THERMOCHEMISTRY or Thermodynamics

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THERMOCHEMISTRY- A
By Dr. Hisham E Abdellatef
http//www.staff.zu.edu.eg/ezzat_hisham/browseMyFi
les.asp?path./userdownloads/physical20chemistry
20for20clinical20pharmacy/
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THERMOCHEMISTRY
The study of heat released or required by
chemical reactions or heat changes caused by
chemical reactions
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What is Energy?
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Internal Energy (E)

• Internal Energy, E, The total energy of a system
• E(system) EK (system) EP (system)

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Potential Kinetic Energy

• Potential energy energy a motionless body has
  by virtue of its position.

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Potential Kinetic Energy
Kinetic energy
Energy of Motion EK ½ mv2

• i.e. the kinetic energy of an object depends on
• both its mass and its speed.

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Potential Kinetic Energy
Kinetic energy energy of motion.
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Problem

• What is the kinetic energy of a person whose
  mass is 130 lb (59.0 kg) traveling in a car at 60
  mph (26.8 m/s).

• The SI unit of energy, kg.m2/s2, is the Joule.

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UNITS OF ENERGY

• 1 calorie heat required to raise temp. of 1.00
  g of H2O by 1.0 oC.
• 1000 cal 1 kilocalorie 1 kcal
• But we use the unit called the JOULE
• 1 cal 4.184 joules

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FIRST LAW OF THERMODYNAMICS

• ?E q w

Energy is conserved!
?????????
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SYSTEM
?E q w
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• 1st Law of Thermodynamics
• Energy can neither created not destroyed, only
  transformed from one form to another

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System and Surroundings

• In chemical reactions, heat is often transferred
  from the system to its surroundings, or vice
  versa.

• The substance or mixture of substances under
  study in which a change occurs is called the
  thermodynamic system (or simply the system.)
• The surroundings are everything outside of the
  thermodynamic system.

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SURROUNDINGS
HEAT
HEAT
HEAT
HEAT
SYSTEM
SYSTEM
EXOTHERMIC
ENDOTHERMIC
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Exothermic
If ?E lt 0, Efinal lt Einitial
cellular respiration of glucose
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Endothermic
If ?E gt 0, Efinal gt Einitial
Photosynthesis is an endothermic reaction
(requires energy input from sun)
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Heat of Reaction

• An exothermic process is a chemical reaction or
  physical change in which heat is evolved (q is
  negative).
• An endothermic process is a chemical reaction or
  physical change in which heat is absorbed (q is
  positive).

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Exothermic
Endothermic
6.2
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Enthalpy Diagrams

• Values of DH are measured experimentally.
• Negative values indicate exothermic reactions.
• Positive values indicate endothermic reactions.

An increase in enthalpy during the reaction DH
is positive.
A decrease in enthalpy during the reaction DH is
negative.
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Exothermic Examples

• Oxidation wooden splint burning (giving off
  light, heat, CO2, H2O
• Burning H2 in air,
• body reactions,
• dissolving metals in acid,
• mixing acid and water,
• sugar dehydration

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Endothermic Examples

• Electrolysis (breaking water down into H2 and O2
  by running electricity in it)
• Photosynthesis, pasteurization, canning
  vegetables
• 2 H2 O2 ? 2H2O energy
• 4 g 32 g ? 36 g 136 600 cal
• 2H2O energy ? 2H2 O2
• 36 g 136 600 cal ? 4g 32 g

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Changes in Internal Energy

• If ?E gt 0, Efinal gt Einitial
• Therefore, the system absorbed energy from the
  surroundings.
• This energy change is called endergonic.

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Changes in Internal Energy

• If ?E lt 0, Efinal lt Einitial
• Therefore, the system released energy to the
  surroundings.
• This energy change is called exergonic.

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Enthalpy (H) of the reaction(Comes from Greek
for heat inside)

• the sum of internal energy and the product of
  this pressure and volume.
• H E PV
• E is the internal energy,
• P is the pressure and
• V is the volume of the system.
• It is also called heat content.
• ?H H product H reactants Hp Hr

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Calculation of ?H from ?E

• When the system changes at constant pressure, the
  change in enthalpy, ?H, is
• ?H ?(E PV)
• ?H ?E P?V
• At constant pressure and temperature
• The enthalpy of a chemical is measured in
  kilojoules per mole (kJmol-1).

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?H ?E P?V
For solid and liquid
?H ?E

• At constant volume

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• In case of gases
• ?H ?E P?V (I)
• ?V ?n x V
• ?n no of moles of products - no of moles of
  reactants
• Px?V PVx?n (II)
• But PV RT (for one mole of gas)
• Putting RT in place of PV in equation (II) we get
• P?V RT?n
• Substituting the value of P AV in equation (I) we
  get
• ?H ?E ?n RT

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• R 1.987 cal. (2 cal.)
• or
• 8.314 joules

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Example- 1

• The heat of combustion of ethylene at 17 C and
  at constant volume is -332.19 kcal. Calculate the
  heat of combustion at constant pressure
  considering water to be in liquid state (R 2
  cal.).

