Chapter 5: thermochemistry

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Chapter 5thermochemistry

• By Keyana Porter
• Period 2
• AP Chemistry

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What is thermochemistry?

• Thermodynamics the study of energy and its
  transformations
• Thermochemistry is one aspect of thermodynamics
• The relationship between chemical reactions and
  energy changes
• Transformation of energy (heat) during chemical
  reactions

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5.1Kinetic Potential Energy

• Energy is the capacity to do work or the transfer
  heat
• Objects possess energy in 2 ways
• Kinetic energy due to motion of object
• Potential/Stored energy result of its
  composition or its position relative to another
  object
• Kinetic energy (Ek) of an object depends on its
  mass (m) and speed (v)
• Ek ½ mv2
• Kinetic energy increases as the speed of an
  object and its mass increases
• Thermal energy energy due to the substances
  temperature associated with the kinetic energy
  of the molecules

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5.1Kinetic Potential Energy

• Potential energy is a result of attraction
  repulsion
• Ex an electron has potential energy when it is
  near a proton due to the attraction
  (electrostatic forces)
• Chemical energy due to the stored energy in the
  atoms of the substance

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5.1Units of Energy

• joule (J) SI unit for energy
• 1 kJ 1000 J
• calorie (cal) non-SI unit for energy amount of
  energy needed to raise the temperature of 1 g of
  water by 1oC
• 1 cal 4.184 J (exactly)
• Calorie (nutrition unit) 1000 cal 1 kcal
• A mass of 2 kg moving at a speed of 1 m/s
  kinetic energy of 1 J
• Ek ½ mv2 ½ (2 kg)(1 m/s)2 1 kg-m2/s2 1 J

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5.1System Surroundings

• System (chemicals) portion that is singled out
  of the study
• Surroundings (container and environment including
  you) everything else besides the system
• Closed system can exchange energy, in the form
  of heat work, but not matter with the
  surroundings

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5.1Transferring Energy Work Heat

• Energy is transferred in 2 ways
• Cause the motion of an object against a force
• Cause a temperature change
• Force (F) any kind of push or pull exerted on an
  object
• Ex gravity
• Work (w) energy used to cause an object to move
  against force
• Work equals the product of the force and the
  distance (d) the object is moved
• w F x d
• heat the energy transferred from a hotter object
  to a colder one
• Combustion reactions release chemical energy
  stored in the form of heat

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5.2Internal Energy

• The First Law of Thermodynamics Energy is
  conserved it is neither created nor destroyed
• Internal energy (E) sum of ALL the kinetic and
  potential energy of all the components of the
  system
• The change in internal energy the difference
  between Efinal Einitial
• ?E Efinal Einitial
• We can determine the value of ?E even if we
  dont know the specific values of Efinal and
  Einitial
• All energy quantities have 3 parts
• A number, a unit, and a sign (exothermic versus
  endothermic)

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5.2Relating ?E to Heat Work

• A chemical or physical change on a system, the
  change in its internal energy is given by the
  heat (q) added to or given off from the system
• ?E q w
• both the heat added to and the work done on the
  system increases its internal energy

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Sign Conventions Used and the Relationship Among
q, w, and ?E
Sign Convention for q q gt 0 Heat is transferred from the surroundings to the system (endothermic) q lt 0 Heat is transferred from the system to the surroundings (exothermic)
Sign convention for w w gt 0 Work is done by the surroundings on the system w lt 0 Work is done by the system on the surroundings
Sign of ?E q w q gt 0 and w gt 0 ?E gt 0 q gt 0 and w lt 0 the sign of ?E depends on the magnitudes of q and w q lt 0 and w gt 0 the sign of ?E depends on the magnitudes of q and w q lt 0 and w lt 0 ?E lt 0
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5.2Endothermic Exothermic Processes

• Endothermic system absorbs heat
• Ex melting of ice
• Exothermic system loses heat and the heat flows
  into the surroundings
• Ex freezing of ice
• The internal energy is an example of a state
  function
• Value of any state function depends only on the
  state or condition of the system (temperature,
  pressure, location), not how it came to be in
  that particular state
• ?E q w but, q and w are not state functions

