Thermochemistry

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Thermochemistry

• The study of energy and its transformations

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Definitions - Energy

• Chemical systems contain both kinetic energy and
  potential energy.
• Energy is the capacity to do work or to produce
  heat.
• An example of both is the combustion of
  gasoline. The gaseous products expand and do
  work (moving the pistons of an engine) and the
  reaction also produces heat.

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Definitions - Energy

• Kinetic energy is the energy of motion, and it
  depends upon the mass of the object and its
  velocity. Since molecules, especially those of
  gases, are in motion, they possess kinetic
  energy.

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Definitions- Energy

• Potential energy is energy due to position or
  composition.
• Chemical energy is potential energy due to
  composition. For example, gasoline and oxygen
  have the potential to produce energy if they
  react.

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Definitions

• When examining chemical systems or reactions, we
  consider the system and its surroundings.
• The system is where we put our focus. Typically,
  it is the reactants and products.
• The surroundings include everything else in the
  universe.

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Definitions

• If a reaction results in the evolution of heat,
  energy flows out of the system and into the
  surroundings. These reactions are exothermic.
• The energy lost by the system must be equal to
  the energy gained by the surroundings.

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Internal Energy

• The internal energy (E) of a system is the sum
  of the kinetic and potential energies of all of
  the particles of the system.
• It is generally not possible to determine the
  internal energy of a system, but we can measure
  changes in internal energy.
• Internal energy is changed by the flow of work
  and/or heat.

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Internal Energy the 1st Law

• The First Law of Thermodynamics states that
• The total internal energy of an isolated system
  is constant.
• However, it is not possible to completely
  isolate a system from its surroundings.

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Internal Energy

• Since energy may flow to or from the
  surroundings and the system, we are concerned
  with energy changes rather than the absolute
  value of the internal energy.
• ?E Efinal - Einitial

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Internal Energy

• Internal Energy is a state function. That is,
  it depends solely on the present state of the
  system, and not how it may have gotten to a
  particular state. A state function is
  independent of pathway.

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Internal Energy

• Internal energy (E) is a state function, and
  depends only on the state of the system, and not
  how it got to that state.

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Chemical Reactions and Energy

• Bond breaking always requires energy.
• Bond making always releases energy.
• In exothermic reactions, more energy is
  released in forming the products than is used in
  breaking apart the reactants.

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Exothermic Reactions

• The heat that is released comes from the
  potential energy stored in the bonds of the
  reactants.

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Definitions

• If a reaction involves the absorption of heat
  into the system it is an endothermic reaction.
• More energy is required to break the bonds in
  the reactants than is released in forming the
  products.

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Endothermic Reactions
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Chemical Reactions Work

• In addition to the release or taking in of
  energy (heat) when bonds are broken and formed,
  chemical reactions can do expansion work.
• Expansion work is done when the volume of
  reactants differs significantly from the volume
  of reactants.

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Work

• Expansion work results when a reaction produces
  more gaseous products than reactants, and thus
  pushed back the atmosphere as the reaction
  proceeds.
• If the volume of the system contracts during
  reaction (more gaseous reactants than products),
  work is done by the surroundings on the system

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Expansion Work

• Usually we just consider the volumes of gases
  in a chemical reaction.
• 2 H2O(g) ? 2 H2(g) O2(g)
• Since 2 moles of gaseous reactants produce 3
  moles of gaseous products, the system expands,
  and does work in pushing back the atmosphere.

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Expansion Work
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Expansion Work

• work Force x Distance
• Pressure Force/Area, or
• Force Pressure(Area)
• work Pressure(Area) x Distance
• work Pressure(Area) x ?h
• work Pressure (length x width) x ?h
• work P ?V

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Expansion Work

• work P ?V
• If gases are produced by a reaction and the
  volume expands, the system is doing work on the
  surroundings. The sign, when considering the
  system, must be negative. So,
• work - P ?V

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Energy, Work and Heat

• As the system loses energy to the surroundings,
  it can do so by losing heat (q), and/or doing
  work (w). As a result,
• ?E heat work
• ?E q w
• ?E q -P?V