The chemical equation for the combustion of
ethylene is C2H4 3 O2 2CO2(g) 2H2O
(1) 1 mole 3 moles 2moles
negligible volume No. of moles of the
products 2 No. of moles of the reactants
4 ?n (2-4) -2
?H ?E ?n RT ?H -332.19 2 x
I0-3x -2 x 290 -333.3 kcal
Given that ?E-332. 19 kcal. T 27317
290k R2cal2xlO-3kcals.
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Example 2

• The heat of combustion of carbon monoxide at
  constant volume and at 17 C is -283.3 Kj.
  Calculate its heat of combustion at constant
  pressure(R 8.314 J degree-1 mole-1).

CO(g) ½ O2(g) ?CO2(g) 1 mole ½ mole 1
mole No. of mles of products 1 No. of moles of
reactants 1.5 n No. of moles of products - No.
of moles of reactants 1-1.5 -0.5 ?H ?E
?n x RT ?H -283.3 (-0.5x (8.314x10-3) x 290
- 283.3-1.20 -284.5 KJ Heat of combustion of CO
at constant pressure is -284.5 kJ.
Given that ?E -283.3 kJ T (27317) 290
K. R 8.314 J or 8.314x10-3 KJ
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Problem

• How much heat could be obtained by the combustion
  of 10.0 grams CH4burning in the presence of
  oxygen at constant pressure.

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Thermochemical equations

• factors which affects the quantity of heat
  evolved or absorbed during a physical or a
  chemical transformation.
• Amount of the reactants and products
• Physical state of the reactants and products
• Temperature
• Pressure

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Thermochemical equations

• It must essentially
• be balanced
• give the value of ?E or ?H corresponding to the
  quantities of substances given by the equation
• mention the physical states of the reactants and
  products . The physical states are represented by
  the symbols (s), (L), (g) and (aq) for solid,
  liquid, gas and aqueous states respectively.

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Example of thermochemical equation

• H2 ½ O2 ? H2O ?H -68.32 Kcal
• 1 mole of hydrogen reacts with 0.5 mole of
  oxygen, one mole of water is formed and 68.32
  Kcal of heats evolved at constant pressure.
• But, not specify whether water is in the form of
  steam or liquid
• H2 (g) ½ O2 (g)? H2O (L) ?H -68.32
  Kcal
• H2 (g) ½ O2 (g)? H2O (g) ?H -57.80
  Kcal.

Effect of temperature ?????
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Standard enthalpy change ?H.

• The heat change at
• 298 K and
• one atmosphere pressure
• is called the standard heat change or standard
  enthalpy change. It is denoted by ?H.

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How do we relate change in temp. to the energy
transferred? Heat capacity (J/oC) heat
supplied (J)
temperature (oC)
Heat Capacity heat required to raise temp. of
an object by 1oC
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Specific heat capacity is the quantity of energy
required to change the temperature of a 1g sample
of something by 1oC
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Variation of heat (or enthalpy) of reaction with
temperature Kirchoff's equations.
internal energies of the reactants and products.

• 1. At constant volume,
• ?EE2-E1
• Differentiating this equation with respect to
  temperature at constant volume, we get

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Kirchoff's equations.

• But we have already seen that

heat capacities
heat capacities of the products and reactants
Integrating between temperature T l and T2 ,
we have
?E2- ?E1 ? E2 -?E1 ? Cv
(T2-T1) (3)
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Kirchoff's equations.

• 2. At constant pressure

  ?HH2-H1
And finely . ?H2- ?H1 ?Cp (T2-T1) (6)
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Example 3 3

• The heat of reaction
• ½ H2 ½ Cl2 HCl at 27 C is -22.1
  Kcal.
• Calculate the heat of reaction at 77C. The
  molar heat capacities at constant pressure at 27
  C for hydrogen, chlorine and HCl are 6.82, 7.70
  and 6.80 cal mol-1, respectively.

Here ½ H2 ½Cl- ? HCl ? H -22.1 Kcal ? Cp
Heat capacities of products - Heat capacities of
reactants 6.80-½(6.82) ½ (7.70) 6.80 -
7.26 -0.46 x 10-3 Kcal ?H2- ?H1 ?Cp
(T2-T1) ?H2 - (-22.1) (-0.46 x 10 -3) x 50
-21.123
T2 273 77 350 K T1 273 27 300 K
T2-T1 (350-300)K50K
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HEAT CAPACITY

• The heat required to raise an objects T by 1 C.

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Specific Heat Capacity

• How much energy is transferred due to T
  difference?
• The heat (q) lost or gained is related to
• a) sample mass
• b) change in T and
• c) specific heat capacity

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Specific Heat Capacity

• Substance Spec. Heat (J/gK)
• H2O 4.184
• Ethylene glycol 2.39
• Al 0.897
• glass 0.84

Aluminum
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Specific Heat Capacity

• If 25.0 g of Al cool from 310 oC to 37 oC, how
  many joules of heat energy are lost by the Al?

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Specific Heat Capacity

• If 25.0 g of Al cool from 310 oC to 37 oC, how
  many joules of heat energy are lost by the Al?

where ?T Tfinal - Tinitial q (0.897
J/gK)(25.0 g)(37 - 310)K q - 6120 J
Notice that the negative sign on q signals heat
lost by or transferred OUT of Al.
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• We define the standard state of a substance as
  the state of the pure substance at 1 atm pressure
  and the temperature of interest (usually 25 C).
• The standard enthalpy change (?H) for a reaction
  is the enthalpy change in which reactants and
  products are in their standard states.
• The standard enthalpy of formation (?Hf) for a
  reaction is the enthalpy change that occurs when
  1 mol of a substance is formed from its component
  elements in their standard states.

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End of par A