Objectives/Goals for Today

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Objectives/Goals for Today

• Exam over Chapters 8 9
• Chapter ten notes
• Section 10.1
• Section 10.2
• Section 10.3
• Section 10.4
• Section 10.5
• Section 10.6
• Section 10.7
• Section 10.8

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Chapter TenEnergy Changes in Chemical Reactions
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Section 10.1Energy Energy Changes
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Energy

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System and Surroundings

• System Part we care about
• Reactants Products
• Surroundings Everything
• else in the universe

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Types of Systems

• A?open system (mass and heat pass through)
• B?closed system (heat only pass through)
• C?isolated system (no heat or mass transfer)

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Energy Flow and Reactions

• For chemical reactions to happen spontaneously,
  the final products must be more stable than the
  starting reactants
• Higher energetic substances
• Typically less stable, more reactive
• Lower energetic substances
• Typically more stable, less reactive

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Exothermic and Endothermic

• Thermal energy flows from warmer to cooler

H2O(s) ? H2O(l)
2H2(g) O2(g) ? 2H2O(l)
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Section 10.2Introduction to Thermodynamics
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What is Thermodynamics

• Study of heat and its transformations into other
  energies
• Thermochemistry is a part of this
• Thermodynamics studies changes in the state of a
  system

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State Functions

• State functions are properties that are
  determined by the state of the system, regardless
  of how it was achieved
• Final Initial
• Ex
• Energy
• Pressure
• Volume
• Temperature

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1st Law of Thermodynamics

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Internal Energy (U)

• Has 2 components
• Kinetic energy various types of molecular and
  electron motion
• Potential energy attractive and repulsive
  interactions between atoms and molecules
• ?U U(products) U(reactants)

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Potential and Kinetic
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Calculating ?U

• ?U q w
• q heat (absorbed or released by the system)
• w work (done on or by the system)

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Example

• Calculate the overall change in internal energy
  (?U) for a system that absorbs 188 J of heat and
  does 141 J of work on its surroundings.

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Group Quiz 1

• Convert 723.01 J into calories
• SKETCH and LABEL what an exothermic and
  endothermic energy vs. time graph would look
  like.
• Calculate the overall change in internal energy
  for a system that releases 43 J in heat and has
  37 J of work done on it by its surroundings

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Section 10.3Enthalpy
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Energy and Enthalpy

• Reactions can be carried out in two ways
• In a closed container (constant volume)
• qv ?U
• In an open container (constant pressure)
• qp ? H

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Energy and Enthalpy

• Combustion of propane gas

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Enthalpy of Reaction (?H)

• ?H H(products) H(reactants)
• endothermic
• exothermic

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Enthalpy and Exo and Endo
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Thermochemical Equations

• H2O(s) ? H2O(l) ?H 6.01 kJ/mol
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l) ?H
  -890.4 kJ/mol

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Looking at the Numbers

• CH4(g) 2O2(g) ? CO2(g) 2H2O(l) ?H
  -890.4 kJ/mol
• How much energy is release from 18.4 g of methane
  being burned?
• If 924.3 kJ of energy was released, how many
  grams of water was produced?

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Properties of Enthalpy

• If you change the AMOUNTS in a balanced equation,
  you change the enthalpy the same way
• Ex if coefficients are doubled, so is the
  enthalpy
• If you reverse the equation, you reverse the sign
  of the ?H
• Ex H2O(s) ? H2O(l) ?H 6.01 kJ/mol
• H2O(l) ? H2O(s) ?H -6.01
  kJ/mol

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Section 10.4Calorimetry
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Calorimetry

• Measurement or heat changes within a system
• Using a calorimeter

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Specific Heat vs. Heat Capacity

• Specific Heat (s) amount of heat required to
  raise the temperature of 1 g of a substance by
  1C (ex liquid water is 4.184 J/(gC)
• q (s)(m)(?T)
• Heat Capacity (C) amount of heat required to
  raise the temperature of an object by 1C
• q (C)(?T)

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Example Specific Heats
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Example

• What is the amount of heat (in kJ) required to
  heat 255 g of water from 25.2 C to 90.5 C?

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Coffee-Cup Calorimetry

• Can calculate changes in heat using styrofoam
  cups and known mass of water
• Assuming constant pressure
• Therefore
• qp ms?T ?H

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Coffee Cup Calorimetry

• System reactants and products (the reaction)
• Surroundings water in calorimeter
• For an exothermic reaction
• The system loses heat
• The surroundings gain (absorb) heat

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Example

• A 30.4-g piece of unknown metal is heated up in a
  hot bath to a temperature of 92.4C. The metal
  is then placed in a calorimeter containing 100. g
  of water at 25.0C. After the calorimeter is
  capped, the temperature of the calorimeter raises
  to 27.2C. What was the specific heat of the
  unknown metal?

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Another Example

• Ex 50.0 mL of 1.00 M HCl and 50.0 mL of 1.00 M
  NaOH are mixed in a calorimeter with 100 g of
  water and capped at room temp (25C). The
  reaction reaches a max of 31.7C. What is the
  ?Hrxn?

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Group Quiz 2

• 125.0-g of a metal is heated to 100.0C. It is
  then placed into a calorimeter containing 100.0
  mL (100.0 g) of water at 25.0C and capped. The
  energy is transferred and the max temperature of
  34.1C is reached. What is the specific heat of
  the metal?

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Section 10.5Hesss Law
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Hesss Law
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Hesss Law
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Examples
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Group Quiz 3

• Given the following, determine the ?H for
• 3H2(g) O3(g) ? 3H2O(g)

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Section 10.6Standard Enthalpies of Formation
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Enthalpy of Formation

• Standard Enthalpy of Formation (?Hf) heat
  change that results when 1 mole of a compound is
  formed from its constituent elements in their
  standard states
• Standard State means stable form
• 1 atm and 25C typically
• Example O(g) (249.4), O2(g) (0), O3(g) (142.2)

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Enthalpy of Formation
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Standard Enthalpy of Reaction

• ?Hrxn enthalpy of a reaction under standard
  conditions

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An Example

• When we know reactions go to completion or can be
  done in one step, we can use a direct method
• Ex Calculate ?Hrxn for
• 2SO(g) 2/3O3(g) ? 2SO2(g)
• From Appendix 2 SO(g) (5.01), O3(g) (142.2),
  SO2(g) (-296.4)

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More on Enthalpy of Reaction

• When a reaction is too slow or side reactions
  occur, enthalpy of reaction can be calculated
  using Hesss Law

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Section 10.7Bond Enthalpy and the Stability of
Covalent Molecules
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Bond Enthalpy

• Recall when bonds are made, energy is given off
  (exo) when bonds break, energy is needed (endo)
• Bond Enthalpy the measure of stability of a
  molecule
• Enthalpy change associated with breaking a
  particular bond in 1 mole of gaseous molecules
• H2(g) ? H(g) H(g) ?H 436.4 kJ/mol

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Bond Enthalpy

• The higher the bond enthalpy, the stronger the
  bond
• The bonds in different compounds have different
  bond enthalpies
• Ex OH bond in water vs. OH bond in methanol
  are different
• Therefore, we speak of AVERAGE bond enthalpy

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Section 10.8Lattice Energy and the Stability of
Ionic Solids
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Lattice Energy

• Recall amount of energy required to convert 1
  mole of ionic solid to its constituent ions in
  the gas phase
• Ex NaCl(s) ? Na(g) Cl-(g)