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Energy, Work and Heat

• ?E q P?V
• Solving for q, the heat change,
• q ?E P?V

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Energy and Enthalpy

• Chemical reactions are sometimes carried out in
  a sealed vessel with a fixed volume, a bomb
  calorimeter. In this apparatus, the volume of
  the reaction mixture (system) cannot change. As
  a result, at constant volume,
• qv ?E

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Energy and Enthalpy

• Many chemical reactions are carried out in an
  open vessel. In this case, the reaction is
  performed at constant pressure, and the volume of
  the system is free to change during the course of
  the reaction. The heat transferred at constant
  pressure, qp, is defined as
• qp ?E P?V

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Energy and Enthalpy

• qp ?E P?V
• Since an open vessel is such a common
  apparatus, the heat transferred at constant
  pressure is given its own name, the enthalpy
  change, ?H.
• qp ?E P?V ?H

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Enthalpy

• Enthalpy is a state function, and is
  independent of reaction pathway.
• ?H Hfinal-Hinitial
• ?H Hproducts-Hreactants

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Sign Conventions

• Our focus will always be on the system. If
  energy flows out of the system into the
  surroundings, it will have a negative (-) sign.
• If energy flows from the surroundings into the
  system, it will have a positive () sign.

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Energy, Heat Work
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Standard Conditions

• Many reactions are categorized by their
  standard enthalpy change, ?Ho. The degree sign
  indicates standard conditions.
• Standard conditions specify that the reactants
  and products are in the same molar amounts
  represented by the coefficients in the balanced
  chemical reaction.

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Standard Conditions

• In a given experiment, the quantities of
  reactants and enthalpy change will vary, but the
  standard enthalpy change is reported based on
  molar quantities.
• The enthalpy of a reaction will also vary with
  the physical states of reactants or products as
  well as temperature and pressure.

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Standard Conditions

• A thermodynamic standard state refers to a
  specific set of conditions. The standard is used
  so that values of enthalpy changes can be
  directly compared.
• The standard state is the most stable form of a
  substance at 1 atm pressure and 25oC.

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Standard Conditions

• Standard conditions are indicated using a degree
  symbol ( o ). Standard conditions for
  thermochemical data differ from the standard
  conditions used in the gas laws.
• 1. All gases have a pressure of exactly 1 atm.
• 2. Pure substances are in the form that they
  normally exist in at 25oC and 1 atm pressure.
• 3. All solutions have a concentration of
  exactly 1M.

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Standard Conditions

• For example, since oxygen is a diatomic gas at
  25oC, the standard state of oxygen is O2(g) at a
  pressure of 1 atm.

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Calorimetry

• Calorimetry is the science of measuring heat.
  It typically involves measuring temperature
  changes as a substance loses or gains heat.
• Since substances vary in how much their
  temperature changes as heat is lost or gained, it
  is important to know the heat capacity (C) of
  substances involved in the reaction.

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Heat Capacity (C)

• The heat capacity of a substance is the amount
  of heat absorbed, usually in joules, per 1 degree
  (C or K) increase in temperature. The amount
  (mass) of the substance also determines the
  amount of heat lost or gained.
• C heat absorbed _q_
• increase in temp. ?T

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Heat Capacity

• The specific heat capacity is for a gram of a
  substance. It has the units J/oC-g or J/K-g.
• The molar heat capacity is for a mole of a
  given substance. It has the units J/oC-mol or
• J/K-mol.

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Calorimeter Constant

• In measuring heat changes during a reaction,
  any heat absorbed or lost be the calorimeter (the
  apparatus itself) must be considered. If this
  amount of heat is significant, the calorimeter
  constant may be provided or measured. This is
  the heat capacity of the specific apparatus used,
  and is expressed in J or kJ per degree change in
  temperature (K or oC).

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Coffee Cup Calorimetry

• A simple device for determining heat changes of
  aqueous reactions at constant pressure is a
  coffee cup calorimeter. Since the contents are
  open to the atmosphere, the pressure, atmospheric
  pressure, remains constant during the reaction.

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Coffee Cup Calorimetry

• The heat change for the reaction, qp, is equal
  to the enthalpy change for the reaction.
• If heat is given off, it goes towards warming
  up the contents of the calorimeter and toward
  warming up the calorimeter walls, thermometer,
  stirrer, etc.