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5.3Enthalpy

• Enthalpy (H) state function the heat absorbed
  or released under constant pressure
• The change in enthalpy equals the heat (qP)
  gained or lost by the system when the process
  occurs under constant pressure
• ?H Hfinal Hinitial qP
• only under the condition of constant pressure is
  the heat that is transferred equal to the change
  in the enthalpy
• The sign on ?H indicated the direction of heat
  transfer
• value of ?H means it is endothermic
• - value of ?H means it is exothermic

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5.4Enthalpies of Reaction

• Enthalpy of reaction (?Hrxn) the enthalpy change
  that accompanies a reaction
• The enthalpy change for a chemical reaction is
  given by the enthalpy of the products minus the
  reactants
• ?H H(products) H(reactants)
• Thermochemical equations balanced chemical
  equations that show the associated enthalpy
  change
• The magnitude of ?H is directly proportional to
  the amount or reactant consumed in the process

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5.4Enthalpies of Reaction

• The enthalpy change for the reaction is equal in
  magnitude but opposite in sign to ?H for the
  reverse reaction
• CO2(g) 2H2O(l) ? CH4(g) 2O2(g) ?H 890 kJ

CH4(g) 2O2(g)
Enthalpy ?
Reversing a reaction changes the sign but not the
magnitude of the enthalpy change ?H2 - ?H1
?H1 - 890 kJ
?H2 890 kJ
CO2(g) 2H2O(l)
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5.4Enthalpies of Reaction

• The enthalpy change for a reaction depends on the
  state of the reactants and products
• Ex CH4 (g) 2O2 (g) ? CO2 (g) 2H2O (l)
• ?H -890 kJ
• If the product was H2O (g) instead of H2O (l),
  the ?H would be - 820 kJ instead of - 890 kJ

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5.5Calorimetry/Heat Capacity Specific Heat

• Calorimetry the measure of heat flow
• Calorimeter measures heat flow
• Heat capacity the amount of heat required to
  raise its temperature by 1 K
• The greater the heat capacity of a body, the
  greater the heat required to produce a given rise
  in temperature
• Molar heat capacity the heat capacity of 1 mol
  of a substance
• Specific heat the heat capacity of 1 g of a
  substance measured by temperature change (?T)
  that a known mass (m) of the substance undergoes
  when it gains or loses a specific quantity of
  heat (q)

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5.5Calorimetry/Heat Capacity Specific Heat

• specific heat quantity of heat transferred
• (grams of substance) x (temperature
  change)
• q
• m x ?T
• Practice Exercise (BL page 160)
• Calculate the quantity of heat absorbed by 50 kg
  of rocks if their temperature increases by 12.0
  OC. (Assume the specific heat of the rocks is .82
  J/g-K.)

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5.5Calorimetry/Heat Capacity Specific Heat

• Solving the problem
• q (specific heat) x (grams of substance) x ?T
• (0.82 J/g-K)(50,000 g)(285 K)
• 4.9 x 105 J

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5.5Constant-Pressure Calorimetry

• The heat gained by the solution (qsoln) is equal
  in magnitude and opposite in sign from the heat
  of the reaction
• qsoln (specific heat of solution) x (grams of
  solution) x ?T - qrxn
• For dilute aqueous solutions, the specific heat
  of the solution is approx. the same as water
  (4.18 J/g-K)
• Practice Exercise (BL page 161)
• When 50 mL of .100M AgNO3 and 50.0 mL of .100 M
  HCl are mixed in a constant-pressure calorimeter,
  the temperature of the mixture increases from
  22.30oC to 23.11oC. The temperature increase is
  caused by this reaction
• AgNO3 HCl ? AgCl HNO3
• Calculate ?H for this reaction, assuming that the
  combined solution has a mass of 100 g and a
  specific heat of 4.18 J/g-oC

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5.5Constant-Pressure Calorimetry

• Solving the problem
• qrxn -(specific heat of solution) x (grams of
  solution) x ?T
• - (4.18 J/g-oC)(100 g)(0.8 K)
• - 68,000 J/mol

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5.5Bomb Calorimetry (Constant-Volume Calorimetry)

