CH 6: Thermochemistry

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CH 6 Thermochemistry

• Renee Y. Becker
• Valencia Community College
• CHM 1045

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Energy

• Energy is the capacity to do work, or supply
  heat.
• Energy Work Heat
• Kinetic Energy is the energy of motion.
• EK 1/2 mv2 (1 Joule 1 kg?m2/s2)
• (1 calorie 4.184 J)
• Potential Energy is stored energy.

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Ek Ep
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Example 1 KE

• Which of the following has the greatest kinetic
  energy?
• A 12 kg toy car moving at 5 mph?
• A 12 kg toy car standing at the top of a large
  hill?

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Energy

• Thermal Energy is the kinetic energy of molecular
  motion
• Thermal energy is proportional to the temperature
  in degrees Kelvin. Ethermal ? T(K)
• Heat is the amount of thermal energy transferred
  between two objects at different temperatures.

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• In an experiment Reactants and products are the
  system everything else is the surroundings.
• Energy flow from the system to the surroundings
  has a negative sign (loss of energy). (-?E or -
  ?H)
• Energy flow from the surroundings to the system
  has a positive sign (gain of energy). (?E or
  ?H)

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• The law of the conservation of energy Energy
  cannot be created or destroyed.
• The energy of an isolated system must be
  constant.
• The energy change in a system equals the work
  done on the system the heat added.
• DE Efinal Einitial E2 E1 q w
• q heat, w work

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• Pressure is the force per unit area.
• (1 N/m2 1 Pa)
• (1 atm 101,325 Pa)
• Work is a force (F) that produces an objects
  movement, times the distance moved (d)
• Work Force x Distance

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The expansion in volume that occurs during a
reaction forces the piston outward against
atmospheric pressure, P. Work -atmospheric
pressure area of piston distance piston moves
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Example 2 Work

• How much work is done (in kilojoules), and in
  which direction, as a result of the following
  reaction?

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• The amount of heat exchanged between the system
  and the surroundings is given the symbol q.
• q DE PDV
• At constant volume (DV 0) qv DE
• At constant pressure qp DE PDV DH
• Enthalpy change DH Hproducts Hreactants

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Example 3 Work

• The explosion of 2.00 mol of solid TNT with a
  volume of approximately 0.274 L produces gases
  with a volume of 489 L at room temperature. How
  much PV (in kilojoules) work is done during the
  explosion? Assume P 1 atm, T 25C.
• 2 C7H5N3O6(s) ? 12 CO(g) 5 H2(g) 3 N2(g) 2
  C(s)

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• Enthalpies of Physical Change

Enthalpy is a state function, the enthalpy change
from solid to vapor does not depend on the path
taken between the two states.
?Hsubl ?Hfusion ?Hvap
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• Enthalpies of Chemical Change Often called heats
  of reaction (DHreaction).
• Endothermic Heat flows into the system from the
  surroundings and DH has a positive sign.
• Exothermic Heat flows out of the system into the
  surroundings and DH has a negative sign.

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• Reversing a reaction changes the sign of DH for a
  reaction.
• C3H8(g) 5 O2(g) ? 3 CO2(g) 4 H2O(l) DH
  2219 kJ
• 3 CO2(g) 4 H2O(l) ? C3H8(g) 5 O2(g) DH
  2219 kJ
• Multiplying a reaction increases DH by the same
  factor.
• 3 C3H8(g) 15 O2(g) ? 9 CO2(g) 12 H2O(l)
  DH 3(-2219) kJ
• DH -6657 kJ

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Example 4 Heat

• How much heat (in kilojoules) is evolved or
  absorbed in each of the following reactions?
• a) Burning of 15.5 g of propane C3H8(g) 5
  O2(g) ? 3 CO2(g) 4 H2O(l)
• DH 2219 kJ/mole
• b) Reaction of 4.88 g of barium hydroxide
  octahydrate with ammonium chloride
• Ba(OH)28 H2O(s) 2 NH4Cl(s) ? BaCl2(aq) 2
  NH3(aq) 10 H2O(l)
• DH 80.3 kJ/mole

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• Thermodynamic Standard State Most stable form of
  a substance at 1 atm pressure and 25C 1 M
  concentration for all substances in solution.
• These are indicated by a superscript to the
  symbol of the quantity reported.
• Standard enthalpy change is indicated by the
  symbol DH.

