Energy

1
Energy Chemistry

• 2H2(g) O2(g) ? 2H2O(g) heat and light
• This can be set up to provide ELECTRIC ENERGY in
  a fuel cell.
• Oxidation
• 2 H2 ? 4 H 4 e-
• Reduction
• 4 e- O2 2 H2O ? 4 OH-

H2/O2 Fuel Cell Energy, page 288
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Energy Chemistry

• ENERGY is the capacity to do work or transfer
  heat.
• HEAT is the form of energy that flows between 2
  objects because of their difference in
  temperature.
• Other forms of energy
• light
• electrical
• kinetic and potential

• Positive and negative particles (ions) attract
  one another.
• Two atoms can bond
• As the particles attract they have a lower
  potential energy

NaCl composed of Na and Cl- ions.
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Potential Kinetic Energy
Kinetic energy energy of motion.
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Internal Energy (E)

• PE KE Internal energy (E or U)
• Internal Energy of a chemical system depends on
• number of particles
• type of particles
• temperature

• The higher the T the higher the internal energy
• So, use changes in T (?T) to monitor changes in E
  (?E).

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Thermodynamics

• Thermodynamics is the science of heat (energy)
  transfer.

Heat transfers until thermal equilibrium is
established. ?T measures energy transferred.

• SYSTEM
• The object under study
• SURROUNDINGS
• Everything outside the system

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Directionality of Heat Transfer

• Heat always transfer from hotter object to cooler
  one.
• EXOthermic heat transfers from SYSTEM to
  SURROUNDINGS.

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Directionality of Heat Transfer

• Heat always transfers from hotter object to
  cooler one.
• ENDOthermic heat transfers from SURROUNDINGS to
  the SYSTEM.

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Energy Chemistry

• All of thermodynamics depends on the law of
• CONSERVATION OF ENERGY.
• The total energy is unchanged in a chemical
  reaction.
• If PE of products is less than reactants, the
  difference must be released as KE.

Energy Change in Chemical Processes
Potential Energy of system dropped. Kinetic
energy increased. Therefore, you often feel a
Temperature increase.
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HEAT CAPACITY

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

Which has the larger heat capacity?
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Specific Heat Capacity

• How much energy is transferred due to Temperature
  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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Table of specific heat capacities
Substance Phase cp J g-1 K-1 Cp J mol-1 K-1
Air (typical room conditionsA) gas 1.012 29.19
Aluminium solid 0.897 24.2
Argon gas 0.5203 20.7862
Copper solid 0.385 24.47
Diamond solid 0.5091 6.115
Ethanol liquid 2.44 112
Gold solid 0.1291 25.42
Graphite solid 0.710 8.53
Helium gas 5.1932 20.7862
Hydrogen gas 14.30 28.82
Iron solid 0.450 25.1
Lithium solid 3.58 24.8
Mercury liquid 0.1395 27.98
Nitrogen gas 1.040 29.12
Neon gas 1.0301 20.7862
Oxygen gas 0.918 29.38
Uranium solid 0.116 27.7
Water gas (100 C) 2.080 37.47
Water liquid (25 C) 4.1813 75.327
Water solid (0 C) 2.114 38.09
All measurements are at 25 C unless noted. Notable minimums and maximums are shown in maroon text. All measurements are at 25 C unless noted. Notable minimums and maximums are shown in maroon text. All measurements are at 25 C unless noted. Notable minimums and maximums are shown in maroon text. All measurements are at 25 C unless noted. Notable minimums and maximums are shown in maroon text.
Aluminum
A Assuming an altitude of 194 meters above mean
sea level (the worldwide median altitude of
human habitation), an indoor temperature of 23
C, a dewpoint of 9 C (40.85 relative
humidity), and 760 mmHg sea levelcorrected
barometric pressure (molar water vapor content
1.16).
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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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Heat/Energy TransferNo Change in State

• q transferred (sp. ht.)(mass)(?T)

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Heat Transfer

• Use heat transfer as a way to find specific heat
  capacity, Cp
• 55.0 g Fe at 99.8 C
• Drop into 225 g water at 21.0 C
• Water and metal come to 23.1 C
• What is the specific heat capacity of the metal?

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Heating/Cooling Curve for Water
Note that T is constant as ice melts or water
boils
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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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FIRST LAW OF THERMODYNAMICS

• ?E q w

Energy is conserved!
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The First Law of Thermodynamics

• Exothermic reactions generate specific amounts of
  heat.
• This is because the potential energies of the
  products are lower than the potential energies of
  the reactants.

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The First Law of Thermodynamics

• There are two basic ideas of importance for
  thermodynamic systems.
• Chemical systems tend toward a state of minimum
  potential energy.
• Chemical systems tend toward a state of maximum
  disorder.
• The first law is also known as the Law of
  Conservation of Energy.
• Energy is neither created nor destroyed in
  chemical reactions and physical changes.

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SYSTEM
?E q w
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ENTHALPY

• Most chemical reactions occur at constant P, so

Heat transferred at constant P qp qp ?H
where H enthalpy
and so ?E ?H w (and w is usually small) ?H
heat transferred at constant P ?E ?H change
in heat content of the system ?H Hfinal -
Hinitial
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ENTHALPY

• ?H Hfinal - Hinitial

If Hfinal gt Hinitial then ?H is positive Process
is ENDOTHERMIC
If Hfinal lt Hinitial then ?H is negative Process
is EXOTHERMIC
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USING ENTHALPY

• Consider the formation of water
• H2(g) 1/2 O2(g) ? H2O(g) 241.8 kJ

Exothermic reaction heat is a product and ?H
241.8 kJ
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USING ENTHALPY

• Making liquid H2O from H2 O2 involves two
  exothermic steps.