• Bomb calorimeter used to study combustion
  reactions
• Heat is released when combustion occurs, absorbed
  by the calorimeter contents, raising the
  temperature of the water (measured before and
  after the reaction)
• To calculate the heat of combustion from the
  measured temperature increase in the bomb
  calorimeter, you must know the heat capacity of
  the calorimeter (Ccal)
• qrxn - Ccal x ?T

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5.6Hesss Law

• Hesss Law if a reaction is carried out in a
  series of steps, ?H for the reaction will be
  equal to the sum of the enthalpy changes for the
  individual steps
• CH4 (g) 2O2 (g) ? CO2 (g) 2H2O (g) ?H -802
  kJ
• (ADD)2H2O (g) ? 2H2O (l) ?H -88 kJ
• CH4 (g) 2O2 (g) 2H2O (g)? CO2 (g) 2H2O (g)
  2H2O (l)
• ?H -890 kJ

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5.6Hesss Law

• Practice Exercise 5.8 (BL page 165)
• Calculate ?H for the reaction
• 2C (s) H2 ? C2H2 (g)
• Given the following reactions and their
  respective enthalpy changes
• C2H2 (g) 5/2 O2 ? 2CO2 (g) H2O (l) ?H
  -1299.6 kJ
• C (s) O2 (g) ? CO2 (g) ?H -393.5 kJ
• H2 (g) ½ O2 (g) ? H2O (l) ?H -285.8 kJ

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5.6Hesss Law

• Solving the problem
• 2CO2 (g) H2O (l) ? C2H2 (g) 5/2 O2 ?H
  1299.6 kJ
• 2C (s) 2O2 (g) ? 2CO2 (g) ?H -393.5 kJ
  (2)
• -787.0 kJ
• H2 (g) ½ O2 (g) ? H2O (l) ?H -285.8 kJ
• 2C (s) H2 ? C2H2 (g) ?H 226.8 kJ
• if the reaction is reversed, the sign of ?H
  changes
• if reaction is multiplied, so is ?H

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Enthalpy Diagram

• The quantity of heat generated by combustion of 1
  mol CH4 is independent of whether the reaction
  takes place in one or more steps
• ?H1 ?H2 ?H3

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5.7Enthalpies of Formation

• Enthalpies of vaporization ?H for converting
  liquids to gases
• Enthalpies of fusion ?H for melting solids
• Enthalpies of combustion ?H for combusting a
  substance in oxygen
• Enthalpy of formation (?Hf) enthalpy change
  where the substance has been formed from its
  elements
• Standard enthalpy (?H o) enthalpy change when
  all reactants and products are at 1 atm pressure
  and specific temperature (298 K)
• Standard enthalpy of formation (?Hof) the
  enthalpy change for the reaction that forms 1 mol
  of the compound from its elements, with all
  substances in their standard states
• ?Hof 0 for any element in its purest form at
  295 K and 1 atm pressure

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5.7Enthalpies of Formation

• The standard enthalpy change for any reaction can
  be calculated from the summations of the
  reactants and products in the reaction
• ?H orxn n ?H of (products) - m?H
  of (reactants)
• Practice Problem 5.9 (BL page 169)
• Calculate the standard enthalpy change for the
  combustion of 1 mol of benzene, C6H6 (l), to CO2
  (g) and H2O (l).

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5.7Enthalpies of Formation

• Solving the problem
• C6H6 (l) 15/2 O2 (g) ? 6CO2 (g) 3H2O (l)
• ?H orxn 6? H of (CO2) 3? H of (H2O) ?H of
  (C6H6) 15/2 ?H of (O2)
• 6(-393.5 kJ) 3(-285.8 kJ) (49.0 kJ)
  15/2 (0 kJ)
• (-2361 857.4 49.0) kJ
• -3267 kJ

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Extra Equations

• Force mass x 9.8 m/s2
• Internal energy ?E Efinal Einitial
• Entropy ? S Sfinal Sinitial
• Enthalpy ? H Hfinal Hinitial
• Gibbs Free Energy ? G Gfinal Ginitial
• ? S ? H / T