Thermodynamics Standard 7

1
ThermodynamicsStandard 7

• Chemistry.
• Ms. Siddall.

2
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

3
Summary 1

• Describe the difference between heat and
  temperature

4
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.

5

• 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)

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Summary 3

• Consider an ice cube dropped into a glass of warm
  water.
• Ice cube system
• Water surroundings
• Does heat flow into the system or out of the
  system?
• What is gaining energy (system or surroundings)?

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Summary 4
Standard 7b exothermic endothermic process

• Endothermic Process A process in which energy is
  absorbed.
• Example Water boiling
• H2O(l) heat ? H2O(g)
• In an endothermic process heat is a reactant.

product
reactants
In an endothermic process which has more energy
reactants or products?
8
Summary 5

• Exothermic Process A process in which energy is
  released.
• Example
• A fire
• 3C 2O2 ? heat 2CO CO2
• In an exothermic process heat is a product

products
reactants
In an exothermic process which has more energy,
reactants or products?
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Energy diagram
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)

H2O(g)
exothermic
endothermic
Increasing energy
H2O(l)
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Transition State energy diagram
activation energy energy needed to form
transition state (activated complex)
Transition state
Energy released when products form
reactants
Total energy released during reaction
energy
products
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Summary 7

• 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

• Draw a transition state energy diagram for an
  endothermic reaction

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Summary 8

• 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

If your body burns about 2,000 food calories a
day, approximately how many joules of energy is
that?
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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 exothermic ?H endothermic
?H change in enthalpy Enthalpy energy/heat
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• N2(g) 2O2(g) ? 2NO2(g) ?H 68KJ
• N2(g) 2O2(g) 68KJ ? 2NO2(g)

Endothermic reaction
Energy is a reactant
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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?

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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.

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Summary 10

• Is the example of ice melting an endothermic
  process or an exothermic process?

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Physical state
gas
evaporating
break hydrogen bonds
Condensing
endothermic
liquid
exothermic
Energy released intermolecular attractions take
over
melting
freezing
break lattice structure
solid
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Summary 11

1. In which phase do the molecules have the most
   energy? (solid, liquid, or gas)
2. Is the process of condensing endothermic or
   exothermic?
3. Is the process of vaporization endothermic or
   exothermic?

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Freezing/boiling point graph for water.
Standard 7d solving problems
Energy absorbed no temp change physical change
boiling ?Hvap
110
100
steam
melting ?Hfus
Water (CH2O(l))
Temperature (C)
Energy absorbed Change in temperature Change
in K.E.
0
ice
-10
energy
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Summary 12

• Which two sections of the graph show no
  temperature change.
• Why is there no temperature change in these
  sections?

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Standard 7d solving problems

• Latent Heat of fusion. (latent heat hidden
  heat)
• ?Hfus The energy released when 1g of a
  substance is frozen OR the energy needed when 1g
  of a substance is melted.
• ?Hfus enthalpy of fusion (J/g)
• Fusion freezing (liquid ? solid)
• also used for melting (solid ? liquid)

Summary 13 What does fusion mean?
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Example freezing water

• How much energy is released when 10g water
  freezes? (?HfusH2O 334J/g)

334J
3340J 3.34kJ
10g H2O(s)
1g H2O(s)
Summary 14
How much energy is needed to melt 100g of water?
(show calculation)
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• Latent Heat of vaporization
• ?Hvap The energy needed when 1g of a substance
  is evaporated OR the energy released when 1g of a
  substance is condensed.
• ?Hvap enthalpy of vaporization (J/g)
• vaporization evaporating (liquid ? gas)
• also used for condensing (gas ?liquid)

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Summary 15

• What does vaporization mean?
• What does condensation mean?

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Summary 16
Example Boiling water

• How much energy is needed to boil 10g water?
  (?HvapH2O 2260J/g)

2260J
22600J 22.6kJ
10g H2O(l)
J
1g H2O(l)
How much energy is released when 100g of water
vapor is condensed? (show work)
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Heat Capacity.

• C specific heat capacity
• The amount of heat energy needed to raise the
  temperature of 1g of a substance by 1C
• Example CH2O(l) 4.18J/gC
• It takes 4.18J of energy to raise the temperature
  of 1g of water by 1C
• 1 calorie 4.18J

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Summary 17

• How much energy is needed to raise the
  temperature of 1g of water by 1C? (give your
  answer in joules and calories)

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Example.

• How much energy is needed to raise the
  temperature of 5g water from 22C to 24C?
  (CH2O(l) 4.18J/gC)

5g H2O(l)
4.18J H2O(l)
2C
41.8J
J
g C
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Summary 18

• How much energy is released when 10g water cools
  from 40C to 30C?

10C
10g H2O(l)
4.18J H2O(l)
418.J
J
g C
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Measuring specific heat capacity for different
compounds
Thermometer Measures temperature change for water
qenergy released by metal energy absorbed
by water
Unknown compound heated to 100C and placed in
the cold water
H2O
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Summary 19

• How much energy (q) is released by a metal if the
  temperature of 100g of water in the calorimeter
  rises from 20C to 30C?

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Measuring the heat of a reaction q (q energy
released or absorbed by water)

• Thermometer measures temperature change for
  water
• T ? exothermic
• T ? endothermic

Reaction chamber 3H2 N2 ? NH3 heat of
reaction is absorbed by water
100g H2O
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• Example 10g NH3 are produced in the above
  reaction. The temperature rises from 20.0C to
  30.0C.
• Calculate q (energy) for the reaction.
• Is the reaction endothermic or exothermic?
• Calculate ?H (J/g) for this reaction
• Calculate ?H (mol/g) for this reaction

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Kinetic energy distribution diagram
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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

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Summary 20

• Explain why more particles evaporate from a cup
  of hot water compared to a cup of cold water.

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Standard 7e Apply Hesss Law to calculate
enthalpy change in a reaction

• Hesss 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) ?H 181kJ
• 2NO(g) O2(g) ? 2NO2(g) ?H -113kJ
• Find the sum of the 2 equations

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• N2(g) O2(g) ? 2NO(g) ?H 181kJ
• 2NO(g) O2(g) ? 2NO2(g) ?H -113kJ
• N2(g) O2(g) 2NO(g) O2(g) ? 2NO(g) 2NO2(g)
  
• N2(g) 2O2(g) ? 2NO2(g)
• ?H
• notes
• You can reverse reactions (change sign of ?H)
• You can multiply or divide equations (do same to
  ?H)

(-113kJ)
68kJ


181kJ
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Hess summary

• Complete questions 66, 74, 81 84 on page 536
  537