Chemistry 100 Chapter 5

1
Chemistry 100 Chapter 5

• Energy Relationships in Chemistry

2
Thermochemistry

• Thermodynamics the study of energy and its
  transformations.
• Thermochemical changes energy changes
  associated with chemical reactions.
• System that specific part of the universe of
  interest to us.
• Surroundings the part of the universe not
  contained in the system.

3
3 types of Systems

• open system exchanges mass and energy
• closed system exchanges energy but no mass
• isolated system no exchange of either mass or
  energy

4
Three Types of Systems
5
Different Types of Energy

• Energy the ability to do work.
• Thermal energy associated with the random
  motions of atoms and molecules
• Heat energy transfer of thermal energy between
  two objects at different temperature.

6
Energy (contd)

• Chemical energy energy stored within the
  structural units of chemical substance.
• Potential energy the ability of an object to
  do work because of its position in a field of
  force.

7

• Kinetic Energy the work that can be performed
  by a moving object.
• The unit of energy
• 1 Joule (J)
• 1 kg m2/s2
• An older unit of energy
• 1 calorie (cal)
• 4.184 J exactly

8
The Law of Conservation of Energy

• The law of conservation of energy
• Energy is neither created nor destroyed in
  ordinary chemical and physical processes
• Converted from one type into another.

9

• This is also stated in terms of the first law of
  thermodynamics.

?E internal energy change of the system Ef and
Ei ? the energy of the final and initial states,
respectively
10
First Law of Thermodynamics

• Chemical reactions either absorb or release
  energy.
• Two terms
• Exothermic reaction heat is released to the
  surroundings.
• Endothermic reaction heat is supplied to the
  system by the surroundings.

11
Exothermic
12
Endothermic
13
The First Law Restated

• chemical systems examine the conversion of
  heat energy into work.

14
Signs for Heat and Work

• Work done by system on surroundings
• w -
• Work done by surroundings on system
• w
• q lt 0, heat flows to surroundings
• Exothermic -
• q gt 0, heat flows to system
• Endothermic

15
Pressure-volume Work

• Pressure volume work
• w -Pop ?V -Pop (Vf -Vi)
• This is the type of work done by the pistons in
  our automobile engines!
• The greater the magnitude of Pop, the gas has to
  "work harder" to obtain the same volume change.

16
Pressure-Volume Work
17
State and Path Functions

• ?E, ?H, ?V are examples of state functions.
• State functions numerical value doesnt depend
  on how the process is carried out.
• Work (w) and q (heat) are path functions
• The amount of work done or heat released depends
  on how the system changes states.

18

• Examine a chemical reaction.
• C (s) O2 (g) ? CO2 (g)
• ?E ECO2 (g) EC(s) EO2(g)
• This reaction has a negative enthalpy change (?H
  -393.5 kJ).

19

• From the first law
• ?surrE ?sysE 0
• ?surrE -?sysE
• The energy "lost" from the system is "gained" in
  the surroundings.

20
Enthalpies of Formation Standard Reaction
Enthalpies

• The enthalpy change for the reaction
• ?rH ?H(products) - ?H(reactants)
• We cannot measure the absolute values of the
  enthalpies!!
• How do we measure enthalpies (or heat contents)
  of chemical species?

21
The Formation Reaction

• A "chemical thermodynamic reference point."
• For CO and CO2
• C (s) O2 (g) ? CO2 (g)
• C (s) ½ O2 (g) ? CO (g)
• The "formation" of CO and CO2 from its
  constituent elements in their standard states
  under standard conditions.

22
The Formation Reaction

• The formation reaction
• For the formation of 1.00 mole of Na2SO3(s)
• 2 Na(s) S(s) 3/2 O2 (g) ? Na2SO3 (s)

The formation enthalpy of Na2SO3(s),
symbolised ?fH?Na2SO3 (s)
23
Standard Conditions for Thermodynamic Reactions

• The degree sign, either ? or ?, indicates
  standard conditions
• P 1.00 atm
• aqueous species 1.00 mol/L
• T temperature of interest (note 25?C or 298 K
  is used in the tables in your text).

24
The Significance of the Formation Enthalpy

• ?fH is a measurable quantity!
• Compare CO (g) with CO2 (g)
• C (s) 1/2 O2 (g) ? CO (g)
• ?fH CO(g) -110.5 kJ/mole
• C (s) O2 (g) ? CO2 (g)
• ?fH CO2(g) - 393.5 kJ/mole
• The formation enthalpy for CO2(g) is larger than
  the formation enthalpy of CO (g).

25
Reactions Enthalpies

• Formation enthalpies thermodynamic reference
  point,
• Formation of the elements from themselves is a
  null reaction ?fH? (elements) 0 kJ / mole.

26
The Combustion of Propane
27
The General Equation

• Calculate enthalpy changes from the formation
  enthalpies as follows.

Reverse a reaction, the sign of the enthalpy
change for the reaction is reversed. Multiply a
reaction by an integer, the enthalpy change is
multiplied by the same integer.
28
The Measurement of Energy Changes Calorimetry

• Calorimetry the measurement of heat and energy
  changes in chemical and physical processes.
• Heat capacity (C) the amount of heat (energy)
  needed to raise the temperature of a given mass
  of substance by 1C.
• Specific heat capacity (s) the amount of heat
  energy (in Joules, J) required to raise 1 g of a
  substance by 1C (units J/g C).