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Example 5

• Is an endothermic reaction a favorable process
  thermodynamically speaking?
• Yes
• No

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Hesss Law

• Hesss Law The overall enthalpy change for a
  reaction is equal to the sum of the enthalpy
  changes for the individual steps in the
  reaction.(not a physical change, chemical change)
• 3 H2(g) N2(g) ? 2 NH3(g) DH 92.2 kJ

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• Reactants and products in individual steps can be
  added and subtracted to determine the overall
  equation.
• (1) 2 H2(g) N2(g) ? N2H4(g)
  DH1 ?
• (2) N2H4(g) H2(g) ? 2 NH3(g)
  DH2 187.6 kJ
• (3) 3 H2(g) N2(g) ? 2 NH3(g)
  DH3 92.2 kJ
• DH1 DH2 DHreaction
• Then DH1 DHreaction - DH2
• DH1 DH3 DH2 (92.2 kJ) (187.6 kJ)
  95.4 kJ

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Example 6 Hesss Law

• The industrial degreasing solvent methylene
  chloride (CH2Cl2, dichloromethane) is prepared
  from methane by reaction with chlorine
• CH4(g) 2 Cl2(g) ?CH2Cl2(g) 2 HCl(g)
• Use the following data to calculate DH (in
  kilojoules) for the above reaction
• CH4(g) Cl2(g) ? CH3Cl(g) HCl(g)
• DH 98.3 kJ
• CH3Cl(g) Cl2(g) ? CH2Cl2(g) HCl(g)
• DH 104 kJ

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• Standard Heats of Formation (DHf) The enthalpy
  change for the formation of 1 mole of substance
  in its standard state from its constituent
  elements in their standard states.
• The standard heat of formation for any element in
  its standard state is defined as being ZERO.
• DHf 0 for an element in its standard state

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Standard Heats of Formation

• Calculating DH for a reaction
• DH DHf (Products) DHf (Reactants)
• For a balanced equation, each heat of formation
  must be multiplied by the stoichiometric
  coefficient.
• aA bB ? cC dD
• DH cDHf (C) dDHf (D) aDHf (A)
  bDHf (B)

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Standard Heats of Formation
Some Heats of Formation, ?Hf (kJ/mol)
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Example 7 Standard heat of formation

• Calculate DH (in kilojoules) for the reaction of
  ammonia with O2 to yield nitric oxide (NO) and
  H2O(g), a step in the Ostwald process for the
  commercial production of nitric acid.

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Example 8 Standard heat of formation

• Calculate DH (in kilojoules) for the
  photosynthesis of glucose and O2 from CO2 and
  liquid water, a reaction carried out by all green
  plants.

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Example 9

• Which of the following would indicate an
  endothermic reaction? Why?
• -?H
• ?H

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Heat of Phase Transitions from ?H?f

• Calculate the heat of vaporization, ?H?vap of
  water, using standard enthalpies of formation
• ?H?f
• H2O(g) -241.8 kJ/mol
• H2O(l) -285.8 kJ/mol

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Calorimetry and Heat Capacity

• Calorimetry is the science of measuring heat
  changes (q) for chemical reactions. There are
  two types of calorimeters
• Bomb Calorimetry A bomb calorimeter measures
  the heat change at constant volume such that q
  DE.
• Constant Pressure Calorimetry A constant
  pressure calorimeter measures the heat change at
  constant pressure such that q DH.

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Constant Pressure
Bomb
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Calorimetry and Heat Capacity

• Heat capacity (C) is the amount of heat required
  to raise the temperature of an object or
  substance a given amount.
• Specific Heat The amount of heat required to
  raise the temperature of 1.00 g of substance by
  1.00C.
• q s x m x ?t
• q heat required (energy)
• s specific heat
• m mass in grams
• ?t Tf - Ti

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Calorimetry and Heat Capacity

• Molar Heat The amount of heat required to raise
  the temperature of 1.00 mole of substance by
  1.00C.
• q MH x n x ?t
• q heat required (energy)
• MH molar heat
• n moles
• ?t Tf - Ti

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Example 10 Specific Heat

• What is the specific heat of lead if it takes 96
  J to raise the temperature of a 75 g block by
  10.0C?

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Example 11 Specific Heat

• How much energy (in J) does it take to increase
  the temperature of 12.8 g of Gold from 56?C to
  85?C?

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Example 12 Molar Heat

• How much energy (in J) does it take to increase
  the temperature of 1.45 x104 moles of water from
  69?C to 94?C?