THERMOCHEMISTRY

1

• CHAPTER 11
• THERMOCHEMISTRY
• The study of heat changes in chemical reactions

2

• ENERGY TRANSFORMATIONS
• Energy- defined as the capacity of doing work or
  supplying heat.
• Energy is detected by observing the effects that
  it has.
• It is weightless, odorless, and tasteless.
• Chemical potential energy is the energy stored
  within the bonds of chemical substances.
• Kinds and arrangements of the atoms in a
  substance determine the amount of energy stored.
• Heat- represented by q, is energy that transfers
  from one substance to another.

3

• IMPORTANT CONCEPT
• When substances have a difference in
  temperature,heat always travels from warmer
  temperature to cooler.
• The flow from a warmer object towards a cooler
  will occur until the temperature of both objects
  becomes the same.

4

• EXOTHERMIC AND ENDOTHERMIC PROCESSES
• When we study heat changes within our universe,
  it is best to define the point of interest as a
  system.
• Everything outside of the system is defined as
  surroundings.
• The following slide demonstrates 3 different
  kinds of systems that can exist in our universe.

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The system is the specific part of the universe
that is of interest in the study.
open
closed
isolated
energy
No exchange
mass energy
Exchange
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• The law of conservation of energy states that in
  any chemical or physical process, energy is
  neither created nor destroyed.
• Energy changes can be accounted for in terms of
  work, stored energy, or heat.
• A process that absorbs heat from its surrounding
  is called an endothermic process.
• A process that releases heat to its surroundings
  is an exothermic process.

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• The lab that we did the other day demonstrated an
  endothermic reaction when the temperature of
  water in the flask increased.
• The increase in water temperature showed that the
  water was absorbing the heat given off by the
  food as it burned.
• The burning food was an exothermic process and it
  released heat to its surroundings.

8
A basic calorimeter
9

• HEAT CAPACITY AND SPECIFIC HEAT
• Exercise generates heat by the breakdown of
  stored glucose and fats and converting them to
  carbon dioxide and water during the process of
  cell respiration.
• If 10 g of glucose is broken down in your body, a
  certain amount of heat is generated.
• If you were to burn the same amount of glucose in
  an open flame, the same amount of heat would be
  produced and the release of water and carbon
  dioxide would be also be released.

10

• The amount of energy stored in the chemical bonds
  of food is thought of in terms of calories.
• A calorie is defined as the quantity of heat
  needed to raise the temperature of water 1C.
• A calorie in the food that we eat is symbolized
  with a capital C. It is equal to one
  kilocalorie, or 1000 calories.

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• The calorie is also related to the joule, the SI
  unit of heat and energy named for the English
  physicist who defined it.
• One joule of heat is raises the temperature of 1
  g of pure water 0.2390C.
• 4.184 J 1 cal

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• Heat Capacity is defined as the amount of heat
  needed to increase the temperature of an object
  exactly 1C.
• The heat capacity depends on two factors
• 1. the mass of the object
• 2. chemical composition
• The greater the mass of the object, the greater
  the heat capacity.

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• To demonstrate how both the mass and the chemical
  composition play a role consider this comparison
• 20 kg puddle and a 20 kg piece of an iron sewer
  cover exposed to the hot burning sun on a hot
  summer day.
• The iron has a lower heat capacity than the water
  because of the molecules that make up the metal.
• As a result more heat is given off by the metal
  as heat is absorbed.

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• The water molecules have a capacity to capture
  the heat and stay cooler than the iron sewer
  cover before heat is given off.
• The specific heat capacity or, stated simply, the
  specific heat of a substance is the amount of
  heat it takes to raise the temperature of
• 1 g of the substance to 1 C.
• 1 calorie of heat raises 1 g of water1C
• 1 calorie of heat raises 1 g of iron 9 C.
• Water has a specific heat 9 times greater than
  iron.

15

• Table 11.6 on page 296 lists the specific heat of
  some common substances.
• The relationship between specific heat and heat
  capacity is as follows
• Heat capacity specific heat X mass in grams

16

• If you examine the table on page 296, you can see
  that water has the highest specific heat compared
  to the other substances listed.
• That means that water has a higher heat capacity
  (can absorb more heat) than any of those
  substances.

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• Specific heat is independent of the amount of
  overall mass of the substance, but mass will play
  a role on its heat capacity. The higher the mass,
  the higher its heat capacity.

