Thermochemistry

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Thermochemistry
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Heat in Chemical Reactions

• Most chemical reactions involve energy to break
  or form bonds.
• Thermochemistry In a chemical reaction, the
  study of changes in heat.
• Heat The energy transferred from one object to
  another due to a difference in temperature.

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How can I use the mole ratio in these problems?

• Because the heats of chemical reactions are
  expressed in kJ/mol, these amounts can be used in
  stoichiometric problems as if it were a mole
  ratio!

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Sample Problems

• Calculate the heat required to melt 25.7 g of
  solid methanol at its melting point. The ?Hfus
  of methanol is 3.22 kJ/mol.

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Sample Problem

• What mass of methane must be burned in order to
  liberate 12.880 kJ of heat? The ?Hcomb of
  methane is -891 kJ/mol.

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Heat and Enthalpy Changes

• Enthalpy the heat content of a system at
  constant pressure
• Represented as H
• Depends on temperature, physical state, and
  composition
• System must be at a constant pressure throughout
  the change

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Enthalpy Change

• ?H SHproducts SHreactants
• Exothermic
• Heat is released
• ?H is NEGATIVE energy is released
• Reaction feels warm
• Extra energy shown on product side of equation
• Endothermic
• Heat is absorbed during a reaction
• ?H is POSITIVE energy is absorbed
• Reaction feels cold
• Extra energy shown on reactant side of equation

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Enthalpy Change Cont.

• the amount of heat that rxn absorbs or releases
  depends on the conditions under which the rxn is
  carried out
• -i.e. temp, pressure, physical states of
  reactants and products
• Use 1 atmosphere as standard pressure and 25
  degrees C as temp.(room temp) for reporting
  enthalpy changes

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Enthalpy Change Cont.

• the standard state of a substance is its pure
  form at standard pressure
• -pure elements must be present in its most stable
  form at standard pressure (standard state)
• Standard enthalpy change- an enthalpy change that
  is measured when reactants in their standard
  states change to products in their standard
  states

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Enthalpy Change Cont.

• Denoted with a superscript 0 (?H degrees)
• If the rxn occurs at temps other than 25 degrees
  C, the standard enthalpy change takes into
  account the heat involved in restoring the
  products to standard conditions after the
  reaction is complete
• The amount of heat that is absorbed or released
  in a rxn depends on the of moles of reactants
  involved

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Sample Problem

• Calculate the ?H of the following reaction
• 2SO2 (g) O2 (g) ? 2SO3 (g)

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Hesss Law

• States that the total enthalpy change for a
  chemical or physical change is the same whether
  it takes one step or several steps
• If a series of rxns are added together the
  enthalpy change for the net rxn will be the sum
  of the enthalpy changes for the individual steps

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Rules for Manipulating Reactions

• If the coefficients of an equation are multiplied
  by a factor, the enthalpy change is multiplied by
  the same factor
• If an equation is reversed, the sign of ?H is
  reversed also

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Sample Problem

• Use the thermochemical equations a and b to
  determine ?H for the decomposition of hydrogen
  peroxide.
• 2H2O2(l) ? 2H2O(l) O2(g)
• 2H2(g) O2(g) ? 2H2O (l) ?H -572 kJ
• H2(g) O2(g) ? H2O2(l) ?H -188kJ

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Sample Problem

• Use reactions a and b to determine ?H for the
  following reaction
• 2CO(g) 2NO(g) ? 2CO2(g) N2(g)
• a. 2CO(g) O2(g) ? 2CO2(g) ?H -566.0 kJ
• b. N2(g) O2(g) ? 2NO(g) ?H 180.6 kJ

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Sample Problem

• Use reactions a, b and c to determine ?H for the
  following reaction.
• H2S(g) 4F2(g) ? 2HF(g) SF6(g)
• a. 1/2 H2(g) 1/2 F2(g) ? HF(g) ?H -273 kJ
• b. S(s) 3F2 (g)? SF6(g) ?H -1220 kJ
• c. H2(g) S(s) ? H2S(g) ?H -21 kJ

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Heat in Physical Changes

• Changes in state are reversible processes that
  can be reversed by adjusting the temperature
• Each change in state requires an energy transfer

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Heat of Fusion

• Symbol ?Hfus
• Defined as heat necessary to convert one mole of
  a solid to a liquid at constant temperature
• ?H mol substance x ?Hfus
• Total heat change the number of moles of a
  substance x its heat of fusion
• Heat of fusion is measured in kJ/mol

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Heat of Solidification

• Symbol ?Hsolid
• Defined as heat necessary to convert one mole of
  liquid to a solid at a constant temperature
• ?H mol substance x ?Hsolid
• Total heat change the number of moles of a
  substance x its heat of solidification
• Heat of solidification is measured in kJ/mol

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Heat of Vaporization

• Symbol ?Hvap
• Defined as heat necessary to vaporize one mole
  of a given liquid
• ?H mol substance x ?Hvap
• Total heat change the number of moles of a
  substance x its heat of vaporization
• Heat of vaporization is measured in kJ/mol

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Heat of Condensation

• Symbol ?Hcond
• Defined as heat necessary to condense one mole
  of a given gas
• ?H mol substance x ?Hcond
• Total heat change the number of moles of a
  substance x its heat of condensation
• Heat of condensation is measured in kJ/mol

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Heat of Solution

• Symbol ?Hsoln
• Defined as heat necessary to dissolve one mole of
  a given substance
• ?H mol substance x ?Hsoln
• Total heat change the number of moles of a
  substance x its heat of solution
• Heat of solution is measured in kJ/mol

