Topic 5: Energetics

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Topic 5 Energetics

• 5.1 Exothermic and endothermic reactions
• 5.2 Calculation of enthalpy changes

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Thermochemistry (Energetics)

• The study of energy involved during chemical
  reactions
• Heat
• the energy of motion of molecules
• All matter has moving particles at stp
• Temperature
• transfer of heat to a substance because of faster
  molecular movement (as long as there is no phase
  change)

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• A temperature change is explained as a change in
  kinetic energy (movement)
• Temperature depends on the quantity of heat (q)
  flowing out or in of the substance.
• Energy flowing in the system Endothermic
• Has a positive value
• Energy entering (feels cool)
• Energy flowing out of the system Exothermic
• Has a negative value
• Energy exiting (feels warm)

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Heat (q)

• qmc ?t
• qheat
• mmass
• ?tchange in temperature (tf-ti)
• cspecific heat capacity (J (g oC)-1)
• Specific heat capacity is the quantity of heat
  required to raise the temperature of a unit mass
  of a substance by one degree Celsius.

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Law of conservation of energy

• ?E universe O
• The total energy of the universe is constant, it
  is not created or destroyed, however it can be
  transferred from one substance to another.
• ?E universe ?E system ?E surroundings

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First Law of thermodynamics

• Any change in energy of a system is equivalent by
  an opposite change in energy of the surroundings.
• ?E system - ?E surroundings
• According to this law, any energy released or
  absorbed by a system will have a transfer of
  heat, q.
• So, q system - q surroundings

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Sample Problem

• 15 g of ice was added to 60.0 g of water. The Ti
  of water was 26.5 oC, the final temperature of
  the mixture was 9.7 oC. How much heat was lost
  by the water?
• qmc ?t
• q(60.0 g) (4.18 J/g oC) (9.7-26.5 oC)
• q - 4213.44 J -4.2 kJ

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Watch this flash video about heat flow

• http//www.mhhe.com/physsci/chemistry/animations/c
  hang_7e_esp/enm1s3_4.swf

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Enthalpy (?H)

• Total kinetic and potential energy of a system
  under constant pressure.
• The internal energy of a reactant or product
  cannot be measured, but their change in enthalpy
  (heat of reaction) can.
• ? H Hproducts Hreactants
• A change in enthalpy occurs during phase changes,
  chemical reactions and nuclear reactions.
• ? H system q surroundings

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Endothermic Reactions Method 1 enthalpy level
diagram
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Endothermic Reactions

• Method 2 enthalpy term outside of the equation
• 2 HgO (s) ? 2 Hg (l) O2(g) ?H181.67 kJ
• Method 3 enthalpy term within the equation
• 2 HgO (s) 181.67 kJ ? 2 Hg (l) O2(g)

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Exothermic ReactionsMethod 1 Enthalpy level
diagram
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Exothermic Reactions

• Method 2
• 4Al(s) 3O2(g)? 2 Al2O3(g) ?H-1675.7 kJ
• Method 3
• 4 Al(s) 3O2(g)? 2 Al2O3(g) 1675.7 kJ
• Neutralization and combustion reactions

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Calorimeters (qwater -qsystem)

• Used to measure the amount of energy involved in
  a chemical reaction.
• To be treated like
• Isolate/closed system
• Specific mass of water used.
• Energy flows to or from the water in the cups
• Measure the temperature change related to the
  water.

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Problem

1. An 25.6 g of an unknown metal with an initial
   temperature of 300 oC, is placed in 150.0 g of
   water with an initial temperature of 35.0 oC. If
   the waters temperature stabilizes at 55.0 oC,
   calculate the specific heat capacity of this
   metal.

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Bomb Calorimeter
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Heat Capacity

• Related to bomb calorimeters
• Unit is (J/oC) because its always with a set
  mass, so it is redundant to repeat the term over.

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15.1 Standard enthalpy changes of reaction

• Higher level
• 15.1.1 Define and apply the terms standard
  state, standard enthalpy change of formation, and
  stand enthalpy change of combustion
• 15.1.2 Determine the enthalpy change of a
  reaction using standard enthalpy changes of
  formation and combustion.

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Standard Molar Enthalpy of Formation

• Standard implies the states of the particle at 1
  atm and at 0oC.
• Quantity of energy released (-) or absorbed ()
  when one mole of a compound is formed directly
  from its elements at standard temperature and
  pressure.
• We use a table to find them.
• Unit for ?Hof kJ/mol
• Watch your states!

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Practice

• What is the standard molar enthalpy of formation
  for the following reaction?
• 2 Na (s) Cl2(g) ?2 NaCl (s) 814 kJ

By definition, the standard molar enthalpy of
formation is for ONE mole of product
formed. ?Hfo -407 kJ/mol for NaCl
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Standard Molar Enthalpy of Combustion (?Hcombo

• Energy changes involved with combustion reactions
  of one mole of a substance.
• Remember that these reactions are only measured
  once cooled to 25oC
• Combustion is a reaction with oxygen as a
  reactant (burning)
• Will need a table of values to use.

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Combustion reaction with alkanes

• Always form water and carbon dioxide.
• Ex CH4 2O2 ? CO2 2H2O
• Remember your alkanes CnH2n2
• Meth C1, Eth C2, Prop C3, But C4, Pent
  C5, Hex C6, Hept C7, Oct C8, Non C9 and
  Dec C10.

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Standard heats of reactions (?Hrxno)

• Measured from all products and reactants at their
  standard states. (if a solution, concentration
  1M)
• All elements at standard state ?Hof 0
• Most compounds have a negative ?Hof
• Use balanced equation, where n number of moles

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Calculating enthalpy changes

• Amount of a substance reacting matters, so can
  use q n?H.
• Remember namount of moles.
• If you are given a mass (g) and molar mass
  (g/mol), then you can solve for n by dividing
  mass by molar mass. (review from Topic 1
  stoichiometry section)

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Practice

• Calculate the ?Horxn for
• 4NH3 (g) 5O2 (g) ? 4NO (g) 6H2O (g)
• CO (g) H2O (g) ? CO2 (g) H2 (g)
• Calculate the ?Hocomb for
• 2CH3OH(l) 3O2(g) ? 2 CO2(g) 4H2O(l)
• 2C2H6 (g) 7 O2 (g) ? 4CO2 (g) 6 H2O (l)

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Practice

• 1/8 S8 (s) H2 (g) ? H2S (g) ?Hrxno -20.2
  kJ
• Is this an endo or exothermic reaction?
• What is the ?Hrxno for the reverse reaction?
• What is the ?H when 2.6 mol of S8 reacts?
• What is the ?H when 25.0 g of S8 reacts?

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Dont Forget to

• Read your text book
• Look over the questions assigned
• Use your course companion
• Use your study guide