Chap 6 Thermochemistry

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Chap 6 Thermochemistry

• F.Y. Hsu

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6.1 Energy

• Literally means work within, however no object
  contains work
• Energy is the capacity to do work, the capacity
  to move something.
• 2 basic types of energy
• Potential (possibility of doing work because
  of composition or position)
• Kinetic (moving objects doing work)

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Energy

• Potential Energy in a gravitational field
• ( position)

• Kinetic Energy energy of motion

PE mgh m mass (kg) g gravity constant (m
s2) h height (m)
KE 1/2mv2 m mass (kg) v velocity (m s1) v2
(m2 s2)
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Work
Work is the product of the force in the direction
of motion and the distance the object is moved
Work force distance ? energy (J)
Collisions in the real world are not perfectly
elastic
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6.2 Thermochemistry Basic Terms

• Thermochemistry is the study of energy changes
  that occur during chemical reactions.
• System the part of the universe being studied.
• Surroundings the rest of the universe.

Universe
Surroundings
Surroundings
System
Surroundings
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Types of Systems

• Open energy and matter can be exchanged with the
  surroundings.
• Closed energy can be exchanged with the
  surroundings, matter cannot.
• Isolated neither energy nor matter can be
  exchanged with the surroundings.

OPEN
CLOSED
ISOLATED
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Internal Energy (U)

• Internal Energy (U) is the total energy contained
  within the system, partly as kinetic energy and
  partly as potential energy
• Kinetic involves three types of molecular motion

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Internal Energy (U)

• Internal Energy (U) is the total energy contained
  within the system, partly as kinetic energy and
  partly as potential energy
• Potential energy involves intramolecular
  interactions and intermolecular interactions

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Heat

• Heat is energy transfer resulting from thermal
  differences between the system and surroundings
• Heat passes spontaneously from the region of
  higher temperature to the region of lower
  temperature
• Thermal equilibrium occurs when the system and
  surroundings reach the same temperature and heat
  transfer stops.

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Figure 6-6 Heat Transfer Illustrated
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Work

• Like heat, work is an energy transfer between a
  system and its surroundings.
• Unlike heat, work is caused by a force moving
  through a distance (heat is caused by a
  temperature difference).
• A negative quantity of work signifies that the
  system loses energy.
• A positive quantity of work signifies that the
  system gains energy.
• There is no such thing as negative energy nor
  positive energy the sign of work (or heat)
  signifies the direction of energy flow.

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Calculating Work for gas

• we will consider only pressure-volume work.
• work (w) P?V
• when a gas expands, ?V is positive and the work
  is negative system loses energy

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6.3 Internal energy (U), States function and the
first law of thermodynamics

• The state of a system refers to its exact
  condition, determined by the kinds and amounts of
  matter present, the structure of this matter at
  the molecular level, and the prevailing pressure
  and temperature

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State function

• A state function is a property that has a unique
  value that depends only the present state of a
  system, and does not depend on how the state was
  reached (does not depend on the history of the
  system).

?U Uf Ui
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First Law of Thermodynamics

• The Law of Conservation of Energy states that in
  a physical or chemical change, energy can be
  exchanged between a system and its surroundings,
  but no energy can be created or destroyed
• The change in U is related to the energy
  exchanges that occur as heat (q) and work (w)

The First Law ?U q w
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First Law Sign Conventions

• Energy entering a system carries a positive sign
  if heat is absorbed by the system, q gt 0. If work
  is done on a system, w gt 0
• Energy leaving a system carries a negative sign
  if heat is given off by the system, q lt 0. If
  work is done by a system, w lt 0

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Example 6.1

• A gas does 135 J of work while expanding, and at
  the same time it absorbs 156 J of heat. What is
  the change in internal energy?

Sol
heat is absorbed by the system this means q is a
positive quantity ? q 156 J. Work is done by
the system this means w is a negative
quantity ?w 135 J. ??U q w 156 J
(135 J) 21 J
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6-4 Heats of Reaction

• The heat of reaction (qrxn) is the quantity of
  heat exchanged between the system and its
  surroundings
• for exothermic reactions,
• (a) in isolated systems, system T ?
• (b) in non-isolated systems, heat is given
  off to the surroundings, i.e., q lt 0
• For endothermic reactions,
• (a) in isolated systems, system T ?
• (b) in non-isolated systems, heat is
  absorbed from the surroundings, i.e., q gt 0

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Figure 6.11 Conceptualizing an Exothermic
Reaction
In an isolated system, all heat is absorbed by
the solution. Maximum temperature rise.
Surroundings are at 25 C
Typical situation some heat is released to the
surroundings, some heat is absorbed by the
solution.
Hypothetical situation all heat is instantly
released to the surroundings. Heat qrxn
In an exothermic reaction, chemical energy in a
system is converted to thermal energy
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Internal Energy Change at Constant Volume

• For a system where the reaction is carried out at
  constant volume, DV 0
• w P?V and ?U q w DU qV.
• All the thermal energy produced by conversion
  from chemical energy is released as heat no P-V
  work is done.