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Coffee Cup Calorimetry

• qreaction qcontents qcal
• qcontents (mass of solution) (?Tsoln)Csoln
• Csoln is the specific heat capacity of the
  reaction mixture. If solutions are aqueous and
  fairly dilute, the specific heat capacity of
  water, 4.18J/oC-g, may be used.

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Coffee Cup Calorimetry

• qreaction qcontents qcal
• qcal Ccal (?T)
• Ccal is the calorimeter heat capacity. It
  includes the heat needed to warm up the walls,
  thermometer and stirrer of the calorimeter, along
  with any heat loss due to leaks.
• In many simple calculations, Ccal is assumed to
  be negligible, and may be ignored.

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Obtaining ?H of Reaction

• The enthalpy change of a reaction, ?H, can be
  obtained from qp.
• First, a sign must be assigned. If the
  temperature increased during the reaction, the
  reaction is exothermic, and q is negative.
• If the temperature decreased during the
  reaction, the reaction is endothermic, and q is
  positive.

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Obtaining ?H of Reaction

• The enthalpy change of a reaction, ?H, can be
  obtained from qp.
• Also, qp is for a specific quantity of
  reactants. Typically, ?Hrxn is for molar
  quantities of reactants. To calculate ?Hrxn from
  qp, you must calculate the heat change per mole
  of reactant.

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Problem Calorimetry

• A coffee cup calorimeter contains 125. grams of
  water at 24.2oC. A 10.5 g sample of KBr, also at
  24.2oC, is added. After dissolving, the mixture
  reaches a final temperature of 21.1oC. Calculate
  ?Hsoln in joules/gram and kJ/mol. Assume the
  specific heat of the solution is 4.18 J/g-oC, and
  no heat is transferred to or from the calorimeter
  or surroundings.

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Constant Volume Calorimetry

• Certain reactions, notably combustion reactions,
  do not lend themselves to open vessels. These
  reactions are usually carried out in a sealed
  reaction vessel called a bomb calorimeter.
• The bomb calorimeter is a rigid steel container
  that is sealed after the reactants have been
  added. The reaction takes place once an
  electrical current is sent through an ignition
  wire to the reaction mixture.

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Bomb Calorimetry

• The steel bomb is immersed in an insulated bath
  containing either water or mineral oil. As the
  combustion reaction releases heat, the heat is
  transferred to the bath.

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Bomb Calorimetry
Reaction vessel
Ignition wire
O2 inlet
gaskets
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Bomb Calorimetry

• Once the bomb has been charged with reactants,
  it is placed in the water or oil bath until it
  reaches a constant temperature.
• A current is sent through the ignition wire,
  and the combustion reaction takes place. The
  heat given off by the reaction is evident from
  the increase in temperature of the water/oil bath.

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Bomb Calorimetry

• The heat generated by the reaction warms up the
  contents of the calorimeter (the bomb,
  thermometer, container walls, stirrer) and the
  water (or oil) bath.
• Usually, the calorimeter constant, which is the
  heat capacity of the entire apparatus, is
  provided, or determined by combusting a substance
  with a known energy of combustion.

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Problem Bomb Calorimetry

• The energy released by combustion of benzoic acid
  is 26.42 kJ/g. The combustion of .1584g of
  benzoic acid increases the temperature of a bomb
  calorimeter by 2.54 oC.
• a) Calculate the calorimeter constant.

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Problem Bomb Calorimetry

• b) 0.2130 g of vanillin (C8H8P3) is burned in
  the same calorimeter with a temperature increase
  of 3.25oC. Calculate the energy of combustion of
  vanillin in kJ/g and kJ/mol.

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Hesss Law

• Enthalpy is a state function. This means that
  a change in enthalpy depends solely on the
  initial and final states (products and
  reactants), and is independent of the reaction
  pathway.

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Hesss Law

• Hesss Law is a method that combines related
  chemical reactions and their enthalpy changes.
  Since enthalpy changes are independent of
  pathway, as long as the net reaction matches the
  reaction of interest, the sum of the enthalpy
  changes will yield ?H for the desired reaction.

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Hesss Law
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Hesss Law

• There are a few basic rules in applying Hesss
  Law
• 1. If a reaction is reversed, the sign of ?H is
  also reversed.
• 2. If the coefficients in a balanced reaction
  are multiplied by an integer, the value of ?H is
  multiplied by the same integer.

