Energy and Chemical Reactions

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Energy and Chemical Reactions

• Chapter 6

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Thermodynamics

• The study of energy and its transformations
• The study of the relationships between chemical
  reactions and energy

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Work, Heat and Energy

• Work - the amount of force applied to an object
  over a distance (w F d)
• Heat - the energy transferred from a hotter
  object to colder one
• Energy - the capacity to do work or transfer
  heat. Can either be kinetic or potential

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Kinetic Energy

• Kinetic energy an object has kinetic energy
  because it is moving therefore, it is known as
  the energy of motion
• Examples thermal, mechanical, electrical and
  sound
• Ek 1/2 mv2

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Potential Energy

• Potential energy - the energy an object possess
  by virtue of its position. Stored energy that
  can be converted into kinetic energy
• Examples chemical, gravitational, electrostatic

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• Types of Energy
• Thermal energy is the energy associated with the
  random motion of atoms and molecules
• Chemical energy is the energy stored within the
  bonds of chemical substances
• Nuclear energy is the energy stored within the
  collection of neutrons and protons in the atom
• Electrical energy is the energy associated with
  the flow of electrons

6.1
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Units of Energy

• SI unit joule (J) (1 kg)(m2/s2)
• 1 cal 4.184 J (exactly)
• The nutritional Calorie (Cal) is 1000 cal

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System and Surroundings

• A system is the part of the universe we are
  interested in studying.
• Surroundings are the rest of the universe (the
  portions of the universe not involved in the
  system)

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Internal Energy, E

• The total energy of a system found as the sum of
  all the kinetic and potential energies of all
  components of the system.
• Absolute internal energy cannot be measured, only
  changes in internal energy.
• Change in internal energy ?E Efinal - Einitial.

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First Law of Thermodynamics

• When a system undergoes a physical or chemical
  change, the change in internal energy (?E) is
  given by the heat added to or absorbed by the
  system (q) plus the work done on or by the system
  (W) ?E q W

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Heat and Work Signs

• Heat flowing from the surrounding to the system
  is positive ( i.e. The system feels cold to the
  touch because it is absorbing heat from your
  hand) q gt 0
• Work done by the surrounding on the system is
  positive W gt 0

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• When heat is added to a system and work is done
  on a system by the surroundings, the change in
  the internal energy of the system is positive ?E
  gt 0

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Endothermic and Exothermic Processes

• An endothermic process is one that absorbs heat
  from the surrounding (it feels cold).
• An exothermic process is one that transfers heat
  to the surrounding (it feels hot).

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Exothermic process is any process that gives off
heat transfers thermal energy from the system
to the surroundings.
Endothermic process is any process in which heat
has to be supplied to the system from the
surroundings.
6.2
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State Functions

• Depends only on the initial and final states of a
  system
• It does not depend on how the internal energy is
  used.
• Example the internal energy of 50g of water at
  25ºC does not matter whether you heat the water
  from 0ºC or cool it from 100º

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Enthalpy of Reactions

• The heat transferred between the system and the
  surrounding during chemical reactions carried out
  under constant pressure is call enthalpy.
• Again, we can only measure the change in
  enthalpy, ?H.

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DH

• Mathematically, ?Hrxn H(products) -
  H(reactants)
• qp
• The subscript on the q only tells us that the
  heat was transferred under constant pressure.

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DH cont.

• For a reaction, Heat transferred from the
  surrounding to the system has a positive enthalpy
  (?H gt 0 for an endothermic reaction)
• Heat transferred from the system to the
  surrounding has a negative enthalpy (?H lt 0 for
  an exothermic reaction)
• Enthalpy is a state function.

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Enthalpy and Stoichiometry

• The magnitude of enthalpy is directly
  proportional to the amount of reactant consumed.
• The enthalpy changes for a reaction and its
  reverse reaction are equal in magnitude but
  opposite in sign.
• Enthalpy change depends on state.

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Example

• The enthalpy change for the formation of water
  from its elements is -285.8 KJ/mol.
• How much energy is evolved if 100.0 grams of
  water are formed?

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Hesss Law

• One of the ways in which the enthalpy change for
  a reaction can be determined is by knowing energy
  changes for related reactions. Hess's Law allows
  us to do this.

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Hesss Law

• If a chemical equation is the sum of multiples of
  other equations, the DH of this equation equals a
  similar sum of multiples of DH's for the other
  equations.

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Example

• Let's say we wanted to know the amount of heat
  that was evolved or absorbed when a mole of
  carbon monoxide was made from its elements.
• (The enthalpy change for a reaction such as this
  is called the heat of formation and is given the
  symbol DHf.)

