raise the temperature of 1 gram of
water from 14 to 15 °C?
1. 4.184 J/g.K
2. 4.184 J/mol.K
3. 2.092 J/g.K
4. 2.092 J/mol.K
Calculate the
with the given data
Using data table 13 to compare the ideal
with the experimental one,
and calculate the Error%(keep 2 decimal places)
Compare with the data table
Not all the heat produced by the reactions is transferred to the new water
The combustion of alcohol is not complete owing to limited oxygen being
available
Not standard conditions
The container is used
in testing the
enthalpy change(H)
of any exothermic or
endothermic
reaction.
Enthalpy change of Neutralization
 We all know that neutralizaiton is and reaction
with exothermic change. Can we measure the
enthalpy change of a neutralization reaction?
Coffee-cup Calorimeter
Things to keep in mind
 In calorimetry experiments, you must have clear in your
mind:
 What is your system?
 What are your surroundings?
 You are measuring changes in energy of your system
indirectly by measuring changes in the temperature
of your surroundings
■ When calculating the heat Q, make sure to use the
mass and specific heat capacity of your surrounding
material . Then relate this heat to the reaction
taking place in your system.
Question
A 5.00-g sample of substance X is burned in a calorimeter
whose heat capacity is 5.00 kJ/ °C. The temperature
increased from 23.0 to 25.0 °C. Calculate the heat
produced by the combustion of this sample.
1. -10.0 kJ
2. -10.0 J
3. 2.00 kJ
4. -2.00 kJ
5. -50.0 kJ
Some more things to consider
 The coffee cup calorimeter the measured
change in enthalpy for a reaction will always be
lower than the actual value, why?(systematic
errors)
Hess’ Law
S
Hess’ Law
 Remember:
enthalpy is a state
function
 Hess’ Law: If a
reaction is carried
out in a series of
steps, ΔH for the
overall reaction
equals the sum of
the enthalpy
changes for the
individual steps
Practice Problem – Hess’ Law
 How can we calculate the enthalpy change for
the reaction, 2 H2(g) + N2(g) → N2H4(g) using
these equations ?
 3H2(g) + N2(g) → 2NH3(g)
 N2H4(g) + H2(g) → 2 NH3(g)
ΔH = -92.2 kJ
ΔH = -187.6 kJ
Clicker Question
Carbon occurs in two forms: as graphite and
diamond. Knowing this:
Cgraphite + O2(g)  CO2(g) ΔH = -393.5 kJ
Cdiamond + O2(g)  CO2(g)
ΔH = -395.4 kJ
Calculate ΔH for Cgraphite  Cdiamond
1. - 788.0 kJ
2. - 1.9 kJ
3. + 1.9 kJ
Extra Credit Problem
What is the energy of reaction for the following process:
6A + 9B + 3D + F → 2G
Given that:
C → A + 2B
∆H = 20.2 kJ/mol
2C + D → E + B
∆H = 30.1 kJ/mol
3E + F → 2G
∆H = -80.1 kJ/mol
 You must show every single step and explain what you are doing. You
must show the units for any numerical value. Any deviation will result in
NO POINTS WHATSOEVER.
 The presentation (look) of your work will also be graded. Let’s begin to
show professionalism. No strikethroughs, no broken pieces of paper, no
other stuff in the page, answers clearly displayed.
Hess’ Law
 FYI:
 Tables exist which list
enthalpies of:
■
■
■
■
■
Combustion
Vaporization
Fusion (melting)
Formation
etc.
 These can be used
together with Hess’ law
to determine the heat
associated with a specific
process.
Calculating Enthalpies of Reaction for
Enthalpies of Formation
We can use Hess’s law to calculate heats of
reaction from heats of formation:
HRxn =  nHf°products –  mHf° reactants
where n and m are the stoichiometric
coefficients
Calculating Enthalpies of Reaction
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
Bond Enthalpies
Starter activity
Can you write equations for the ΔHθc of glucose (C6H12O6)?
Bond breaking and bond making
Chemical reactions involve bond breaking and bond
making.
Bond energy
The quantity of energy needed to break a particular bond in a
molecule is called the bond dissociation enthalpy (H diss), or
bond enthalpy for short. It refers to the enthalpy change when
one mole of bonds of the same type are broken in gaseous
molecules under standard conditions.
Bond energy
H – H (g)  H(g)
+ H(g)
Bond
Number
broken
Number
formed
Average Molar
Bond Enthalpy
kJ mol-1
C-C
1
0
+347
C-H
5
0
+413
C-O
1
0
+358
O-H
1
6
+464
O=O
3
0
+498
C=O
0
4
+805
Here bond enthalpies are defined endothermically.
Bond breaking:
Total endothermic value = (+347 x 1) + (+413 x 5) + (+358 x
1) + (+464 x 1) + (+498 x 3) = +4728 kJ
Bond making:
Total exothermic value = (-464 x 6) + (-805 x 4) = -6004 kJ
Sum total of bond breaking and bond making:
Hc = +4728 + - 6004 = -1276 kJ mol-1
Standard enthalpy change of formation, ΔHθf, 298 is the
enthalpy change when 1 mole of a compound is formed from
its elements under standard conditions (100kPa and 298K),
all reactants and products being in their standard states.
Standard enthalpy change of combustion, ΔHθc,298 is the
enthalpy change when 1 mole of a substance is burned
completely in oxygen under standard conditions (100kPa
and 298K), all reactants and products being in their standard
states.