Thermodynamics

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Thermodynamics
Heat
Work
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Heat is a form of energy

• Mechanical work done on a system produces a rise
  in temperature like heat added to the system.

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Joules Experiment
Joule showed that mechanical energy could
be converted into heat energy.
DT
M
F
Dx
H2O
W FDx
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• Thermodynamics Thermodynamics is concerned with
  interconversions of different forms of energy.
• It was developed as a mathematical tool for
  studying phenomena such as the way in which heat
  energy can be converted into mechanical energy in
  heat engines
• Thermodynamics provides a means for deciding
  whether a process will occur spontaneously
• Spontaneous in this context implies nothing
  about how fast it will take place - e.g.
  combustion of a diamond
• Thermodynamics is linked at a fundamental level
  to the nature of the universe
• The Industrial Revolution depended on heat
  engines, most of which (like these steam engines)
  were of very low efficiency. The development of
  physical theories, and mathematical tools, to
  analyse these systems led to rapid improvements
  in technology.

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• Energy can be divided into two categories -
  kinetic energy and potential energy
• Kinetic energy includes all forms of energy that
  result from movement - either linear motion or
  rotation.
• Heat, which is molecular motion.
• Radiant energy - the kinetic energy of photons of
  light and other electromagnetic radiation
• Mechanical energy
• Electrical energy as currents of moving electrons
  or charged particles
• Potential energy includes all forms of energy
  that are stored.
• Energy stored in chemical bonds
• Energy stored in concentration gradients
• Energy stored as electrical potential (separation
  of charges)
• Energy stored in the nuclei of atoms
• The basic unit of energy is the Joule (J)

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Thermodynamic Systems

• A thermodynamic system is a collection of matter
  which has distinct boundaries. OR
• A real or imaginary portion of universe whish has
  distinct boundaries is called system. OR
• A thermodynamic system is that part of universe
  which is under thermodynamic study.

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Systems
Thermodynamic System A quantity of matter or a
region in space chosen for study. Surroundings Ev
erything external to the system. Boundary Surface
that separates the system from the surrounding.
It may be fixed or movable
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Closed and Open Systems
Closed system (Control mass) A fixed amount of
mass chosen for study (no mass can cross its
boundary). Heat and work can cross the boundary,
volume may also change. e.g. piston
cylinder. Open system (Control volume) A
selected region chosen for study. Mass, heat and
work can cross its boundary, volume may also
change. e.g. water heater, car radiator, turbine,
nozzle. Isolated system A system closed to
mass, heat and work flows. It is not affected by
the surroundings.
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Open Systems
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Properties of a system
Thermodynamic Property A measurable quantity that
defines the condition of a system e.g.
temperature T pressure P mass
m volume V
density ?
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Extensive and Intensive properties

• Properties are of 2 types
• Intensive properties Independent of mass. e.g.
  P, T, v, ?.
• Extensive properties Change with mass. e.g. m, V,
  Energy

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Temperature and 0th law of thermodynamics
Temperature Degree of hotness of coldness 0th
law of thermodynamics When 2 bodies have equality
of temperature with a 3rd body, then they have
equality of temperature with each other.
TA
TB
TC
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Absolute scale of temperature
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Internal Energy (E)

• Definition The total of the kinetic and
  potential energy in a system.
• E Kinetic Energy Potential Energy

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The First Law of Thermodynamics

• Internal Energy
• Internal Energy total energy of a system.
• Involves translational, rotational, vibrational
  motions.
• Cannot measure absolute internal energy.
• Change in internal energy,

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First Law of Thermodynamics

• The change in the internal energy of a system E,
  is equal to the heat input Q minus the work done
  by the system.
• ?E Q-W
• The internal energy is the energy stored in the
  system.
• For an ideal gas the internal energy is the
  kinetic energy of the gas

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The First Law of Thermodynamics

• Relating DE to Heat(q) and Work(w)
• Energy cannot be created or destroyed.
• Energy of (system surroundings) is constant.
• Any energy transferred from a system must be
  transferred to the surroundings (and vice versa).
• From the first law of thermodynamics

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Work

• Work is a form of energy. Its the energy
  involved in moving something. If nothing moves,
  no work is done.
• Work in chemical terms is usually done with
  pressure and volume changes.