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• To calculate the specific heat of a substance
  during an increase in temperature, you divide
  heat input by the temperature change times the
  mass of the substance.
• The formula for specific heat (C) is as follows
• C q__ heat(Joules or calories)
  
• m X ? T mass(g) X change in
  temperature

19

• Sample problem
• The temperature of a piece of copper with a mass
  of 95.4 g increases from 25.0 C to 48.0C when
  the metal absorbs 849 J of heat. What is the
  specific heat of copper?
• Knowns m 95.4 g
• ? T (48.0C 25.0 C) 23.0C
• q 849 J
• C 849 J
• 95.4 g X 23.0 C
• 0.387 J/g X C

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• Open up your text to page 299
• do problems
• 4 through 9
• at your seat

21

• 4. Define energy and explain how energy and heat
  are related.
• Energy is the capacity for doing work or
  supplying heat. Heat is energy that transforms
  between objects across a temperature
  gradient.(the term gradient is used to indicate
  that there is a temperature difference between
  the objects.)
• 5. Explain the difference between heat capacity
  and specific heat.
• The specific heat of a substance is independent
  of its mass. The heat capacity of an object is
  proportional to its mass.

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• 6.Will the specific heat of 50g of a substance be
  the same as, or greater than, the specific heat
  of 10g of the same substance?
• The same as. If you were asked to compare the
  heat capacity then the higher mass would have a
  larger heat capacity.
• By definition, specific heat is the amount of
  energy required to raise 1g of any substances
  temperature 1C. So whether we are talking about
  10 grams of a like substance or 1000 g of that
  same substance, the specific heat will be the
  same

23

• 7. On a sunny day, why does the concrete deck
  around an outdoor swimming pool become hot, while
  the water stays cool?
• Water has a higher heat capacity.
• 8. Using calories, calculate how much heat 32.0 g
  of water absorbs when water is heated from 25.0C
  to 80.0C. How many joules is this?
• q m X ?T X C
• q 32.0g X 55.0C X 1.00cal 1,760 cal
  or1.76kcal
• 1,760 cal 7.36 X 103 J because 1 cal 4.184 J
• 4.184 J X 1760 cal 7.36 X 103

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• 9. A chunk of silver has a heat capacity of 42.8
  J/ C.
• If the silver has a mass of 181g, calculate the
  specific heat of silver.
• Specific heat heat capacity
• mass
• 42.8 J / C
• 181 g
• 2.36 X 10-1 J/g X C

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• BEFORE YOU DO 10,
• YOU NEED TO KNOW THAT
• 1 g of water 1 mL of water
• 1L 1000 mL

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• 10. How many kilojoules of heat are absorbed when
  1.00 L of water is heated from 18C to 85 C.
• Knowns
• 1.00L of water 1000g
• Specific heat of water(from chart) 4.18 J(g X
  C) or 4.18 X 10-3kJ(g X C)
• Note 1000 J 1 kJ
• 4.18 J 1000 will convert J to kJ
• ? T 85C 18 C 67C
• q 1000g X 67C X 4.18 X 10-3 2.8 X 102kJ

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• 11.2 MEASURING AND EXPRESSING HEAT CHANGES
• Calorimetry- the accurate and precise measurement
  of heat change for chemical and physical
  processes.
• In calorimetry, the heat released by the system
  is equal to the heat absorbed by the
  surroundings.
• This obeys the law of conservation of energy.
• A simple calorimeter made with a foam cup and a
  foam lid can be used to test the amount of heat
  change during a chemical reaction.
• The following slide demonstrates how you would
  set this up.

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Constant-Pressure Calorimetry
1. Dissolve reacting chemicals (system) in a
known volume of water(surroundings) take the two
different solutions initial temperatures.
2.Place the two solutions into the calorimeter
and mix them together.
3. After the reaction is complete, take the final
temperature.
4. Because you know the initial and final
temperatures and the heat capacity of water, you
can calculate the amount of heat released(q) or
absorbed in the reaction using the equation for
specific heat.
No heat enters or leaves!
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• Calorimetry experiments can also be performed at
  constant volume using a bomb calorimeter.
• A bomb calorimeter, shown on next slide, measures
  the heat released from a burning compound. The
  calorimeter is a closed system that is, the mass
  of the system is held at a constant.

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• BOMB CALORIMETER- in a bomb calorimeter the
  system is closed in order to prevent mass loss. A
  sample is burned in a constant-volume chamber in
  the presence of oxygen at high-pressure.

Oxygen intake valve
Electrical leads with firing element
thermometer
stirrer
Sample to be burned
water
31

• Sample problem
• A student mixed 50.0 mL of water containing 0.50
  mol HCl at 22.5C with 50.0 mL of water
  containing 0.50 mol NaOH at 22.5C in a foam cup
  calorimeter. The temperature of the resulting
  solution increased to 26.0C. How much heat in kJ
  was released by the reaction?