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Problem Hesss Law

• Calculate ?Ho for the reaction
• C6H4(OH)2(aq)     H2O2(aq)? C6H4O2(aq)    2
  H2O(l)
• Using
• 1)C6H4(OH)2(aq) ?C6H4O2(aq)  H2(g) ?Ho 177.4
  kJ
• 2)  H2 (g)     O2(g)? H2O2(aq)               ?Ho
  191.2 kJ
• 3)  H2 (g)  1/2  O2(g) ? H2O(g)            ?Ho
  241.8 kJ
• 4)  H2O(g) ? H2O(l)                               
  ?Ho 43.8 kJ

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Standard Enthalpies of Formation

• A formation reaction involves combining
  elements, in their standard states, to form one
  mole of a compound.
• A table of standard enthalpies of formation
  (?Hfo) is in appendix of the text. The ?Hfo
  values of most common compounds have been
  determined and tabulated.

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Standard Enthalpies of Formation

• CaCO3(s) has a ?Hfo of -1207 kJ/mol. This is
  the enthalpy change for the reaction
• Ca(s) C(graphite) 3/2 O2(g) ? CaCO3(s)
• Fractional coefficients are acceptable since
  all quantities are molar, and a formation
  reaction produces one mole of a compound.

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Standard Enthalpies of Formation

• ?Hfo values can be used to calculate the
  standard enthalpy changes for many reactions.
• In an application of Hesss Law, it is as if
  the reactants are decomposed into their elements,
  and then the elements are recombined into the
  desired products. Since enthalpies of reaction
  are independent of pathway, this provides an
  accurate way to calculate enthalpies of reactions.

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Standard Enthalpies of Formation
CH4(g) 2O2(g) ? CO2(g) 2 H2O(l)

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Hesss Law and ?Hfo

• ?Hrxno ? ?Hfo(products) - ? ?Hfo(reactants)
• Problem Use standard heats of formation to
  calculate ?Horxn for
• KClO3(s) ? KCl(s) 3/2 O2(g)

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Bond Dissociation Energies

• Bond dissociation energies can be used to
  estimate the enthalpy of a reaction. The
  enthalpy change will approximately equal the
  energy of bonds broken energy of bonds formed.
• ?Ho D(bonds broken) D(bonds formed)

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Bond Energies

• Since bond energies are average values for a
  variety of molecules, they will only provide an
  estimate of the enthalpy change for a specific
  reaction. The following relationship can also be
  used to estimate enthalpies of reaction.
• ?Ho D(reactant bonds) D(product bonds)

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Predicting Spontaneity

• A process is spontaneous if it occurs with no
  outside intervention. An example is the melting
  of ice above a temperature of 0oC. Although the
  melting of ice is endothermic, it will still
  occur on its own if the temperature is high
  enough.
• The spontaneity of a process will depend on the
  enthalpy change, the temperature and the entropy
  change.

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Entropy

• Entropy is a measure of randomness or disorder.
  Spontaneous reactions or processes may involve
  an increase in entropy.
• In the example of melting ice, the liquid water
  is more random in structure than the solid. As a
  result, ?S, the entropy change, is positive.
• ?S Sfinal-Sinitial

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Entropy

• ?S Sfinal-Sinitial
• Entropy values for pure substances can be
  calculated, and typically have the units J/mol-K.
• Gases have higher entropy values than liquids,
  and mixtures have greater entropy than pure
  substances.

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Spontaneity

• A process will always be spontaneous if it
  releases heat and increases in entropy.
• A process will never be spontaneous if it
  absorbs heat and involves a decrease in entropy.
• For other combinations of enthalpy and entropy
  changes, temperature will play a role.

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Free Energy

• The free energy change, ?G, is used to predict
  if a process is spontaneous. It considers the
  enthalpy change, entropy change and temperature.
• ?G ?H - T?S

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Free Energy

• If ?G is negative, the process is spontaneous
  at the specified temperature. If positive, the
  process is not spontaneous, and if ?G 0, the
  system is at equilibrium.
• ?G ?H - T?S

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Free Energy

• At equilibrium, the forward process occurs at
  the same rate as the reverse process. Neither
  direction is favored.