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Internal Energy Change at Constant Pressure

• For a system where the reaction is carried out at
  constant pressure,
• DU qP PDV or
• DU PDV qP (w PDV )
• Most of the thermal energy is released as heat.
• Some work is done to expand the system against
  the surroundings (push back the atmosphere).

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Enthalpy (?)

• Enthalpy (H) is the sum of the internal energy
  and the pressurevolume product of a system
• qP ?H ?U P?V
• Enthalpy is an extensive property (depends on how
  much of the substance is present)
• Enthalpy is a state function. U, P, and V are all
  state functions, therefore H must be a state
  function also

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Figure 6.14 Enthalpy Diagrams

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

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Reversing a Reaction

• DH changes sign when a process is reversed.
• Therefore, a cyclic process has the value DH 0.

Same magnitude different signs.
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?H in Stoichiometric Calculations

• For problem-solving, heat evolved (exothermic
  reaction) can be thought of as a product. Heat
  absorbed (endothermic reaction) can be thought of
  as a reactant.
• We can generate conversion factors involving DH.
• For example, the reaction

H2(g) Cl2(g) ? 2 HCl(g) ?H 184.6 kJ
184.6 kJ 1 mol Cl2
184.6 kJ 2 mol HCl
184.6 kJ 1 mol H2
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Example 6.5

• What is the enthalpy change associated with the
  formation of 5.67 mol HCl(g) in this reaction?
• Ex H2(g) Cl2(g) ?? 2 HCl(g) ?H 184.6 kJ

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6.5 Calorimetry

• Calorimetry is a technique used to measure heat
  exchange in chemical reactions
• A calorimeter is the device used to make heat
  measurements
• In a calorimeter we measure the temperature
  change of water or a solution to determine the
  heat absorbed or evolved by a reaction.

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Heat capacity Specific Heat

• The heat capacity (C) of a system is the quantity
  of heat required to change the temperature of the
  system by 1?
• C q/DT (units are J/C)
• Molar heat capacity is the heat capacity of one
  mole of a substance.
• The specific heat (s) is the heat capacity of one
  gram of a pure substance (or homogeneous
  mixture).
• s C/m q/(mDT)
• q s m DT

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Specific Heat

• q mass x specific heat x DT
• If DT is positive (temperature increases), q is
  positive and heat is gained by the system.
• If DT is negative (temperature decreases), q is
  negative and heat is lost by the system.
• The calorie, while not an SI unit, is still used
  to some extent.
• Water has a specific heat of 1 cal/(g oC).
• 4.184 J 1 cal
• One food calorie (Cal or kcal) is actually equal
  to 1000 cal.

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Table 6.1 Specific Heats of substances at 25?
Many metals have low specific heats.
The specific heat of water is higher than that of
almost any other substance.
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Measuring Enthalpy Changes for Chemical Reactions

• For a reaction carried out in a calorimeter, the
  heat evolved by a reaction is absorbed by the
  calorimeter and its contents.
• qrxn qcalorimeter
• qcalorimeter mass x specific heat x
  ?T
• By measuring the temperature change that occurs
  in a calorimeter, and using the specific heat and
  mass of the contents, the heat evolved (or
  absorbed) by a reaction can be determined and the
  enthalpy change calculated

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Example 6.11

• Ex A 50.0-mL sample of 0.250 M HCl at 19.50 C
  is added to 50.0 mL of 0.250 M NaOH, also at
  19.50 C, in a calorimeter. After mixing, the
  solution temperature rises to 21.21 C. Calculate
  the heat of this reaction.
• Sol
• (1)The solution volumes are additive. The volume
  of NaCl(aq) that forms is equal to 50.0 mL 50.0
  mL 100.0 mL.
• (2)The NaCl(aq) is sufficiently dilute that its
  density and specific heat are about the same as
  those values for pure water 1.00 g/mL and 4.18 J
  g1 C1, respectively.

