Energy

1
Chapter 5

• Energy
• Thermodynamics
• (rev. 0910)

2
Definition

• Thermodynamics-
• is the study of energy transformations.

3
Chemical Reactions

• Chemical reactions involve not just the
  conversion of reactants into products, but also
  involve an energy change in the form of heatheat
  released as the result of a reaction, or heat
  absorbed as a reaction proceeds.
• Energy changes accompany all chemical reactions
  and are due to rearranging of chemical bonding.

4
Making Bonds

• Addition of energy is always a requirement for
  the breaking of bonds but the breaking of bonds
  in and of itself does not release energy.
• Energy release occurs when new bonds are formed.

5
Bond Energy

• If more energy is released when new bonds form
  than was required to break existing bonds, then
  the difference will result in an overall release
  of energy.
• If, on the other hand, more energy is required to
  break existing bonds than is released when new
  bonds form, the
• difference will result overall in energy being
  absorbed.

6
Overall Reaction Energy

• Whether or not an overall reaction releases or
  requires energy depends upon the final balance
  between the breaking and forming of chemical
  bonds.

7
Energy is...

• the ability to do work or produce heat.
• conserved.
• made of heat and work.
• a state function. (Energy is a property that is
  determined by specifying the condition or state
  (e.g., temperature, pressure, etc.) of a system
  or substance.)
• independent of the path, or how you get from
  point A to B.

8
Energy

• While the total internal energy of a system (E)
  cannot be determined, changes in internal energy
  (E) can be determined.
• The change in internal energy will be the amount
  of energy exchanged between a system and its
  surroundings during a physical or chemical
  change.
• ? E E final - E initial

9
Definitions

• Work is a force acting over a distance.
• Heat is energy transferred between objects
  because of temperature difference. (Heat is not a
  property of a system or substance and is not a
  state function. Heat is a processthe transfer of
  energy from a warm to a cold object.)

10
System vs Surroundings

• The Universe is divided into two halves.
• system and the surroundings.
• In a chemistry setting, a system includes all
  substances undergoing a physical or chemical
  change.
• The surroundings would include everything else
  that is not part of the system.

11
Heat

• Most commonly, energy is exchanged between a
  system and its surroundings in the form of heat.
• Heat will be transferred between objects at
  different temperatures.
• Thermochemistry is the study of thermal energy
  changes.

12
Exo vs Endo

• Exothermic reactions release energy to the
  surroundings.
• Endothermic reactions absorb energy from the
  surroundings.

13
Heat
Potential energy
14
Heat
Potential energy
15
Three Parts

• Every energy measurement has three parts
• A unit ( Joules of calories).
• A number.
• a sign to tell direction.
• negative - exothermic
• positive- endothermic

16
Surroundings
System
Energy
DE lt0
17
Surroundings
System
Energy
DE gt0
18
Same rules for heat and work

• Heat given off is negative.
• Heat absorbed is positive.
• Work done by the system on the surroundings is
  negative.
• Work done on the system by the surroundings is
  positive.

19
First Law of Thermodynamics

• The energy of the universe is constant.
• It is also called the
• Law of conservation of energy.
• q heat w work
• In a chemical system, the energy exchanged
  between a system and its surroundings can be
• accounted for by heat (q) and work (w).
• DE q w
• Take the systems point of view to decide signs.

20
Conservation of Energy

• Energy exchanged between a system and its
  surroundings can be considered to off set one
  another.
• The same amount of energy leaving a system will
  enter the surroundings (or vice versa), so the
  total amount of energy remains constant.

21
Metric Units

• The SI (Metric System) unit for all forms of
  energy is the joule (J).

