Thermochemistry

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Thermochemistry

• Topic 5

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Words you need to know

• Thermochemistry
• Thermodynamics
• Energy
• Heat
• calorie/Calorie
• Joule
• Energy transfers occur btwn the system and its
  surroundings

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calorie v. Calorie

• calorie is the amount of heat energy needed to
  raise the temperature of 1 g of water 1oC
• Calorie is 1000 calories
• Food energy is measured in Calories
• Fats are 9 Calories/g
• Carbs and proteins are 4 Calories/g

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1st Law of Thermodynamics

• aka Law of conservation of energy
• Energy in the universe is constant, cannot be
  created or destroyed
• Energy can be converted to different forms

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Exothermic Processes

• System transfers heat to its surroundings
• Temperature of the surroundings increases
• feels hot
• Potential energy of the system is converted to
  heat energy that is released (system does not
  cool off)

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Endothermic Processes

• System absorbs heat from the surroundings
• Temperature of the surroundings decreases
• feels cool
• Heat energy from the surroundings is converted
  into potential energy in the system

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Practice

1. Imagine an ice cube melting in your hand. Is the
   melting of ice endothermic or exothermic?
   Explain using the terms system and surroundings
2. Imagine warming your hands near a campfire. Is a
   campfire an endothermic or exothermic process?
   Explain.

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Enthalpy (H)

• Property of a system that explains heat flow
  between the system and its surroundings (constant
  P)
• State function (middle steps dont matter, just
  the beginning and end of a reaction)

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Enthalpy changes (DH)

• Defined as heat absorbed by the system during a
  physical or chemical change
• DH is positive for endothermic rxns (because heat
  is absorbed by the system)
• DH is negative for exothermic rxns (because heat
  is released by the system)
• Expressed in kJ or kJ/mol
• Magnitude of DH is directly proportional to moles
  of reactants and products

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Types of Enthalpy Changes

• DHrxn heat of reaction of any chemical reaction
• DHcomb heat of combustion for combustion
  reactions (rxns with O2) only
• DHfus heat of fusion when a solid melts

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More Types

• DHvap heat of vaporization when a liquid
  vaporizes
• DHBDE bond dissociation energy or the heat
  required to break a bond
• DHf heat of formation or the heat change when a
  compound is formed from its elements
• DHsols heat of solution or the heat change when
  a solute dissolves in a solvent

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Practice

• How many kJ of heat are absorbed when 25.0 g of
  methane burn in air? Methane has a DHcomb of
  -802 kJ/mol.

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Reversible Reactions

• If DH is positive for the forward reaction, then
  it will be equal in size but opposite in sign for
  the reverse reaction
• H2O (s) ? H2O (l) DH 6.0 kJ/mol
• H2O (l) ? H2O (s) DH -6.0 kJ/mol

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Keep in mind

• The sign for DH is positive or negative depending
  on the direction of energy flow
• The sign does NOT indicate a positive or negative
  value for energy
• What scientific law requires that the magnitude
  of the heat change for a forward and reverse
  reactions be the same with opposite signs?
  Explain.

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DH and Phase Changes

• DH for phase changes from solid to liquid and
  liquid to gas are ALWAYS positive
  (endothermic/absorb heat)
• DH for phase changes from gas to liquid and
  liquid to solid are ALWAYS negative
  (exothermic/release heat)
• Why is this? What is the sign of DH for the
  process of sublimation? Deposition?

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Calorimetry

• aka. measurement of heat flow
• Calorimeter measure heat flow
• Heat capacity (C) amount of heat required to
  raise the temperature of any object 1oC.
  Expressed in J/K or J/oC
• We will more often use specific heat capacity
  (Cp) which is the capacity of 1 g of a substance
• Water has a Cp of 1 cal/gK or 4.184 J/gK

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Explain

• Why is the Cp of water 1 cal/gK? What is the
  conversion between calories and joules?

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Example

• What is the molar heat of combustion of liquid
  ethanol if the combustion of 9.03 grams of
  ethanol causes a calorimeter to increase in
  temperature by 3.54 K? The heat capacity of the
  calorimeter is 75.8 kJ/K.

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Hesss Law

• If a reaction is carried out in a series of steps
  the DH of the overall reaction is equal to the
  sum of the DHs for each individual step.
• Useful for determining DH for reactions that are
  difficult to measure directly, like sulfur
  trioxide

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Hesss Law with SO3

• The overall reaction is
• 2S (s) 3O2 (g) ? 2SO3 (g)
• The reaction occurs in 2 measurable steps
• S (s) O2 (g) ? SO2 (g) DHrxn -269.9 kJ
• 2SO2 (g) O2 (g) ? 2SO3 (g) DHrxn -196.6 kJ
• To get the total DH for the reaction, manipulate
  the equation steps like an algebraic equation.
• Whatever you do to the reaction, you must also do
  to DH.

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Enthalpies of Formation

• Reaction that produces 1 mole of a substance from
  its constituent elements in their most stable
  thermodynamic state
• To form 1 mole of HI, the equation looks like
  this
• ½ H2 ½ I2 ? HI DH 25.94 kJ

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Standard Heat of Formation (DHof)

• Heat absorbed when 1 mole of a substance is
  formed from its elements in their standard states
  at 25oC and 1 atm.
• There is a BIG table in the back of your book
  listing standard heats of formation.

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Example

• Write the thermochemical equation associated with
  the standard heat of formation of AlCl3. What is
  the DHof for this equation?

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Things to know

• The table of standard heats of formation includes
  elements, ions, and compounds.
• The DHof of pure elements is always zero.

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One last thing

• Hesss law allows us to calculate the DHorxn for
  just about any reaction.
• Breaking the overall reaction into the formation
  reactions for both products and reactants and
  then putting them all together like this
• DHorxn SnDHof products SmDHof reactants

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Example

• Calculate the standard enthalpy change for the
  combustion of 1 mole of liquid ethanol. Tip
  carefully watch the signs of DHof.

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Another Example

• What is
• The enthalpy of sublimation of solid calcium?
• The heat of solution of gaseous ammonia?
• The bond dissociation energy of hydrogen gas?
• The heat change when gaseous bromine condenses to
  a liquid?
• Write chemical equations to illustrate your
  answers.