H2 O2 gas
H2O vapor
Liquid H2O
Making H2O from H2 involves two steps. H2(g)
1/2 O2(g) ? H2O(g) 242 kJ
H2O(g) ? H2O(l) 44 kJ H2(g) 1/2 O2(g) ?
H2O(l) 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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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.
Active Figure 6.18
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Hesss Law

• Hesss Law of Heat Summation, ?Hrxn ?H1 ?H2
  ?H3 ..., states that the enthalpy change for a
  reaction is the same whether it occurs by one
  step or by any (hypothetical) series of steps.
• Hesss Law is true because ?H is a state
  function.
• If we know the following ?Hos

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Hesss Law

• For example, we can calculate the ?Ho for
  reaction 1 by properly adding (or subtracting)
  the ?Hos for reactions 2 and 3.
• Notice that reaction 1 has FeO and O2 as
  reactants and Fe2O3 as a product.
• Arrange reactions 2 and 3 so that they also
  have FeO and O2 as reactants and Fe2O3 as a
  product.
• Each reaction can be doubled, tripled, or
  multiplied by half, etc.
• The ?Ho values are also doubled, tripled, etc.
• If a reaction is reversed the sign of the ?Ho is
  changed.

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Hesss Law

• Hesss Law in a more useful form.
• For any chemical reaction at standard conditions,
  the standard enthalpy change is the sum of the
  standard molar enthalpies of formation of the
  products (each multiplied by its coefficient in
  the balanced chemical equation) minus the
  corresponding sum for the reactants.

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Hesss Law
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Hesss Law

• Given the following equations and ?Ho?values
• calculate ?Ho for the reaction below.

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Hesss Law

• Use a little algebra and Hesss Law to get the
  appropriate ?Ho?values

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Some Thermodynamic Terms

• Notice that the energy change in moving from the
  top to the bottom is independent of pathway but
  the work required may not be!
• Some examples of state functions are
• T (temperature), P (pressure), V (volume), ?E
  (change in energy), ?H (change in enthalpy the
  transfer of heat), and S (entropy)
• Examples of non-state functions are
• n (moles), q (heat), w (work)

???H along one path ???H along another path

• This equation is valid because ?H is a STATE
  FUNCTION
• These depend only on the state of the system and
  not how it got there.
• V, T, P, energy and your bank account!
• Unlike V, T, and P, one cannot measure absolute
  H. Can only measure ?H.

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Some Thermodynamic Terms

• The properties of a system that depend only on
  the state of the system are called state
  functions.
• State functions are always written using capital
  letters.
• The value of a state function is independent of
  pathway.
• An analog to a state function is the energy
  required to climb a mountain taking two different
  paths.
• E1 energy at the bottom of the mountain
• E1 mgh1
• E2 energy at the top of the mountain
• E2 mgh2
• ?E E2-E1 mgh2 mgh1 mg(?h)

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Standard States and Standard Enthalpy Changes

• Thermochemical standard state conditions
• The thermochemical standard T 298.15 K.
• The thermochemical standard P 1.0000 atm.
• Be careful not to confuse these values with STP.
• Thermochemical standard states of matter
• For pure substances in their liquid or solid
  phase the standard state is the pure liquid or
  solid.
• For gases the standard state is the gas at 1.00
  atm of pressure.
• For gaseous mixtures the partial pressure must be
  1.00 atm.
• For aqueous solutions the standard state is 1.00
  M concentration.

• ?Hfo standard molar enthalpy of formation
• the enthalpy change when 1 mol of compound is
  formed from elements under standard conditions.
• See Table 6.2 and Appendix L

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Enthalpy Values
Depend on how the reaction is written and on
phases of reactants and products

• H2(g) 1/2 O2(g) ? H2O(g) ?H -242 kJ
• 2H2(g) O2(g) ? 2H2O(g) ?H -484 kJ
• H2O(g) ? H2(g) 1/2 O2(g) ?H
  242 kJ
• H2(g) 1/2 O2(g) ? H2O(l) ?H -286 kJ

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?Hfo, standard molar enthalpy of formation

• H2(g) ½ O2(g) ? H2O(g) ?Hf (H2O, g) -241.8
  kJ/mol
• C(s) ½ O2(g) ? CO(g) ?Hf of CO -
  111 kJ/mol
• By definition, ?Hfo 0 for elements in their
  standard states.

Use ?Hs to calculate enthalpy change for
H2O(g) C(graphite) ? H2(g) CO(g)
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Using Standard Enthalpy Values

• In general, when ALL enthalpies of formation are
  known,

?Horxn ? ?Hfo (products) - ? ?Hfo (reactants)
Remember that ? always final initial
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Using Standard Enthalpy Values

• Calculate the heat of combustion of methanol,
  i.e., ?Horxn for
• CH3OH(g) 3/2 O2(g) ? CO2(g) 2 H2O(g)
• ?Horxn ? ?Hfo (prod) - ? ?Hfo (react)