22
Heat and Work

• DE q w
• - q is exothermic -q -?H
• q is endothermic
• -w is done by the system
• w is done on the system
• Note
• ?H stands for enthalpy which is the heat of
  reaction

23
Practice Problem

• A gas absorbs 28.5 J of heat and then performs
  15.2 J of work. The change in internal
• energy of the gas is
• (a) 13.3 J
• (b) - 13.3 J
• (c) 43.7 J
• (d) - 43.7 J
• (e) none of the above

24
Answer

• (b) E q w
• 28.5 J - 15.2 J 13.3 J

25
Practice Problem

• Which of the following statements correctly
  describes the signs of q and w for the following
  exothermic process at 1 atmosphere pressure and
  370 Kelvin?
• H2O(g) ? H2O(l)
• (a) q and w are both negative
• (b) q is positive and w is negative
• (c) q is negative and w is positive
• (d) q and w are both positive
• (e) q and w are both zero

26
Answer

• (c). An exothermic indicates q is negative and
  the gas is condensing to a liquid so it is
  exerting less pressure on its surroundings
  indicating w is positive.

27
What is work?

• Work is a force acting over a distance.
• w F x Dd
• P F/ area
• d V/area
• w (P x area) x D (V/area) PDV
• Work can be calculated by multiplying pressure by
  the change in volume at constant pressure.
• Use units of literatm or Latm

28
Pressure and Volume Work

• Work refers to a force that moves an object over
  a distance.
• Only pressure/volume work (i.e., the
  expansion/contraction of a gas) is of
  significance in chemical systems and only when
  there is an increase or decrease in the amount of
  gas present.

29
Work needs a sign

• If the volume of a gas increases, the system has
  done work on the surroundings.
• work is negative
• w - PDV
• Expanding work is negative.
• Contracting, surroundings do work on the system W
  is positive.
• 1 Latm 101.3 J

30
Example

• When, in a chemical reaction, there are more
  moles of product gas compared to
• reactant gas, the system can be thought of as
  performing work on its surroundings (making w lt
  0) because it is pushing back, or moving back
  the atmosphere to make room for the expanding
  gas.
• When the reverse is true, w gt 0.

31
Compressing and Expanding Gases

• Compressing gas
• Work on the system is positive
• Work is going into the system
• Expanding gas
• Work on the surroundings is negative
• Work is leaving the system

32
Clarification Info

• If the reaction is performed in a rigid
  container, there may be a change in pressure,
• but if there is no change in volume, the
  atmosphere outside the container didnt move
  and without movement, no work is done by or on
  the system.
• If there is no change in volume (V 0), then no
  work is done by or on the system (w 0) and the
  change in internal energy will be entirely be due
  to the heat involved ( ?E q).

33
Examples

• What amount of work is done when 15 L of gas is
  expanded to 25 L at 2.4 atm pressure?
• If 2.36 J of heat are absorbed by the gas above.
  what is the change in energy?
• How much heat would it take to change the gas
  without changing the internal energy of the gas?

34
Enthalpy

• The symbol for Enthalpy is H
• H E PV (thats the definition)
• at constant pressure.
• DH DE PDV
• the heat at constant pressure qp can be
  calculated from
• DE qp w qp - PDV
• qp DE P DV DH

35
DH DE PDV

• Using DH DE PDV
• the heat at constant pressure qp can be
  calculated from
• DE qp w (if w - PDV then)
• DE qp PDV (now rearrange)
• qp DE P DV DH

36
Examples of Enthapy Changes

• KOH(s) ? K(aq) OH-1 (aq)
  ?Hsolution - 57.8 kJ mol1
• C3H8(g) 5O2(g) ? 3CO2(g) 4H2O(l)
  ?Hcombustion -2221kJ/mol
• H2O(s) ? H2O(l)
  ?Hfusion 6.0 kJ/mol
• Fe2O3(s) 2Al(s) ? Al2O3(s) 2Fe(s)
  ?Hreaction - 852 kJ/mol
• Ca(s) O2(g) H2(g) ? Ca(OH)2(s)
  ?Hformation - 986 kJ/mol1

37
3 Methods

• There are a variety of methods for calculating
  overall enthalpy changes that you should be
  familiar with.
• The three most common are the
• the use of Heats of Formation
• Hesss Law
• the use of Bond Energies

38
Heat of Reaction

• To compare heats of reaction for different
  reactions, it is necessary to know the
  temperatures at which heats of reaction are
  measured and the physical states of the reactants
  and products.
• Look in the Appendix of the textbook to find
  Standard Enthapy tables.

