Thermodynamics
Standard 7
Chemistry.
Ms. Siddall.
Standard 7a: ‘heat flow’
Chemical Thermodynamics = the
movement of heat in a chemical
reaction.
 Temperature = a measure of the
average kinetic energy of particle motion
 Heat = The transfer of energy from a
hotter object to a colder object
(sometimes called ‘heat flow’)
– temperature measures energy
– Heat measures energy transfer
Summary 1

Describe the difference between heat
and temperature
Energy transfer


Particle vibrations increase when a
particle gains energy
Vibrations are transferred to surrounding
particles
Summary 2
Describe how energy is transferred
between atoms.
Identifying heat transfer:
 System: experiences a change
 Surroundings: causes a change
e.x.
hot coffee (system) cools
because it transfers heat to the
air, the cup, the table & the
whole universe! (surroundings)
Summary 3
Consider an ice cube dropped into a
glass of warm water.
 Ice cube = system
 Water = surroundings
1. Does heat flow into the system or
out of the system?
2. What is gaining energy (system or
surroundings)?
Standard 7b: exothermic & endothermic
process
Endothermic Process: A process in which
energy is absorbed.
 Example:
– Water boiling
– H2O(l) + heat  H2O(g)
reactants

product
In an endothermic process heat is a reactant.
Summary 4

In an endothermic process which has
more energy; reactants or products?
Exothermic Process: A process in which
energy is released.
 Example:
– A fire
– 3C + 2O2  heat + 2CO + CO2
reactants

products
In an exothermic process heat is a product
Summary 5

In an exothermic process which has
more energy, reactants or products?
exothermic
Increasing energy
H2O(l)
endothermic
Energy diagram
H2O(g)
Summary 6

Draw an energy diagram for the campfire
reaction.
– Show reactants and products.
– Draw only one arrow from reactants to
products and label the arrow (endothermic or
exothermic)
Transition State energy diagram
activation energy = energy needed to
form transition state (activated complex)
Transition state
energy
reactants
products
Energy released
when products form
Total energy
released during
reaction
Transition State: An intermediate state



that can occur during a reaction
Also called an ‘activated complex’
An exothermic reaction is not always
spontaneous because energy is
needed to form a transition state.
e.x. a spark is needed to start a fire
Summary 7

Draw a transition state energy diagram
for an endothermic reaction
Measuring heat flow.
 Energy is measured in joules (J) or
calories (cal)
 Example: 334J of energy are needed
to melt 1g of ice.
 1 calorie (c) = 4.18J
 1 food calorie (C) = 1000 calories =
4180J
Summary 8

If your body burns about 2,000 food
calories a day, approximately how
many joules of energy is that?
Energy released = exothermic
KJ = kilojoules = 1000J
Showing a change in energy:
 S(s) + O2(g)  SO2(g) + energy
 S(s) + O2(g)  SO2(g) + 297KJ
 S(s) + O2(g)  SO2(g)
∆H = -297KJ
∆H = change in enthalpy
Enthalpy = energy/heat
-∆H = exothermic
+∆H = endothermic

N2(g) + 2O2(g)  2NO2(g)
∆H = +68KJ

N2(g) + 2O2(g) + 68KJ  2NO2(g)
Endothermic reaction
Energy is a reactant
Summary 9


Write an equation to show water
melting. Use ∆H to show energy.
(it takes 5.9kJ of energy to melt ice)
Is ∆H negative or positive? Why?
Standard 7c: energy of phase change
Phase Change: The physical state of a
compound changes
 The same compound is observed before
and after the change
 Example: ice melting H2O(s)  H20(l)
 There is no temperature change.
 Energy is used to overcome intermolecular
attractions.
Summary 10

Is the example of ice melting an
endothermic process or an exothermic
process?
Condensing
endothermic
break hydrogen
bonds
solid
break lattice
structure
liquid
evaporating
gas
melting
freezing
Energy released intermolecular
attractions take over
exothermic
Physical state
Summary 11
1.
2.
3.
In which phase do the molecules
have the most energy? (solid, liquid,
or gas)
Is the process of condensing
endothermic or exothermic?
Is the process of vaporization
endothermic or exothermic?
Standard 7d:
Freezing/boiling point graph for water. solving problems
Temperature (°C)
110
Energy absorbed
= no temp change
= physical change
boiling
ΔHvap
100
steam
melting
ΔHfus
0
-10
ice
energy
Water (CH2O(l))
Energy absorbed
= Change in temperature
= Change in K.E.
Summary 12