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Heat of Formation

• Symbol ?Hf
• Defined as heat necessary to form one mole of a
  compound from its elements
• ?H mol substance x ?Hf
• Total heat change the number of moles of a
  substance x its heat of formation
• Heat of formation is measured in kJ/mol

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Specific Heat

• Relates temperature changes to heat changes
• Defined as the amount of heat energy required to
  increase the temperature of one gram of a
  substance by one degree Celsius
• Specific heat -Varies if pressure and temperature
  are not kept constant
• A physical property
• Varies depending on the substance
• Symbol is Cp, units are J/gC

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Specific Heat

• Must be measured!
• Substances with low specific heats require less
  energy to feel hot than those with high specific
  heats
• Specific heat can be used to calculate changes in
  heat

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Substance Specific Heat J/gC
Water (liquid) 4.184
Water (solid) 2.03
Water (steam) 2.01
Ethanol (liquid) 2.44
Aluminum (solid) 0.897
Granite (solid) 0.803
Iron (solid) 0.449
Lead (solid) 0.129
Silver (solid) 0.235
Gold (solid) 0.129
Copper (solid) 0.385
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Heat Calculations

• Change in heat of a substance can be calculated
  using the following equation
• q m?TCp
• q change in heat
• m mass of the substance
• ?T change in temperature of the substance
• Cp specific heat of the substance

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Total Energy Changes

• The amount of heat (q) involved in a reaction is
  positive () if the sample warms up. The sample
  is gaining heat.
• The amount of heat (q) involved in a reaction is
  negative (-) if the sample cools off. The sample
  is releasing, or losing, heat.

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Specific Heat Problems

• Use the sheet which has the chart on it to find
  the specific heat of the element or substance in
  the problem
• Solve using algebra and the equation
• q m?TCp

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Sample Problems
If the temperature of 34.4 g of ethanol increases
from 25.0 ?C to 78.8 ?C, how much heat has been
absorbed by the ethanol? The specific heat of
ethanol is 2.44 J/(g??C)
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Heat in a Chemical Reaction

• Energy can be converted into other forms
• Measures of changes in heat energy can be made in
  a calorimeter
• Changes in heat energy can be used to calculate
  specific heat
• These heats refer to the total flow of energy
  during a chemical change

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Sample Problems
A 4.50 g nugget of pure gold absorbed 276 J of
heat. What was the final temperature of the gold
if the initial temperature was 25 ?C ? The
specific heat of gold is 0.129 J/(g??C).
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Sample Problems
A 155-g sample of an unknown substance was heated
from 25.0?C to 40.0 ?C. In the process, the
substance absorbed 5696 J of energy. What is the
specific heat of the substance?
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Calorimeter Questions

• Transfer of heat is measured by measuring the
  difference in temperature transferred to water
  from an object
• Specific heat of water (4.184 J/gC) and its mass
  is used to solve the problem.

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Sample Problem

• A piece of metal is placed in a calorimeter, and
  causes the 335 g of water to increase in
  temperature from 21.0C to 50.1C. What is the
  amount of energy released by the piece of metal?

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Activation Energy

• Defined as the minimum amount of energy that must
  be supplied to a system to start a chemical
  change.
• Endothermic reactions must have a source from
  which to draw their energy (usually their
  surroundings)

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Spontaneity

• Spontaneous change a change that proceeds on
  its own, without any outside intervention
• Occurs primarily in one direction
• Does not mean the reaction will occur quickly

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Spontaneity

• Some can occur when a small amount of energy is
  added to the system
• Some can be reversed if conditions change
• The more energy released, the lower enthalpy is
  and therefore the more likely the reaction will
  be spontaneous

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Entropy and Stability

• Entropy (S) a measure of the randomness or
  disorder of the system
• The tendency of nature is toward more disorder
• Its effects increase with temperature
• Measured in units of J/K

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Entropy

• ?S Sproducts Sreactants
• If entropy of a system increases during a
  reaction or process, Sproducts ? Sreactants and
  ?S is positive
• If entropy of a system decreases,
  Sproducts lt Sreactants and ?S is negative

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Entropy

• When a solid turns to a liquid and a liquid to a
  gas, entropy increases (?S ? 0)
• When a gas dissolves in a liquid, entropy
  decreases (?S ? 0)

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Sample Problem

• What is the entropy change for the single
  replacement reaction between sodium chloride and
  fluorine? SNaF 51.5 J/mol K

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Sample Problem

• What is the change in entropy for the
  decomposition of potassium chlorate? SKClO3(s)
  143.7 J/mol K

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Gibbs Free Energy

• Spontaneous rxns release energy that can perform
  work
• Free energy change represents the maximum work
  that a spontaneous rxn can perform
• Two ways to calculate Gibbs Free Energy
• ?G ?Gproducts - ?Greactants
• ?G ?H-T?S

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Gibbs Free Energy Cont.

• If ?G is negative the rxn is spontaneous
• If ?G is positive the rxn is not spontaneous and
  requires sustained input of energy to make it
  occur
• If ?G0 the rxn is at equilibrium
• Exothermic rxns in which the products are more
  disordered than its reactants, ?G is always
  negative-regardless of temp

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Gibbs Free Energy Cont.

• Endothermic rxns in which the products are more
  ordered than its reactants, ?G is always
  positive- regardless of temp
• ?G for a nonspontaneous rxn is the minimum amount
  of work that must be performed by an external
  source to make the rxn occur

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• Ex. Calculate the free energy change for the
  following rxn.
• Determine if the reaction will be spontaneous or
  not.
• C(s) 2H2(g) ? CH4(g)
• ?S-80.7 J/mol K ?H -75.0k J/mol T298K