39
Standard Enthalpy of Formation

• Measurements have been made and tables
  constructed of Standard Enthalpies of Formation
  with reactants in their standard states.
• Use the symbol DHºf
• Standard state is the most stable physical state
  of reactants at
• 1 atmosphere pressure
• specified temperatureusually 25 C
• 1 M solutions
• For solids which exist in more than one
  allotropic form, a specific allotrope must be
  specified.

40
?Hformation

• It is important to recognize that the
  ?Hformation (abbreviated as ?Hf) is really just
  the heat of reaction for a chemical change
  involving the formation of a compound from its
  elements in
• their standard states.

41
Standard Enthalpies of Formation

• The standard heat of formation is the amount of
  heat needed to form 1 mole of a compound from its
  elements in their standard states.
• See the table in the Appendix
• Remember For an element the value is 0

42
Equation Practice

• You need to be able to write the equation
  correctly before solving the problem.
• Try
• What is the equation for the formation of NO2 ?
  Try writing the equation.

43
Practice Answer

• ½N2 (g) O2 (g) NO2 (g)
• You must make one mole to meet the definition.

44
Since we can manipulate the equations

• We can use heats of formation to figure out the
  heat of reaction.
• Lets do it with this equation.
• C2H5OH 3O2(g) 2CO2 3H2O
• which leads us to this rule.

45
Hesss Law

• Definition
• When a reaction may be expressed as the algebraic
  sum of other reactions, the enthalpy change of
  the reaction is the algebraic sum of the enthalpy
  changes for the combined reactions.

46
Hess Law

• Enthalpy is a state function.
• It is independent of the path.
• We can add equations to come up with the desired
  final product, and add the DH values.
• Two rules to remember
• If the reaction is reversed the sign of DH is
  changed
• If the reaction is multiplied, so is DH

47
Enthalpy

• As enthalpy is an extensive property, the
  magnitude of an enthalpy change for a chemical
  reaction depends upon the quantity of material
  that reacts.
• This means
• if the amount of reacting material in an
  exothermic reaction is doubled, twice the
  quantity of heat energy will be released.

48
For the oxidation of sulfur dioxide gas

• SO2(g) ½O2(g) ? SO3(g)
• ?H - 99 kJ/mol
• Doubling the reaction results in
• 2SO2(g) O2(g) ? 2SO3(g)
• ?H - 198 kJ/mol
• Notice that if you double the reaction, you must
  double the ?H value.

49
Sign Change of ?H

• 2SO2(g) O2(g) ? 2SO3(g)
• ?H - 198 kJ/mol
• If the reaction is written as an endothermic
  reaction
• 2SO3(g) ? 2SO2(g) O2(g)
• ?H 198 kJ/mol

50
Tips for Hesss Law Problems

• It is always a good idea to begin by looking for
  species that appear as reactants and products in
  the overall reaction.
• This will provide a clue as to whether a reaction
  needs to be reversed or not.
• Second, consider the coefficients of species that
  appear in the overall reaction.
• This will help determine whether a reaction needs
  to be multiplied before the overall summation.