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The First Law of Thermodynamics
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QUIZ

• Calculate the energy change for a system
  undergoing a process in which 15.4 kJ of heat
  flows and where 6.3 kJ of work is done on the
  system.

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ANSWER

• ?E q w
• q - 15.40 J
• w 06.30 J
• ?E - 15.40 J 06.30 J - 09.10 J

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The First Law of Thermodynamics

• Exothermic and Endothermic Processes
• Endothermic absorbs heat from the surroundings.
• An endothermic reaction feels cold.
• Exothermic transfers heat to the surroundings.
• An exothermic reaction feels hot.

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Pressure-volume work

• P f / A
• Work done by the system (gas)
• w f ?x
• w P ?V

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• Example 2
• A gas expands by 0.50 L against a constant
  pressure of 0.50 atm at 25 C. What is the work
  in erg done by the system? (1.0 atm 1.013 x 106
  dyne/cm2)
• Solution
• W - P?V
• - (0.50 x 1.013 x 106 dyne/cm2) x 500 cm3
• - 2.50 x 103 dyne/cm
• - 2.50 x 103 erg

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Reversible work

• Isothermal work expansion against variable
  pressure.
• n number of moles, R gas constant (8.314
  JK-1mol-1, 1.987 cal K-1mol-1, 0.0802 L.atm.
  K-1mol-1)
• T absolute temperature

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• Example
• What is the maximum work done in the isothermal
  reversible expansion of 2 moles of an ideal gas
  from 1 to 5 litres at 25 C?
• Solution
• 7976.43 J

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Special Forms of the 1st law

• ?E q w
• Case 1 Isothermal process
• ?E 0, hence, q - w
• Case 2 Isochoic process
• w 0, ?E qv
• Case 3 Adiabatic process
• q 0, ?E w
• Case 4 Isobaric
• ?E q w

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Energy Changes in ChemicalReactions

• How are energy changes measured?
• One Answer Calorimetry, q Ccal?T
• What thermodynamic quantities do we get?
• Constant Volume qv ?E.
• Constant Pressure qP ?H
• In most calorimetry, ?T is very small, initial
• and final states are at nearly constant T

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Chemical Reactivity

• What drives chemical reactions? How do they
  occur?
• The first is answered by THERMODYNAMICS and the
  second by KINETICS.
• Have already seen a number of driving forces
  for reactions that are PRODUCT-FAVORED.
• formation of a precipitate
• gas formation
• H2O formation (acid-base reaction)
• electron transfer in a battery

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Chemical Reactivity

• But energy transfer also allows us to predict
  reactivity.
• In general, reactions that transfer energy to
  their surroundings are product-favored.

So, let us consider heat transfer in chemical
processes.
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Enthalpy of Reaction from Heat of Formation

• Standard Heat of Formation (?Hof)
• The amount of energy gained or lost when1 mole
  of the substance is formed from its elements
  under standard conditions(25C, 1 atm 101.3
  kPa)

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• Example 2
• Standard enthalpies of formation are C2H5OH(l)
  -228, CO2 (g) -394, and H2O(l) -286 kJ/mol.
  Calculate the enthalpy of the reaction,
• C2H5OH (l) 3O2 (g) 2 CO2 (g) 3 H2O (l)
• Solution
• ?HR 3x?Hf(H2O (l)) 2x?Hf(CO2(g)) 1x?Hf
  C2H5OH(l) 3x?HfO2(g)
• 3molX-286 kJ/mol 2molx-394
  kJ/mol-1molx-228 kJ/mol 3molx0.0 kJ/mol
• - 858 (-788) -228 - 1546 228
  - 1418 kJ

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Heat of Reactions from Bond Energies

• When bonds are formed, energy is released.
• In order to break bonds, energy must be absorbed
• Exothermic Products have stronger bonds than
  the reactants.
• Endothermic Products have weaker bonds than the
  reactants.