• C q__ heat(Joules or calories)
  
• m X ? T mass(g) X change in
  temperature
• q m X ?T X C
• Mass 100g(50mL 50mL converted to grams)
• ?T 26.0 22.5 3.5C
• C 4.18 J 1000 0.00418 kJ
• q 100 g X 3.5C X .00418kJ 1.5 kJ

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• THERMOCHEMICAL EQUATIONS
• You can treat heat change during a chemical
  reaction like any other reactant and product.
• An equation that includes the heat change is
  called a thermochemical equation
• A heat of reaction is the heat change for the
  reaction exactly as it is written.
• You will usually see the heat of reaction written
  as
• ?H, which is the heat change at constant
  pressure.
• The physical states of the matter must also be
  given.
• conditions are at STP1 atmosphere or(101.3 kPa)
  and at room temperature (23C)

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• Some reactions will be endothermic and some
  exothermic.
• Consider the heat of reaction of the following
  equation which demonstrates what occurs when
  calcium oxide is mixed with water to create
  concrete.
• CaO(s) H2O(l)?Ca(OH)2(s) 65.2kJ
• If the heat is expressed on the product side,
  then heat is released.
• This means it is an exothermic reaction
• and ?H -65.2kJ
• Note that a negative signs is used to indicate
  exothermic

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• An example of an endothermic reaction would be
  the use of baking soda(sodium hydrogen carbonate)
  when baking.
• NaHCO3 decomposes when it is heated and releases
  carbon dioxide,
• it is the carbon dioxide that causes the cake to
  raise. Because this is an endothermic reaction,
  ?H will be written with a positive sign.

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• The thermochemical equation for this process is
  written as follows
• Note the heat is on the reactant side of the
  equation
• Heat on the reactant side means that it is an
  endothermic reaction.
• 2NaHCO3(s) 129kJ ? Na2CO3(s) H20(g)CO2(g)
• Using this balanced equation, calculate the
  kilojoules of heat required to decompose
• 2.24 mol NaHCO3

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• Solution

• 2NaHCO3(s)129kJ?Na2CO3(s)H20(g)CO2(g)
• Knowns 2.24 mol NaHCO3 decomposes
• Unknowns ?H ?kJ
• From the equation set up an equality
• 2 mol NaHCO3 129kJ
• Now set up a conversion
• 2.24 mol NaHCO3 X 129kJ
• 2 mol NaHCO3
• 144kJ

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• Chemistry problems involving enthalpy changes are
  similar to stoichiometry problems.
• The amount of heat absorbed or released during a
  reaction depends on the number of moles of the
  reactants involved.
• In the previous endothermic reaction, the
  potential energy of the products is higher than
  the potential energy of the reactants.

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• It is important to state the physical states of
  the reactants and products because it can mean
  different ? Hs
• H2O(l)? H2(g) ½ O2(g) ?H 285.8kJ
  H2O(g)?H2(g) ½ O2(g) ?H 241.8kJ
• A difference of 44.0kJ
• It makes sense that it would require a greater
  amount of energy for water as a liquid to split
  into oxygen and hydrogen gas than it would for a
  water vapor molecule to split into oxygen and
  hydrogen gas.

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• Heat of Combustion is the heat of reaction for
  the complete burning of one mole of a substance.
• Heats of combustion for some common substances
  are listed in the table on page 305
• The combustion of natural gas which is mostly
  methane(CH4) is an exothermic process where I mol
  of CH4 will release 890kJ of heat.

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• SEAT WORK
• WORK ON PROBLEMS
• 15 TO 19
• ON PAGE 306

41

• 15. When 2 mol of solid Mg combines with 1 mol of
  O2, 2 mol of solid MgO is formed and 1204 kJ of
  heat is released.
• 2Mg(s) O2 ? 2MgO(s) 1204kJ
• 16. Ethanol(C2H5OH) when burned reacts with O2 to
  produce CO2(g) and H2O(g) How much heat is
  released when 12.5 g of C2H5OH burns?
• C2H5OH(l) 3O2 ? 2CO2(g) 3H2O(g) ?H1235kJ
• 1 mol C2H5OH 46.0g
• 46.0g C2H5OH 1235 kJ
• 12.5g C2H5OH X 1235kJ 3.36 X 102
  kJ
• 46.0g C2H5OH

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• 17. Explain the term heat of reaction
• The heat absorbed or released in a chemical
  change.
• 18. Calculate the heat change(?H) in kJ for the
  conversion of 15.0 g H2
• Note need to consider or sign
• H2(g) F2(g) ? 2HF(g) ?H -536 kJ
• Given 15.0 gH2
• ? kJ H2
• 2g H2 -536 kJ
• 15.0 g H2 X -536 kJ -4.02 X 103 kJ
• 2 g H2

43

• 19. Why is it important to give the physical
  state of a substance in a thermochemical
  reaction?
• Because a phase change will indicate
• that an energy change has taken place
• and some energy changes (solid to gas)
• can require more energy than others
• (liquid to gas)

44

• ENERGY TRANSFORMATIONS
• Energy- defined as the capacity of doing work or
  supplying heat.
• Energy is detected by observing the effects that
  it has.
• It is weightless, odorless, and tasteless.
• Chemical potential energy is the energy stored
  within the bonds of chemical substances.
• Kinds and arrangements of the atoms in a
  substance determine the amount of energy stored.
• Heat- represented by q, is energy that transfers
  from one substance to another.