51
Example

• C(s) O2(g) ? CO2(g) ?Hf - 394
  kJ/mol
• 2H2(g) O2(g) ? 2H2O(l)
  ?Hf - 572 kJ/mol
• CO2(g) 2H2O(l) ? CH4(g) 2O2(g) ?Hf 891
  kJ/mol
• -------------------------------------------------
  ----
• C(s) 2H2(g) ? CH4(g) ?Hf - 75
  kJ/mol

52
Hesss Law Example
Given
DHº 77.9kJ
DHº 495 kJ
DHº 435.9kJ
Calculate DHº for this reaction
53
Example

• Given
  calculate DHº for this reaction

DHº -1300. kJ
DHº -394 kJ
DHº -286 kJ
54
Problems to Try

• Try 12-3 Practice Problems

55
O2
NO2
-112 kJ
180 kJ
H (kJ)
NO2
68 kJ
N2
2O2
56
Calorimetry

• Calorimetry is the study of the heat released or
  absorbed during physical and chemical reactions.
• For a certain object, the amount of heat energy
  lost or gained is proportional to
• the temperature change. The initial temperature
  and the final temperature in the calorimeter
• are measured and the temperature difference is
  used to calculate the heat of reaction.

57
Equipment Calorimeter

• There are two kinds of calorimeters
• constant pressure
• bomb

58
Calorimetry

• Constant pressure calorimeter
• (called a coffee cup calorimeter)
• A coffee cup calorimeter measures DH.
• The calorimeter can be an insulated cup, full of
  water.

59
Definitions

• Heat capacity is the amount of energy required to
  raise the temperature of an object 1 kelvin or 1
  C.
• Specific heat capacity is the heat capacity of
• 1 gram of a substance.
• Molar heat capacity is the heat capacity of
• 1 mole of a substance.

60
Heat Capacity

• Heat capacity is an extensive property, meaning
  it depends on the amount presenta large amount
  of a substance would require more heat to raise
  the temperature 1 K than a small amount of the
  same substance.

61
Specific Heat Capacity

• Definition
• The specific heat capacity of each substance is
  an intensive property which relates the heat
  capacity to the mass of the substance.

62
Specific Heat Capacity c

• q mass x c x DT
• Or written as
• q mcDT
• This is the main equation for calorimetry
  calculations
• mass will be in grams
• Units for c are J/g K or J/g C
• The specific heat of water is 1 cal/g ºC

63
Molar Heat Capacity

• Molar Heat Capacity is the heat capacity of 1
  mole of a substance.
• molar heat capacity c/moles
• heat molar heat x moles x DT
• Remember that heat is shown as ?H
• Make the units work and youve done the problem
  right.

64
Sign of q

• If a process results in the sample losing heat
  energy, the loss in heat is designated as q is
  negative.
• The temperature of the surroundings will
  increase during this exothermic process.
• If the sample gains heat during the process,
  then q is positive. The temperature of the
• surroundings will decrease during an endothermic
  process.
• The amount of heat that an object gains or loses
  is directly proportional to the change in
• temperature.

65
Examples

• The specific heat of graphite is 0.71 J/gºC.
• Calculate the energy needed to raise the
  temperature of 75 kg of graphite from 294 K to
  348 K.

66
Extra Problem

• A 46.2 g sample of copper is heated to 95.4ºC and
  then placed in a calorimeter containing 75.0 g of
  water at 19.6ºC. The final temperature of both
  the water and the copper is 21.8ºC.
• What is the specific heat of copper?

67
Calorimetry

• Constant volume calorimeter is called a bomb
  calorimeter.
• Material is put in a container with pure oxygen.
  Wires are used to start the combustion. The
  container is put into a container of water.
• The heat capacity of the calorimeter is known and
  tested.
• Since DV 0, PDV 0, DE q

68
Bomb Calorimeter

• thermometer
• stirrer
• full of water
• ignition wire
• Steel bomb
• sample

69
Properties

• intensive properties are not related to the
  amount of substance.
• Examples density, specific heat, temperature.
• Extensive property - does depend on the amount of
  substance.
• Examples Heat capacity, mass, heat from a
  reaction.

70

• Collegeboard. (2007-2008). Professional
  Development workshop materials Special focus
  thermochemistry. http//apcentral.collegeboard.com
  /apc/public/repository/5886-3_Chemistry_pp.ii-88.p
  df