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Title: Chapter 5: thermochemistry


1
Chapter 5thermochemistry
  • By Keyana Porter
  • Period 2
  • AP Chemistry

2
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

3
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

4
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

5
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

6
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

7
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

8
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)

9
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

10
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
11
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

12
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

13
5.4 Enthalpies 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

14
5.4 Enthalpies 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)
15
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

16
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)

17
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.)

18
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

19
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

20
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

21
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

22
5.6 Hesss 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

23
5.6 Hesss 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

24
5.6 Hesss Law
  • Solving the problem
  • 2CO2 (g) H2O (l) ? C2H2 (g) 5/2 O2 ?H
    1299.6 kJ
  • 2 x 2C (s) 2O2 (g) ? 2CO2 (g) ?H
    2(-393.5 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

25
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

26
5.7 Enthalpies 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

27
5.7 Enthalpies 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).

28
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

29
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
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