Energy Changes in Chemical Reactions

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Chapter 10

• Energy Changes in Chemical Reactions

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Energy Changes in Chemical Rxns

• Most reactions give off or absorb energy
• Energy is the capacity to do work or supply heat.
• Heat transfer of thermal (kinetic) energy
  between two systems at different temperatures
  (from hot to cold)

Metal bar in water
Metal bar drilled
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Types of Energy

• Work (w) energy transfer when forces are applied
  to a system
• Heat (q) energy transferred from a hot object to
  a cold one
• Radiant energy? heat from the sun
• Thermal energy ? associated with motion of
  particles
• Potential energy ? energy associated with
  objects position or substances chemical bonds
• Kinetic energy ? energy associated with objects
  motion

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Heat versus Temperature

• Describe the difference between the two.
• SI unit of energy J
• 1 watt 1 J/s, so a 100 Watt bulb uses 100 J
  each second
• We often use the unit of kJ to refer to chemical
  heat exchanges in a reaction. 1 kJ 1000 J
• Energy is also reported in calories
• Amount of energy needed to raise 1 gram of water
  by 1oC
• 1 cal 4.184 J 1 Cal 4184 J
• Cal (or kcal) is used on food labels

Molecular heat transfer
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Energy and Energy Conservation

• Heat form of energy transferred from object at
  higher temperature to one at lower temperature
  (from hot object to cold object)
• Thermochemistry study of heat changes in
  chemical reactions, in part to predict whether or
  not a reaction will occur
• Thermodynamics study of heat and its
  transformations
• First Law of Thermodynamics Energy can be
  converted from one form to another but cannot be
  created or destroyed

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System and Surroundings

• System loses heat (negative) gains heat
  (positive)

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Endothermic vs Exothermic

• Endothermic reaction q is positive (q gt 0)
• Reaction (system) absorbs heat
• Surroundings feel cooler
• Exothermic reaction q is negative (q lt 0)
• Reaction releases heat
• Surroundings feel warmer

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Enthalpies of Reaction

• Determine if the following processes are
  endothermic or exothermic
• Combustion of methane
• Reacting Ba(OH)2 with NH4Cl
• Neutralization of HCl
• Melting
• CaCO3 (s) ? CaO (s) CO2 (g)

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Answers to Enthalpies of Reaction

• Combustion of methane exothermic
• Reacting Ba(OH)2 with NH4Cl endothermic
• Neutralization of HCl exothermic
• Melting endothermic
• CaCO3 (s) ? CaO (s) CO2 (g) endothermic
• Combustion, neutralization, and combination
  reactions tend to be exothermic
• Decomposition reactions tend to be endothermic
• Melting, boiling, and sublimation are endothermic

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Applications of heat emission/absorption
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Specific Heat and Heat Capacity

• Specific heat (sp. ht.) amount of heat required
  to raise 1 gram of substance by 1oC
• Use mass, specific heat, and DT to calculate the
  amount of heat gained or lost
• q msDT ? ms C ? q CDT
• Heat capacity (C) amount of heat required to
  raise the temperature of a given quantity of a
  substance by 1oC C q / DT J / oC
• Molar heat capacity (Cm) amount of heat that can
  be absorbed by 1 mole of material when
  temperature increases 1oC q (Cm) x (moles of
  substance) x (DT) J / mol oC

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Specific Heat and Heat Capacity
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Practice Problem

• Calculate the amount of heat transferred when 250
  g of H2O (with a specific heat of 4.184 J/goC)
  is heated from 22oC to 98oC.
• q msDT
• Is heat being put into the system or given off by
  the system?
• If a piece of hot metal is placed in cold water,
  what gains heat and what loses heat? Which one
  will have a positive q value and which will have
  a negative q value?

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Practice Problem

• 34.8 g of an unknown metal at 25.2oC is mixed
  with 60.1 g of H2O at 96.2 oC (sp. ht. 4.184
  J/goC). The final temperature of the system
  comes to 88.4oC. Identify the unknown metal.
• Specific heats of metals
• Al 0.897 J/goC
• Fe 0.449 J/goC
• Cu 0.386 J/goC
• Sn 0.228J/goC

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Calorimetry and Heat Capacity

• Heat changes in a reaction can be determined by
  measuring the heat flow at constant pressure
• Apparatus to do this is called a calorimeter.
• Heat evolved by a reaction is absorbed by water
  heat capacity of calorimeter is the heat capacity
  of water.

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Example

• A 28.2 gram sample of nickel is heated to 99.8oC
  and placed in a coffee cup calorimeter containing
  150.0 grams of water at 23.5oC. After the metal
  cools, the final temperature of the metal and
  water is 25.0oC.
• qabsorbed qreleased 0
• Which substance absorbed heat?
• Which substance released heat?
• Calculate the heat absorbed by the substance you
  indicated above.

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Group Quiz 25

• A hot piece of copper (at 98.7oC, specific heat
  0.385 J/goC) weighs 34.6486 g. When placed in
  room temperature water, it is calculated that
  915.1 J of heat are released by the metal.
• What gains heat?
• What loses heat?
• What is the final temperature of the metal?
• Watch signs!!!!