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Bomb Calorimetry Reactions at Constant Volume

• Some reactions, such as combustion, cannot be
  carried out in a coffee-cup calorimeter.
• In a bomb calorimeter, a sample of known mass is
  placed in a heavy-walled bomb, which is then
  pressurized with oxygen.
• Since the reaction is carried out at constant
  volume
• qrxn qcalorimeter ?U
• (qP ?H ?U P?V)

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Example 6.13

• In a preliminary experiment, the heat capacity of
  a bomb calorimeter assembly is found to be 5.15
  kJ/C. In a second experiment, a 0.480-g sample
  of graphite (carbon) is placed in the bomb with
  an excess of oxygen. The water, bomb, and other
  contents of the calorimeter are in thermal
  equilibrium at 25.00 C. The graphite is ignited
  and burned, and the water temperature rises to
  28.05 C. Calculate ?H for the reaction
• C(graphite) O2(g) ?? CO2(g) ?H ?

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• Solution

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6.6 Hesss Law of Constant Heat Summation

• The heat of a reaction is constant, regardless of
  the number of steps in the process
• ?Hoverall S?Hs of individual reactions
• When it is necessary to reverse a chemical
  equation, change the sign of ?H for that reaction
• When multiplying equation coefficients, multiply
  values of ? H for that reaction

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Figure 6.19 Hesss Law An Enthalpy Diagram
We can find DH(a) by subtracting DH(b) from DH(c)
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Example 6.14

• Ex Calculate the enthalpy change for reaction
  (a) given the data in equations (b), (c), and
  (d).
• (a) 2 C(graphite) 2 H2(g) ? C2H4(g) DH
  ?
• (b) C(graphite) O2(g) ? CO2(g)
  DH 393.5 kJ
• (c) C2H4(g) 3 O2 ? 2 CO2(g) 2 H2O(l)
  DH 1410.9 kJ
• (d) H2(g) ½ O2 ? H2O(l) DH
  285.8 kJ

Sol
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6.7 Standard state

• The standard state of a solid or liquid substance
  is the pure element or compound at 1 atm pressure
  and the temperature of interest
• Gaseous standard state is the ideal gas at 1
  atm pressure and the temperature of interest
  (usually 25 C).

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6.7 Standard Enthalpies of Formation

• 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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Table 6.2 Standard Enthalpiesof Formation at 25
oC
Note The standard enthalpy of formation of a
pure element in its reference form is 0 Ex O2
(g) ., C (s), H2(g) at 1atm 25?
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Calculations Based onStandard Enthalpies of
Formation

• DHrxn Snp x DHf(products) Snr x
  DHf(reactants)
• The symbol S signifies the summation of several
  terms.
• The symbol n signifies the stoichiometric
  coefficient used in front of a chemical symbol or
  formula.
• In other words
• Add all of the values for DHf of the products.
• Add all of the values for DHf of the reactants.
• Subtract 2 from 1

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Example 6.15

• Ex Synthesis gas is a mixture of carbon monoxide
  and hydrogen that is used to synthesize a variety
  of organic compounds. One reaction for producing
  synthesis gas is
• 3 CH4(g) 2 H2O(l) CO2(g) ? 4 CO(g) 8
  H2(g) ?H ?
• Use standard enthalpies of formation from Table
  6.2 to calculate the standard enthalpy change for
  this reaction.

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• Solution

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Ionic Reactions in Solution

• We can apply thermochemical concepts to reactions
  in ionic solution by arbitrarily assigning an
  enthalpy of formation of zero to H(aq).

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Example 6.18

• Ex H(aq) OH(aq) ? H2O(l) ?H 55.8 kJ
• Use the net ionic equation just given, together
  with ?Hf 0 for H(aq), to obtain ?Hf for
  OH(aq).

Solution
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Spontaneous reaction

• A reaction that occurs (by itself) when the
  reactants are brought together under the
  appropriate conditions is said to be spontaneous.
• A spontaneous reaction isnt necessarily fast
  (rusting diamond ? graphite etc. are slow).

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Combustion Fuels

• Fossil Fuels Coal, Natural Gas, and Petroleum
• A fuel is a substance that burns with the release
  of heat.
• These fossil fuels were formed over a period of
  millions of years from organic matter that became
  buried and compressed under mud and water.
• Foods Fuels for the Body
• The three principal classes of foods are
  carbohydrates, fats, and proteins.
• 1 Food Calorie (Cal) is equal to 1000 cal (or 1
  kcal).

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Figure 6.21 Glyceryl Trilaurate A Typical Fat
Fats are esters (R-COO-R)
This particular fat is saturated all the CC
bonds are single bonds.