John A. Schreifels

1
Chapter 6

• Thermochemistry

2
Overview

• Understanding Heats of Reaction
• Energy and its Units
• Heat of Reaction
• Enthalpy and Enthalpy Change
• Thermochemical Equations
• Stoichiometry and Heats of Reaction
• Measuring Heats of Reaction
• Uses of Heats of Reaction
• Hesss Law
• Standard Enthalpies of Formation
• Fuels Foods, etc.

3
Energy and Its Units

• Thermochemistry the study of the energy changes
  that take place during a reaction.
• Reactions generally proceed in whichever
  direction that will produce products with lower
  energy than the reactants.
• Heat and Energy
• Heat energy transferred from hotter to colder
  one.
• Kinetic energy the energy of movement of matter
• . Units Joule 1 kgm2/s2.
• E.g. what is the kinetic energy of 50.0 kg person
  running at a velocity of 20 m/s.
• Potential energy stored energy. E.g. water at
  the top of a mountain, a compressed spring, a
  chemical bond.

4
Energy Changes and Energy Conservation

• First law of Thermodynamics Energy is neither
  created nor destroyed but may be converted from
  one form to another.
• Energy forms
• Thermal energy a form of kinetic energy energy
  transfer results in a temperature change.
• Chemical energy a form of potential energy.
  Energy is stored in chemical bonds and released
  when a compound reacts.
• During reaction, energy is usually transformed
  from chemical to thermal energy.
• First law can be written as
• E q w
• where q heat involved in the process and w
  work done by or to the system.
• Work can be electrical or pressure volume

5
Internal Energy and The First Law of
Thermodynamics

• Internal Energy, E, is the sum of the potential
  and kinetic energy of a system.
• System - that part of the universe upon which we
  are focusing, e.g. reactions.
• Surroundings - eveything else in the universe
  which is not the system.
• State function - property depending only upon
  initial and final states and not upon path.
• Extent of transfer of energy is ?E Efinal ?
  Einitial.
• System is the reference point and a negative sign
  indicates that energy is flowing from the system
  to the surrounding.
• exothermic (exo out of). Heat flows from
  system to surroundings.
• endothermic (endo into). Heat flows from
  surroundings to system.
• C(gr) O2(g) ? CO2(g) 393.5 kJ Exothermic
• CO2(g) 393.5 kJ ? C(gr) O2(g) Endothermic

6
Sign conventions

• Sign of DE will depend upon the sign of q and w.

DE q w
7
Internal Energy and The First Law of
Thermodynamics2

• The conditions of measurement must be included
  when discussing the total internal energy since
  it is related to
• chemical identity of reactants and products
• their temperature, pressure, and physical state.
• Internal energy of a system is a state function.
  
• State Function a property of the system which
  depends only in the initial and final states and
  is independent of the history of the system.
• Several energy functions to be discussed have
  this property.

8
Expansion Work

• Work force acting over some distance w ? d x
  F (referenced to the system).
• During reactions often there is an expansion of
  gases against some pressure where pressure is
  equal to the force per unit area
• or .
• Work is obtained by substitution
• w ? d x F ? d x (PxA) or
• w ? P?V.
• The first law can be restated as E q ? P?V.
• This equation indicates that the amount of heat
  involved in a reaction will be reduced by the
  amount of work being done for a given change in
  the internal energy.
• E.g. Calculate the work done when during a
  reaction the gaseous products cause the volume to
  change from 22.4 L to 44.8 L against a constant
  pressure of 1.00 atm.

9
Expansion Work2

• If work is performed at constant temperature,
  then the amount of work performed will depend
  upon the change in the number of moles (?n)
• Modifications of the ideal gas law (PV nRT
  where n mol and R 8.3145 J/molK) lead to an
  alternative way of determining work. ?P?V ??nRT
  
• The presence of solids and liquids need not be
  considered since the molar volume of either a
  solid or liquid is about 1000x smaller than the
  molar volume of a gas.
• E.g. determine the work performed during the
  combustion of methane at 1.00 atm and 298.15 K.
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l)

10
Energy and Enthalpy

• From the first law q ?E P?V.
• With no change in volume the equation simplifies
  to qV ?E.
• At constant pressure qP ?E P?V.
• There are times when both volume and pressure can
  change the heat involved in the reaction is then
  a more complicated function of ?E.
• Enthalpy the heat output at constant pressure.
  H E PV.
• In general, ?H ?E P?V V?P.
• At constant pressure, a change in enthalpy is
  given by
• ?H ?E P?V qP.
• Normally, ?H and ?E are fairly close to each
  other in magnitude. In the combustion of propane
  (see book), ?E ?2043 kJ, ?H ?2041 kJ and w
  ?P?V ?2kJ.

