Chapter 6 Thermochemistry 6'1 6'2 Notes

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Chapter 6 Thermochemistry6.1 6.2 Notes

• AP Chemistry

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Goals 35 45 aah!

• Describe energy in terms of the capacity to do
  work or produce heat.
• Understand the law of conservation of energy.
• Describe energy as a state function.
• Distinguish between exothermic and endothermic
  processes.
• Distinguish between potential and kinetic energy.
• Distinguish between heat and temperature.
• Describe energy flow in the universe between
  systems and surroundings.
• Understand the first law of thermodynamics.
• Perform energy calculations on systems with both
  heat and work energy.
• Describe the enthalpy of a system as the flow of
  heat energy at constant pressure.
• Use calorimetry to determine the heat of reaction
  at constant volume and pressure.

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The Nature of Energy

• Energy capacity to do work or release heat
• Law of Conservation of Energy energy can be
  converted from one form to another, but cant be
  created or destroyed
• The energy of the universe is constant

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Kinetic vs. Potential Energy

• Kinetic Energy - energy an object has due to its
  motion and depends upon the objects mass and
  velocity (KE(1/2)mv2)
• Potential Energy energy due to position or
  composition

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Chemical Energy

• Chemical Energy potential energy stored within
  the bonds of a substance
• System part of the universe we are focusing on
  (reactant and products in a reaction)
• Surroundings everything else in the universe
  (room reaction is taking place in)

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Heat vs. Temperature

• Heat energy transferred between two objects due
  to temperature differences between them
• Temperature measure of the average kinetic
  energy of particles in matter

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Exothermic vs. Endothermic

• Exothermic heat is evolved, or given off
  energy flows out of the system
• Endothermic heat is absorbed energy flows into
  the system
• The energy gained by the surrounding must be
  equal to the energy lost by the system

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Exothermic Reactions

• Potential energy stored in the chemical bonds is
  being converted to thermal energy via heat

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Endothermic Reactions

• Thermal energy is being converted into potential
  energy stored in the chemical bonds

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Thermodynamics

• Study of energy and the changes it undergoes
• 1st law of thermodynamics the energy of the
  universe is constant

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Internal Energy

• Sum of kinetic and potential energy of the
  particles in the system
• Energy heat work
• ?E q w
• If pressure volume work is done, then w -P?V

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Sign of q and w

• q is () if energy flows into the system
  (endothermic)
• q is (-) if energy flows out of the system
  (exothermic)
• w is (-) if system does work on surroundings
  (expanding)
• w is () if surroundings does work on system
  (compression)

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Internal Energy Example

• A system releases 125 kJ of heat and does 29 kJ
  of work. Calculate ?E for this process.

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PV Work Example

• Calculate the work associated with the expansion
  of a gas from 36 L to 57 L at a constant external
  pressure of 1.5 atm.

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Internal Energy, Heat, and Work Example

• A hot air balloon is inflated by heating the air
  inside it. The volume of the balloon increases
  from 2.00x106 L to 4.00x106 L by the addition of
  5.50x108 J of heat energy. Assuming the balloon
  expands against a constant pressure of 1 atm,
  calculate ?E for this process. (Use 1 Latm
  101.3 J)

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Enthalpy

• ?H E PV
• At constant pressure, ?H q, or the change in
  enthalpy of a system is equal to the energy flow
  as heat
• When studying reactions at constant pressure, the
  terms heat of reaction and enthalpy are
  interchangeable
• ?H ?Hproducts - ?Hreactants

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Enthalpy Example

• On average, each person in the USA consumes 110 g
  of sugar daily. If all of this sugar is assumed
  to be glucose, how much heat energy from glucose
  is transferred to one person per day at constant
  pressure?
• C6H12O6 6 O2 ? 6 CO2 6 H2O 2870 kJ

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Calorimetry

• Science of measuring heat
• Calorimeter device used to measure heat change
• Heat Capacity (C)
• C (heat absorbed/increase in temperature)
• Specific Heat Capacity (c or s) amount of heat
  necessary to raise 1 gram of a substance by 1 oC
• Molar Heat Capacity - amount of heat necessary to
  raise 1 mole of a substance by 1 oC

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Constant-Pressure vs. Constant Volume Calorimetry

• Constant-Pressure heat of reaction is equal to
  the enthalpy change
• ?H q specific heat x mass x ?temperature
• q mc?T
• Constant-Volume heat of reaction is equal to
  the internal energy change
• ?E q ?temperature x heat capacity of
  calorimeter
• q C?T

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Constant-Pressure Calorimetry Example

• A 2.00 L of 0.500 M Ba(NO3)2 solution at 22.0 oC
  is mixed with 1.00 L of 1.00 M Na2SO4 in a
  calorimeter to produce a white precipitate. The
  temperature of the solution increased to 28.1 oC.
  Assume this reaction took place at constant
  pressure and the calorimeter absorbed a
  negligible amount of heat (Density of the final
  solution is equal to 1.0 g/mL and its specific
  heat capacity is 4.18 J/(goC)). Calculate the
  enthalpy change per mole of BaSO4 formed.

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Constant-Volume Calorimetry Example

• 12.0 g of H2 gas was burned in a bomb calorimeter
  causing a temperature increase of 17.0 oC. The
  bomb calorimeter has a heat capacity of 11.3
  kJ/oC. What is the energy of combustion of
  hydrogen per mole?