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Enthalpies of Physical/Chemical Changes

• Enthalpy (H) describes heat flow into and out of
  a system under constant pressure
• Enthalpy (a measure of energy) is heat
  transferred per mole of substance.
• At constant pressure,
• qp DH Hproducts Hreactants
• DH gt 0 ? endothermic (net absorption of energy
  from environment products have more internal
  energy)
• DH lt 0 ? exothermic (net loss of energy to
  environment reactants have more internal energy)

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Heating a Pure Substance (Water)

• Why does T become constant during melting and
  evaporating?
• Melting, vaporization, and sublimation are
  endothermic
• We can calculate total heat needed to convert a
  15 gram piece of ice at -20oC to steam at 120oC.

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Enthalpies of Phase Changes

• Heat of fusion (DHfus) Amount of heat required
  to melt (solid ? liquid)
• Heat of vaporization (DHvap) Amount of heat
  required to evaporate (liquid ? gas)
• Heat of sublimation (DHsub) Amount of heat
  required to sublime (solid ? gas)
• Why are there no values for DHfreezing,
  DHcondendsation, or DHdeposition?

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Thermochemical Equations

• Shows both mass and enthalpy relationships
• 2Al (s) Fe2O3 (s) ? 2Fe (s) Al2O3 (s)
  DHo -852 kJ
• Amount of heat given off depends on amount of
  material
• 852 kJ of heat are released for every 2 mol Al, 1
  mol Fe2O3, 2 mol Fe, and 1 mol Al2O3

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Thermochemical Equations

• 2Al (s) Fe2O3 (s) ? 2Fe (s) Al2O3 (s)
  DHo -852 kJ
• How much heat is released if 10.0 grams of Fe2O3
  reacts with excess Al?
• What if we reversed the reaction?
• Heat would have to be put in to make the reaction
  proceed
• 2Fe (s) Al2O3 (s) ? 2Al (s) Fe2O3 (s) DHo
  852 kJ

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Hesss Law

• If a compound cannot be directly synthesized from
  its elements, we can add the enthalpies of
  multiple reactions to calculate the enthalpy of
  reaction in question.
• Hesss Law change in enthalpy is the same
  whether the reaction occurs in one step or in a
  series of steps
• Look at direction of reaction and amount of
  reactants/products

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Hesss Law

• Value changes sign with direction

Figure 8.5
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Hesss Law

• Values of enthalpy change
• For a reaction in the reverse direction, enthalpy
  is numerically equal but opposite in sign
• Reverse direction, heat flow changes endothermic
  becomes exothermic (and vice versa) sign of DH
  changes
• Proportional to the amount of reactant consumed
• Twice as many moles twice as much heat half as
  many moles half as much heat
• DHT DH1 DH2 DH3 .

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Enthalpy of Chemical Reaction

• Thermochemical equation
• H2(g) I2(s) ? 2HI(g) DH 53.00 kJ
• Two possible changes
• Reverse the equation
• 2HI(g) ? H2(g) I2(s) DH -53.00 kJ
• Double the amount of material
• 2H2(g) 2I2(s) ? 4HI(g) DH 106.00 kJ

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Hesss Law

• Calculate DHo for
• 2NO (g) O2 (g) ? N2O4 (g) DHo ?
• N2O4 (g) ? 2NO2 (g) DHo 57.2 kJ
• NO (g) ½ O2 (g) ? NO2 (g) DHo -57.0 kJ

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Hesss Law

• We can use known values of DHo to calculate
  unknown values for other reactions
• P4 (s) 3 O2 (g) ? P4O6 (s) DH -1640.1 kJ
• P4 (s) 5 O2 (g) ? P4O10 (s) DH -2940.1 kJ
• What is DHo for the following reaction?P4O6 (s)
  2 O2 (g) ? P4O10 (s) DH ?

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Hesss Law
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Hesss Law

• Given
• 2NH3(g) ? N2H4(l) H2(g) DH 54 kJ
• N2(g) H2(g) ? NH3(g) DH -69 kJ
• CH4O(l) ? CH2O(g) H2(g) DH -195 kJ
• Find the enthalpy for the following reaction
• N2H4(l) CH4O(l) ? CH2O (g) N2(g) 3H2(g)
• DH ? kJ

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Group Quiz 26

• Given the following equations
• 2CO2 (g) ? O2 (g) 2CO (g) DH 566.0 kJ
• ½ N2 (g) ½ O2 (g) ? NO (g) DH 90.3 kJ
• Calculate the enthalpy change for
• 2CO (g) 2NO (g) ? 2CO2 (g) N2 (g) DH ?