11
Enthalpies of Physical and Chemical Change

• Enthalpies of Physical Change
• Heating a substance increases the temperature
  the amount of heat absorbed is proportional to
  the heat capacity of the species being heated.
• Amount of energy absorbed during phase change is
  proportional to the heat of phase change.

• Sum the heats in each portion of the curve to
  determine overall heat.
• The heat for converting a solid directly to a gas
  is called the heat of sublimation and is equal to
  the sum of the heats of fusion and vaporization
  at the same temperature.

12
Enthalpies of Chemical Change

• ?H is an extensive property its value depends
  upon the amount of reactants.
• ?H is attached to the chemical equation to
  indicate the amount of heat involved in the
  reaction.
• E.g. the combustion of methane
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l) ?H ? 890kJ
• 2CH4(g) 4O2(g)? 2CO2(g) 4H2O(l) ?H ? 1780kJ
• E.g.2 determine the amount of heat that would be
  evolved when 150 g of methane is burnt.
• Reversing reaction changes the sign of the heat.
• CO2(g) 2H2O(l) ? CH4(g) 2O2(g) ?H 890kJ.

13
Calorimetry and Heat Capacity

• Calorimeter a device that measures the change
  in the heat content or internal energy.
• Atmospheric pressure
• Bomb calorimeter
• Heat capacity the amount of heat absorbed by a
  substance to raise the temperature by a given
  amount.
• Heat transferal to a substance like a solid or a
  liquid, causes a change in temperature that is
  proportional to the amount of heat involved.
• H2O absorbs 4.18 J for every gram and C
• Al absorbs 0.902 J for every gram and C
• The amount of heat absorbed is directly
  proportional to amount of absorbing species
• where s specific heat capacity, C molar heat
  capacity and
• ?T Tfinal ? Tinitial.

14
Calorimetry and Heat Capacity2

• Energy change from any source such as reactions
  or phase change can be measured with heat
  capacity.
• E.g. How much heat is required to heat 500.0 g of
  water from 20.0C to 100.0C.
• The enthalpy change in the system is the negative
  of the heat of the calorimeter.
• E.g. exothermic reactions gives off heat to
  calorimeter. ?H ? qcalorimeter.
• E.g.2 When 2.00 g of ethanol was burned, all of
  the reaction energy was used to heat water in a
  calorimeter. Determine ?H for the reaction if
  the temperature of 200.0 g of water increased
  from 25.0C to 89.0C.
• Heat capacity of a whole calorimeter is used for
  complicated calorimeters such as the bomb
  calorimeter.
• E.g. 800.0 J of heat caused the temperature of a
  calorimeter was found to increase by 2.0 K. In
  some other reaction, the temperature of the
  calorimeter was found to increase by 5.0 K.
  Calculate the heat of the reaction.

15
Hesss Law

• Hesss law when a reaction at constant
  temperature and pressure can be written as the
  summation of a series of reactions, the enthalpy
  change, ?H, of the reaction is equal to the
  summation of the ?Hs of the individual
  reactions.
• E.g. determine the heat of formation of NO2(g)
• ½ N2(g) O2(g) ? NO2 ?
• Forming NO2(g) from N2(g) can be thought of as 2
  step process

16
Hesss Law2

• Missing steps in a sequence can be determined
  using Hesss law.
• E.g. determine the heat for methanol
  decomposition to its elements from the heat of
  combustion and the other given reactions. Heat
  of combustion is
• CH3OH(g) O2(g) ? CO2(g) H2O(l) ?H ?726.4 kJ.

DH DH1 DH2 DH3 or DH1 -726.4
393.51 571.66 238.77 kJ
17
Standard Heats of Formation

• Standard state the pure form of a substance at 1
  atm usually at 25C.
• Standard reaction enthalpies, ?H, difference in
  enthalpy between products and reactants of a
  reaction each in their standard states.
• Standard heat (enthalpy) of formation the
  standard reaction enthalpy per mol for the
  synthesis of a compound from its elements.
• Since reaction enthalpy depends upon conditions
  of experiment, it is usually reported at some
  reference condition, ?H
• Most tables present enthalpy data in its standard
  state and as the heat of formation.
• E.g. (HCl) is ?92.3 kJ and the reaction is
• ½H2(g) ½Cl2(g) ? HCl(g) ?92.3 kJ
• ?H of pure elements in their most stable form
  under standard conditions is defined as zero.
  E.g. Na(g), Na(s) C(g), C(gr), C(d).
• of elements in another form often given.
• Na(s) ? Na(g) ?H 107.8 kJ/mol. Also called
  the enthalpy of sublimation.

18
Calculations with Heat of Formation

• ?H of a reaction can be obtained from of all
  reactants and products.
• E.g. Determine the heat of combustion of ethanol,
  CH3CH2OH, from heats of formation in the book.
• Solution CH3CH2OH 3O2 ? 2CO2(g)
  3H2O(l) DHc ?
• For any general reaction such as aA bB ? cC
  dD,