29

• General expression for heat capacity
• C m s
• m is the mass of the substance (in grams).
• Molar heat capacity
• Cm M s
• M molar mass of the substance
• s its specific heat capacity.

30
The Calorimeter

• A calorimeter a device which contains water
  and/or another substance with a known capacity
  for absorbing energy (heat).
• Calorimeters are adiabatic systems.
• All energy changes take place within the
  calorimeter.

31
Adiabatic System

• Adiabatic system thermally insulated from the
  rest of the universe
• No heat exchange between system and
  surroundings!
• For an adiabatic system,
• qtot qrxn qH2O qcal 0
• ?-qrxn qH2O qcal

32
The Constant Volume (Bomb) Calorimeter
?E qv
33
The Constant Pressure Calorimeter
?H qp
34
Relating the Enthalpy to the Internal Energy

• The enthalpy and the internal energy both
  represent quantities of heat.
• ?E qv.
• ?H qp.
• ?E and ?H are related as follows
• ?H ?E Pop ?V
• ?V the volume change for the reaction.

35

• For reactions involving gases
• ?V ?ng /(RT Pop)
• ?ng ? np (g) - ? nr (g)
• For most reactions, ?ng is small.
• The difference between the internal energy change
  and the enthalpy change is small.

36
Other important Enthalpy changes

• Many other important processes have associated
  enthalpy changes.
• The measurement of the heat changes for these
  process can give us some insight into the changes
  in intermolecular forces that occur during the
  transformation.

37
Heat of dilution and solution.

• ?solH the heat absorbed or given off when a
  quantity of solute is dissolved in a solvent.
• ?solH H(soln) - H(component)
• H(component) H (solid) H(solvent)

38

• For the process,
• HCl (aq, 6 M) ? HCl (aq, 1 M).
• A significant amount of heat is released when the
  acid solution is diluted.
• This is the enthalpy of dilution of the acid.
• ?dilH H(soln 2) H(soln ,1)

39
Lattice Enthalpies

• Look at the following process.
• NaCl (s) ? Na (g) Cl- (g)
• ?H ?latH 788 kJ/mole ? the lattice enthalpy
• A very endothermic reaction!
• Due to the strength of the ionic bond!

40
Latent Heats

• Latent heats are the enthalpy changes associated
  with phase transitions.
• H2O (l) ? H2O (g)
• ?rH ?vapH ? the enthalpy of vapourization.
• H2O (s) ? H2O (l)
• ?rH ?fusH ? the enthalpy of fusion.
• H2O (s) ? H2O (g)
• ?rH ?subH ? the enthalpy of sublimation.

41
Foods and Fuels

• Most of the chemical reactions that produce heat
  are combustion reactions.
• Note all combustion reactions are exothermic.
  
• Fuel values are generally reported as positive
  quantities.
• Obtaining fuel values calorimetry.

42
Calories, Food Calories, and Kilojoules

• When we read our cereal boxes, we may see the
  following
• 1 bowl cereal 30 g cereal 132 Cal (490 kJ).
• Isnt 1 calorie 4.184 J (not 4.184 kJ)?
• The fuel values of foods are reported as food
  calories (Cal).
• 1.00 food calorie (Cal) 1000 thermal calories
  (cal) 4184 J 4.184 kJ.

43
Combustion of Carbohydrates and Fats

• Most of the energy our body needs comes form the
  combustion of sugar and fats.
• For the glucose (blood sugar) combustion
• C6H12O6 (s) 6 O2 (g) ? 6 CO2 (g) 6 H2O (l)
• ?rH? -2816 kJ
• This energy is supplied quickly to the body!
• Average fuel value of carbohydrates 17 kJ/g.

44
Fats

• The combustion (metabolism) of fats also produces
  CO2 and H2O.
• The combustion of tristearin
• C57H110O6 (s) 163/2 O2 (g) ? 57 CO2 (g) 55
  H2O (l).
• ?rH? -37.8 x 104 kJ

45
Fuel Value of Fats

• Fats are the bodys energy stockpiles!
• Insoluble in water.
• Average fuel value 38 kJ/g about twice that
  of the carbohydrates.

46
Caloric Contents

• For proteins average fuel value 17 kJ/g,
  about the same value as for the carbohydrates.
• The relative amounts of proteins, fats, and
  carbohydrates in foods determines the caloric
  content.

47
Fossil Fuels

• Coal, petroleum, and natural gas are known as
  fossil fuels. They are collectively the major
  source of energy for commercial and personal
  consumption.
• Fossil fuels are mixtures of many different kinds
  of organic compounds.
• The fuel values of fossil fuels is directly
  related to the amount of carbon and hydrogen in
  the fuel.

48
Hydrogen As a Fuel

• Hydrogen has a huge fuel value (142 kJ/g).
• The combustion product is innocuous water.
• Obviously, there are problems!
• Two major difficulties with H2 as a fuel source.
• Where do we get the hydrogen?
• How do we store the hydrogen?