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Standard Heats of Formation

• Standard heat of formation (DHof) heat needed
  to make 1 mole of a substance from its stable
  elements in their standard states
• DHof 0 for a stable (naturally occurring)
  element
• Which of these have DHof 0?
• CO(g), Cu(s), Br2(l), Cl(g), O2(g), O3(g), O2(s),
  P4(s)
• Do the following equations represent standard
  enthalpies of formation? Why or why not?
• 2Ag (l) Cl2 (g) ? 2AgCl (s)
• Ca (s) F2 (g) ? CaF2 (s)

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Standard Enthalpies of Formation

• Can use measured enthalpies of formation to
  determine the enthalpy of a reaction (use
  Appendix B in back of book)
• DHorxn SnDHof (products) SnDHof (reactants)
• S sum n number of moles (coefficients)
• Direct calculation of enthalpy of reaction if the
  reactants are all in elemental form
• Sr (s) Cl2 (g) ? SrCl2 (g)
• DHorxn DHof (SrCl2) DHof (Sr) DHof
  (Cl2) -828.4 kJ/mol

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Standard Enthalpies of Formation

• Some Common Substances (25oC)

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Heats of Formation

• DHorxn S DHof,products - S DHof,reactants
• Calculate values of DHo for the following rxns
• 1) CaCO3 (s) ? CaO (s) CO2 (g)
• 2) 2C6H6 (l) 15O2 (g) ? 12CO2 (g) 6H2O (l)
• DHof values
• CaCO3 -1207.1 kJ/mol CaO -635.5 kJ/mol
  CO2 -393.5 kJ/mol C6H6 49.0 kJ/mol
  H2O(l) -285.8 kJ/mol

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Group Quiz 27

• Use Standard Heat of Formation values to
  calculate the enthalpy of reaction for
• C6H12O6(s) ? C2H5OH(l) CO2(g)
• Hint Is the equation balanced?
• DHof (C6H12O6(s)) -1260.0 kJ/mol
• DHof (C2H5OH(l)) -277.7 kJ/mol
• DHof (CO2(g)) -393.5 kJ/mol

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Bond Dissociation Energies

• Bond Dissociation Energy (or Bond Energy, BE)
  energy required to break a bond in 1 mole of a
  gaseous molecule
• Reactions generally proceed to form compounds
  with more stable (stronger) bonds (greater bond
  energy)

H2 Bond Energy
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Bond Dissociation Energies

• Bond energies vary somewhat from one mole- cule
  to another so we use average bond dissociation
  energy (D)
• H-OH 502 kJ/mol Avg O-H 453
• H-O 427 kJ/mol kJ/mol
• H-OOH 431 kJ/mol

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Bond Dissociation Energies
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Bond Dissociation Energies

• DHorxn SBE (reactants) - SBE (products)
  endothermic
  exothermic
• energy input
  energy released
• SBE(react) gt SBE(prod) ? endothermic
• SBE(react) lt SBE(prod) ? exothermic
• Use only when heats of formation are not
  available, since bond energies are average values
  for gaseous molecules

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Heats of Reaction

• Use bond energies to calculate the enthalpy
  change for the following reaction N2(g)
  3H2(g) ? 2NH3(g)
• DHrxn BEN ? N 3BEH-H -6BEN-H
• DHrxn 945 3(436) 6(390) -87 kJ
• measured value -92.2 kJ
• Why are the calculated and measured values
  different?

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Heats of Reaction

• Use bond energies to calculate the enthalpy
  change for the decomposition of nitrogen
  trichloride NCl3 (g) ? N2 (g)
  Cl2 (g)
• How many distinct bond types are there in each
  molecule?
• How many of each bond type do we need to
  calculate DHrxn?
• BE(N-Cl) 200 kJ/mol
• BE(NN) 945 kJ/mol
• BE(Cl-Cl) 243 kJ/mol

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Answer

• 6(N-Cl) -1(N N) -3(Cl-Cl)
• 6(200) -(945) -3(243) -474 kJ

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Thermochemistry Calculation Summary

• Use q msDT (s J/goC)
• If given mass of reactant, convert to moles and
  multiply by enthalpy to find total heat
  transferred
• If given multiple equations with enthalpies, use
  Hesss Law
• If given DHof values products reactants
• If given bond energy (BE) values
• reactants -products

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Practice Problems

• Identify how to set up the following problems
• Calculate the DHo of reaction for
• C3H8 (g) 5O2 (g) ? 3CO2 (g) 4H2O (l)
• DHof C3H8(g) -103.95 kJ/mol DHof CO2(g) -393.5
  kJ/mol DHof H2O(l) -285.8 kJ/mol
• 8750 J of heat are applied to a 170 g sample of
  metal, causing a 56oC increase in its
  temperature.  What is the specific heat of the
  metal? Which metal is it?

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Practice Problems

• C2H4(g ) 6F2(g) ? 2CF4(g) 4HF(g) DHo ?
• H2 (g) F2 (g) ? 2HF (g) DHo -537 kJ
• C (s) 2F2 (g) ? CF4 (g) DHo -680 kJ
• 2C (s) 2H2 (g) ? C2H4 (g) DHo 52.3 kJ
• Use average bond energies to determine the
  enthalpy of the following reaction
• CH4 (g) Cl2 (g) ? CH3Cl (g) HCl (g)
• (BEC-Cl 328 kJ/mol)

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The End