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• O2(g) 2O(g) ?H 490.4 kJ
• H2(g) 2H(g) ?H 431.2 kJ
• H2O(g) 2H(g) O(g) ?H 915.6 kJ
• 2H2(g) O2(g) 2H2O(g)    ?H ?
• ?H reaction bonds broken Energy (absorbed)
  - bonds formed Energy (released)
• 2 H-H OO 4 O-H
• 2 x 431.2 1 x 490.4 4 x 457.8
• 862.4 490.4 1831.2
• 1532.8 1831.2
• - 478.4 kJ

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Heat of Combustion

• IS the Heat Evolved when One mole of a Fuel is
  Burned in Enough Oxygen

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• EXample
• How much heat is evolved when 54.0 g glucose
  (C6H12O6) is burned according to this equation?
• ?Hcomb -2808 kJ/mol, C 12, O 16, H 1
• C6H12O6(s) 6O2(g) 6H2O(l)
  6CO2(g)
• Answer
• Mass of glucose 54.0 g C6H12O6
• ?Hcomb -2808 kJ
• Energy (q) ?
• Convert grams of C6H12O6 to moles of C6H12O6
• 54.0 g / 180 g mol-1 0.300 mol C6H12O6
• ? H 0.300 mol x - 2808 kJ/mol - 842 kJ

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Heat at constant VOLUME

• If we consider a system at constant volume. ?V is
  zero. If ?V is zero, then the work is zero.
• At constant volume the heat is equal to ?U.
• We can measure heat at constant volume in a Bomb
  Calorimeter.
• The calculations are the same for bomb
  calorimetry as for coffee-cup calorimetry.
• The heat calculated is a measure of the INTERNAL
  ENERGY change instead of Enthalpy.

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Finding ?H from ?E

• We can find ?E from bomb calorimetry. But that is
  at constant volume.
• How do we find ?H?
• Here ?n is the change in number of moles of GAS
  in the balanced chemical equation.
• ?E qp w ?H - P?V ?H - ?nRT

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• ?nRT
• This is an amount of energy that would be
  represented as work if work could be done.
• This can be zero if there is no change in number
  of moles of gas.

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Application

• For the reaction
• Mg (s) ½ O2 (g) MgO (s)
• the enthalpy change is -601.2 kJ. What is the
  internal energy change for this reaction? How
  many g of magnesium must react to effect an
  internal energy change of -22.4 kcal?

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Hesss Law

• The Enthalpy of a Reaction is the Same if it
  takes place in One or More than One step.

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Hesss Law Energy Level Diagrams
Forming H2O can occur in a single step or in a
two steps. ?Htotal is the same no matter which
path is followed.
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USING ENTHALPY

• Making H2O from H2 involves two steps.
• H2(g) 1/2 O2(g) ---gt H2O(g) 242 kJ
• H2O(g) ---gt H2O(liq) 44 kJ
• --------------------------------------------------
  ---------------
• H2(g) 1/2 O2(g) --gt H2O(liq) 286 kJ
• Example of HESSS LAW
• If a rxn. is the sum of 2 or more others, the net
  ?H is the sum of the ?Hs of the other rxns.

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QUIZs

• The standard heat of formation, ?Hof, for sulfur
  dioxide (SO2) is -297 kJ/mol. How many kJ of
  energy are given off when 25.0 g of SO2 (g) is
  produced from its elements? 
• The heat of reaction for the combustion of 1 mol
  of ethyl alcohol is -9.50 102 kJC2H5OH (l)
  3 O2 (g) ? 2 CO2 (g) 3 H2O (l) 9.5 102 kJ.
• How much heat is produced when 11.5 g of
  alcohol is burned?
• ?H for the complete combustion of 1 mol of
  propane is
• -2.22 103 kJ
• C3H8 (g) 5 O2 (g) ? 3 CO2 (g) 4 H2O (l)
• Calculate the heat of reaction for the
  combustion of 33.0 g of propane.

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