Chapter 16 - Energy and Chemical Change

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Chapter 16 - Energy and Chemical Change
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Energy and Heat

• Thermochemistry - concerned with heat changes
  that occur during chemical reactions
• Energy - capacity for doing work or supplying
  heat
• weightless, odorless, tasteless
• if within the chemical substances- called
  chemical potential energy

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Energy and Heat

• Gasoline contains a significant amount of
  chemical potential energy
• Heat - represented by q, is energy that
  transfers from one object to another, because of
  a temperature difference between them.
• only changes can be detected!
• flows from warmer ? cooler object

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Exothermic and Endothermic Processes

• Essentially all chemical reactions, and changes
  in physical state, involve either
• release of heat, or
• absorption of heat

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Exo- and Endothermic Processes

• In studying heat changes, think of defining these
  two parts
• the system - the part of the universe on which
  you focus your attention
• the surroundings - includes everything else in
  the universe

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Exo- and Endothermic Processes

• Together, the system and its surroundings
  constitute the universe
• Thermochemistry is concerned with the flow of
  heat from the system to its surroundings, and
  vice-versa.

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Exothermic and Endothermic Processes

• The Law of Conservation of Energy states that in
  any chemical or physical process, energy is
  neither created nor destroyed.
• All the energy is accounted for as work, stored
  energy, light, or heat.

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Exothermic and Endothermic Processes

• Heat flowing into a system from its surroundings
• defined as positive
• q has a positive value
• called endothermic
• system gains heat as the surroundings cool down

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Exo- and Endothermic Processes

• Heat flowing out of a system into its
  surroundings
• defined as negative
• q has a negative value
• called exothermic
• system loses heat as the surroundings heat up

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Exothemic and Endothermic

• Every reaction has an energy change associated
  with it
• Exothermic reactions release energy, usually in
  the form of heat.
• Endothermic reactions absorb energy
• Energy is stored in bonds between atoms

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Heat Capacity and Specific Heat

• A calorie is defined as the quantity of heat
  needed to raise the temperature of 1 g of pure
  water 1 oC.
• Used except when referring to food
• a Calorie, written with a capital C, always
  refers to the energy in food
• 1 Calorie 1 kilocalorie 1000 cal.

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Heat Capacity and Specific Heat

• The calorie is also related to the joule, the SI
  unit of heat and energy
• named after James Prescott Joule
• 4.184 J 1 cal
• Heat Capacity - the amount of heat needed to
  increase the temperature of an object exactly 1 oC

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Heat Capacity and Specific Heat

• Specific Heat Capacity - the amount of heat it
  takes to raise the temperature of 1 gram of the
  substance by 1 oC (abbreviated c)
• often called simply Specific Heat
• Water has a HUGE value, compared to other
  chemicals
• Movie

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Specific Heats of Common Substances are provided
in a table in your book.
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Heat Capacity and Specific Heat

• For water, c 4.184 J/(g oC), and also c 1.00
  cal/(g oC)
• Thus, for water
• it takes a long time to heat up, and
• it takes a long time to cool off!
• Water is used as a coolant!

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Heat Capacity and Specific Heat

• To calculate, use the formula
• q c x mass (g) x ?T
• heat abbreviated as q
• ?T change in temperature
• c Specific Heat
• Units are either J/(g oC) or cal/(g oC)

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Energy Conversion Factors are given in a table in
your book.
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Section 16.2Heat in Chemical Reactions and
Processes

• OBJECTIVES
• Calculate heat changes in chemical and physical
  processes.

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Calorimetry

• Calorimetry - the accurate and precise
  measurement of heat change for chemical and
  physical processes.
• The device used to measure the absorption or
  release of heat in chemical or physical processes
  is called a Calorimeter

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Calorimetry

• Foam cups are excellent heat insulators, and are
  commonly used as simple calorimeters
• Fig. 16.5, page 497

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Calorimetry

• Calorimetry experiments can be performed at a
  constant volume using a device called a bomb
  calorimeter
• For systems at constant pressure, the heat
  content is the same as a property called Enthalpy
  (H) of the system

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Calorimetry

• Changes in enthalpy ?H
• q ?H These terms will be used interchangeably
  in this textbook
• Thus, q ?H m x c x ?T
• ?H is negative for an exothermic reaction
• ?H is positive for an endothermic reaction

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In terms of bonds
O
C
O
Breaking this bond will require energy.
Making these bonds gives you energy.
In this case making the bonds gives you more
energy than breaking them.
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Exothermic

• The products are lower in energy than the
  reactants
• Releases energy

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C O2 CO2
395 kJ
395kJ
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Endothermic

• The products are higher in energy than the
  reactants
• Absorbs energy
• Movie

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CaCO3 CaO CO2
CaCO3 176 kJ CaO CO2
176 kJ
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Sec. 16.23 Chemistry Happens in

• MOLES
• An equation that includes energy is called a
  thermochemical equation
• CH4 2O2 CO2 2H2O 802.2 kJ
• 1 mole of CH4 releases 802.2 kJ of energy.
• When you make 802.2 kJ you also make 2 moles of
  water

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Thermochemical Equations

• A heat of reaction is the heat change for the
  equation.
• The physical state of reactants and products must
  also be given.
• Standard conditions for the reaction are 101.3
  kPa (1 atm.) and 25 oC

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CH4 2 O2 CO2 2 H2O 802.2 kJ

• If 10. 3 grams of CH4 are burned completely, how
  much heat will be produced?

1 mol CH4
802.2 kJ
10. 3 g CH4
16.05 g CH4
1 mol CH4
515 kJ
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CH4 2 O2 CO2 2 H2O 802.2 kJ

• How many liters of O2 at STP would be required to
  produce 23 kJ of heat?
• How many grams of water would be produced with
  506 kJ of heat?