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Example

• If we knew the enthalpy change for this reaction
• CO2(g) ---------------gt CO(g) 1/2O2 (g)
• DH 283.0 KJ
• and we knew the enthalpy change for this
  reaction
• C(s) 1O2 (g) ---------------gt CO2 (g) DH
  -393.5 KJ
• We could add them together and get the enthalpy
  change for the formation of CO(g)
• C(s) 1/2O2 (g) ---------------gt CO(g) DH
  -110.5 KJ

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Example

• CO2(g) ------gt CO(g) 1/2O2(g) DH 283.0
  KJ
• C(s) O2(g) -----gt CO2(g) DH
  -393.5 KJ
• __________________________________________________
  ____
• C(s) 1/2O2(g) ------gt CO(g) DH
  -110.5 KJ

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• When applying Hess's Law to obtain the DH for a
  reaction, the known (given) reactions can be
  multiplied by any factor including a -1 which
  reverses them. The corresponding DH is
  multiplied by the same factor.

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Example

• If we wanted to find the DH for the following
  reaction
• 2SO2(g) O2 (g) ----------gt SO3(g)
• and, we were given the following reactions
• SO2 (g) ----------gt S(s) O2 (g) DH 296.8
  KJ
• 2SO3(g) ----------gt 2S(s) 3O2 (g)DH
  791.4 KJ
• In order to obtain the reaction we want we would
  have to multiply the first reaction by 2 and we
  would have to reverse the second reaction. In
  each case we have to do the same operation on the
  DH values.

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The Result

• 2SO2(g) -----gt 2S(s) 2O2 (g) DH
  2(296.8 KJ)
• 493.6 KJ
• 2S(s) 3O2 (g)----gt2SO3 (g) DH -1(791.4 KJ)
• -791.4 KJ
• ______________________________________________
• 2SO2 (g) O2 (g) ----------gt SO3(g) DH
  -197.8 KJ

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Standard Molar Enthalpies of Formation

• If a compound is formed from its constituent
  elements, then the enthalpy change for the
  reaction is called the enthalpy of formation.
• Standard conditions (standard state) refer to the
  substance at 1 atm pressure and 25ºC (298 K)
• Standard enthalpy , ?Hº, is the enthalpy measured
  when everything is in its standard state.

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• Standard enthalpy of formation of a compound,
  ?Hºf , is the enthalpy change for the formation
  of 1 mole of compound with all substances in
  their standard states. All elements have an
  enthalpy of formation equal to zero.

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Enthalpy Change for a Reaction

• DHrxn
• ? DHf products - ? DHf reactants

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Why it works

• CH4(g) 4Cl2(g) -----? CCl4(l) 4HCl(g)

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Example
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Calorimetry

• The measurement of heat flow
• A calorimeter is the device used to measure heat
  flow.

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Heat Capacity and Specific Heat

• Heat capacity is the amount of energy required to
  raise the temperature of an object by 1ºC.
• Molar heat capacity is the heat capacity of 1 mol
  of a substance.
• Specific heat, or specific heat capacity, is the
  amount of energy required to raise the
  temperature of 1 g of substance by 1ºC.

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Calculating Heat

• Heat q
• q (specific heat) (grams of substance) ?T

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The specific heat (s) of a substance is the
amount of heat (q) required to raise the
temperature of one gram of the substance by one
degree Celsius.
The heat capacity (C) of a substance is the
amount of heat (q) required to raise the
temperature of a given quantity (m) of the
substance by one degree Celsius.
C ms
Heat (q) absorbed or released
q msDt
q CDt
Dt tfinal - tinitial
6.4
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Constant-Pressure Calorimetry

• The most common technique use atmospheric
  pressure as the constant pressure
• Recall that ?H q
• Easiest method use a coffee-cup calorimeter
• qsoln (specific heat of solution) (grams of
  solution) ?T -qrxn

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Example
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Bomb Calorimetry (Constant Volume)

• Reactions can be carried out under conditions of
  constant volume instead of constant pressure.
• Constant-volume calorimetry is carried out in a
  bomb calorimeter.
• The most common type of reaction studied under
  these conditions is combustion.

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• If we know the heat capacity of the calorimeter,
  Ccal, then the heat of reaction, qrxn -Ccal
  ?T
• Since the reaction is carried out under constant
  volume, q corresponds to ?E rather than ?H (for
  most reactions, the difference in ?E and ?H is
  small).

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Bomb Calorimeter Calculations

• -qrxn qbomb qwater
• -qrxn Ccal ?T S.H. mass ?T
• -qrxn (Ccal S.H. mass) ?T

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Example
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Notable Facts about Enthalpy

• Enthalpies for reactions can be given special
  names
• Formation Reaction - All reactants and elements,
  one mole of product compound -DHf
• Combustion Reaction - Something combining with
  oxygen - generally a hydrocarbon - DHcomb

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More Notable Facts

• Equations to determine the enthalpy of a reaction
  can be rearranged to determine any of the
  components of the equation.
• The enthalpy of formation of all elements in
  their standard states is equal to zero.
• The enthalpy of formation of H(aq) is also equal
  to zero.