Which two sections of the graph show
no temperature change.
Why is there no temperature change
in these sections?
Standard 7d: solving problems
Latent Heat of fusion. (latent heat = hidden heat)
ΔHfus = The energy released when 1g of a
substance is frozen OR the energy needed
when 1g of a substance is melted.
ΔHfus = enthalpy of fusion (J/g)
Fusion = freezing (liquid  solid)
• also used for melting (solid  liquid)
•
Summary 13

What does ‘fusion’ mean?
Example: freezing water

How much energy is released when 10g
water freezes? (ΔHfusH2O = 334J/g)
10g H2O(s)
334J
1g H2O(s)
= 3340J
=J
= 3.34kJ
Summary 14

How much energy is needed to melt
100g of water? (show calculation)
Latent Heat of vaporization
ΔHvap = The energy needed when 1g of a
substance is evaporated OR the energy
released when 1g of a substance is
condensed.
ΔHvap = enthalpy of vaporization (J/g)
•
vaporization = evaporating (liquid  gas)
•
also used for condensing (gas liquid)
Summary 15


What does vaporization mean?
What does condensation mean?
Example: Boiling water

How much energy is needed to boil 10g
water? (ΔHvapH2O = 2260J/g)
10g H2O(l)
2260J
1g H2O(l)
= 22600J
=J
= 22.6kJ
Summary 16

How much energy is released when
100g of water vapor is condensed?
(show work)
Heat Capacity.

C = specific heat capacity
• The amount of heat energy needed to
raise the temperature of 1g of a
substance by 1°C

Example: CH2O(l) = 4.18J/g°C
• It takes 4.18J of energy to raise the
temperature of 1g of water by 1°C
 1 calorie = 4.18J
Summary 17

How much energy is needed to raise
the temperature of 1g of water by
1°C? (give your answer in joules and
calories)
Example.

How much energy is needed to raise
the temperature of 5g water from
22°C to 24°C? (CH2O(l) = 4.18J/g°C)
5g H2O(l)
4.18J H2O(l) 2°C
g °C
==41.8J
J
Summary 18

How much energy is released when 10g
water cools from 40°C to 30°C?
10g H2O(l)
4.18J H2O(l) 10°C
g °C
==418.J
J
Measuring specific heat capacity for
different compounds
Thermometer: Measures
temperature change for water
‘q’=energy released by metal
=energy absorbed by water
H2O
Unknown compound:
heated to 100°C and
placed in the cold water
Summary 19

How much energy (q) is released by a
metal if the temperature of 100g of
water in the calorimeter rises from
20°C to 30°C?
Measuring the heat of a reaction: ‘q’
(q = energy released or absorbed by water)
Thermometer: measures
temperature change for water
• T  = exothermic
• T  = endothermic
100g H2O
Reaction chamber:
3H2 + N2  NH3
heat of reaction is
absorbed by water
Example: 10g NH3 are produced in the above
1.
2.
3.
4.
reaction. The temperature rises from
20.0°C to 30.0°C.
Calculate ‘q’ (energy) for the reaction.
Is the reaction endothermic or
exothermic?
Calculate ΔH (J/g) for this reaction
Calculate ΔH (mol/g) for this reaction
Kinetic energy distribution diagram
Kinetic energy distribution diagram
T1 = low temperature = low energy
 T2 = higher temperature = higher
energy


Emin = minimum energy needed to
escape.
– More T2 particles have Emin
– Less T1 particles have Emin
Summary 20

Explain why more particles evaporate
from a cup of hot water compared to a
cup of cold water.
Standard 7e: Apply Hess’s Law to
calculate enthalpy change in a
reaction


Hess’s Law: If a series of reactions are
added together the enthalpy change for the
net reaction will be the sum of the enthalpy
changes for the individual steps.
E.x. N2(g) + 2O2(g)  2NO2(g)
• N2(g) + O2(g)  2NO(g) ΔH = +181kJ
• 2NO(g) + O2(g)  2NO2(g) ΔH = -113kJ

Find the sum of the 2 equations…
N2(g) + O2(g)  2NO(g) ΔH = +181kJ
 2NO(g) + O2(g)  2NO2(g) ΔH = -113kJ

N2(g) + O2(g) + 2NO(g) + O2(g)  2NO(g) + 2NO2(g)
N2(g) + 2O2(g)  2NO2(g)
ΔH = +181kJ + (-113kJ) = +68kJ
notes:


You can reverse reactions (change sign of ΔH)
You can multiply or divide equations (do same
to ΔH)
Hess summary

Complete questions 66, 74, 81 & 84
on page 536 & 537