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Summary, so far...
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Enthalpy

• The heat content a substance has at a given
  temperature and pressure
• Cant be measured directly because there is no
  set starting point
• The reactants start with a heat content
• The products end up with a heat content
• So we can measure how much enthalpy changes

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Enthalpy

• Symbol is H
• Change in enthalpy is DH (delta H)
• If heat is released, the heat content of the
  products is lower
• DH is negative (exothermic)
• If heat is absorbed, the heat content of the
  products is higher
• DH is positive (endothermic)

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Energy
Change is down
DH is lt0
Reactants
Products

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Energy
Change is up
DH is gt 0
Reactants
Products

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Heat of Reaction

• The heat that is released or absorbed in a
  chemical reaction
• Equivalent to DH
• C O2(g) CO2(g) 393.5 kJ
• C O2(g) CO2(g) DH -393.5 kJ
• In thermochemical equation, it is important to
  indicate the physical state
• H2(g) 1/2O2 (g) H2O(g) DH -241.8 kJ
• H2(g) 1/2O2 (g) H2O(l) DH -282.5 kJ

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Heat of Combustion

• The heat from the reaction that completely burns
  1 mole of a substance.

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Section 16.3Heat in Changes of State

• OBJECTIVES
• Calculate heat changes that occur during melting,
  freezing, boiling, and condensing.

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Heats of Fusion and Solidification

• Molar Heat of Fusion (?Hfus) - the heat absorbed
  by one mole of a substance in melting from a
  solid to a liquid
• Molar Heat of Solidification (?Hsolid) - heat
  lost when one mole of liquid solidifies Movie

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Heats of Fusion and Solidification

• Heat absorbed by a melting solid is equal to heat
  lost when a liquid solidifies
• Thus, ?Hfus -?Hsolid
• Note Table 16.6, page 502
• Sample Problem 16-4, page 504

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Heats of Vaporization and Condensation

• When liquids absorb heat at their boiling points,
  they become vapors.
• Molar Heat of Vaporization (?Hvap) - the amount
  of heat necessary to vaporize one mole of a given
  liquid.

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Heats of Vaporization and Condensation

• Condensation is the opposite of vaporization.
• Molar Heat of Condensation (?Hcond) - amount of
  heat released when one mole of vapor condenses
• ?Hvap - ?Hcond

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Heats of Vaporization and Condensation

• The large values for ?Hvap and ?Hcond are the
  reason hot vapors such as steam is very dangerous
• You can receive a scalding burn from steam when
  the heat of condensation is released!

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Heats of Vaporization and Condensation

• H20(g) ? H20(l) ?Hcond - 40.7kJ/mol

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Heat of Solution

• Heat changes can also occur when a solute
  dissolves in a solvent.
• Molar Heat of Solution (?Hsoln) - heat change
  caused by dissolving one mole of a substance
• Sodium hydroxide provides a good example of an
  exothermic molar heat of solution

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Heat of Solution

• NaOH(s) ? Na1(aq) OH1-(aq)
• ?Hsoln - 445.1 kJ/mol
• The heat is released as the ions separate and
  interact with water, releasing 445.1 kJ of heat
  as ?Hsoln thus becoming so hot it steams!

H2O(l)
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Section 16.4Calculating Heat Changes

• OBJECTIVES
• Apply Hesss law of heat summation to find heat
  changes for chemical and physical processes.

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Hesss Law

• If you add two or more thermochemical equations
  to give a final equation, then you can also add
  the heats of reaction to give the final heat of
  reaction.
• Called Hesss law of heat summation

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Why Does It Work?

• If you turn an equation around, you change the
  sign
• If H2(g) 1/2 O2(g) H2O(g) DH-285.5 kJ
• then, H2O(g) H2(g) 1/2 O2(g) DH
  285.5 kJ
• also,
• If you multiply the equation by a number, you
  multiply the heat by that number
• 2 H2O(g) 2 H2(g) O2(g) DH 571.0 kJ

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Why does it work?

• You make the products, so you need their heats of
  formation
• You unmake the products so you have to subtract
  their heats.
• How do you get good at this?

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Applying Hesss Law

• 2H2O2(l) ? 2H2O(l) O2(g)
• Can be shown in two equations
• a) 2H2(g)O2(g) ? 2H2O(l) ?H -572 kJ
• b) H2(g) O2(g) ? H2O2(l) ?H -188 kJ
• Reverse the equation b) to make H2O2 a reactant
  and then double it.
• c) 2H2O2(l)?2H2(g)2O2(g) ?H 376 kJ
• Note how ?H -188 kJ was doubled and its sign was
  reversed.

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Applying Hesss Law

• Adding equations a) and c) reactants and products
  together yields
• 2H2O2(l) 2H2(g) O2(g) ?
• 2H2(g) 2O2(g) 2H2O(l)
• Canceling like items of both sides reduces the
  equation to
• 2H2O2(l) ? 2H2O(l) O2(g)
• Adding equations a) and c) enthalpies yields
• 376 kJ -572 kJ -196 kJ

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Standard Heats of Formation

• The DH for a reaction that produces one mol of a
  compound from its elements at standard conditions
• Standard conditions are 25C and 1 atm.
• The symbol is

• The standard heat of formation of an element is
  zero element 0
• This includes the diatomics like O2

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Using Standard Heat of Formation

• Table 16.7, page 510 has standard heats of
  formation for numerous compounds
• The heat of a reaction can be calculated by
  subtracting the heats of formation of the
  reactants from the products

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Example Standard Heat of Formation

• CH4(g) 2 O2(g) CO2(g) 2 H2O(g)