Thermochemistry: Chemical Energy

1
Chapter 8

• Thermochemistry Chemical Energy

2
Energy

• Energy capacity to supply heat or do work
• Energy Heat Work
• E q w
• 2 types of Energy
• Potential Energy
• Kinetic Energy

3
Energy

• Two fundamental kinds of energy.
• Potential energy is stored energy.
• Kinetic energy is the energy of motion.
• Law of Conservation of Energy
• Energy can be converted from one kind to another
  but never destroyed

4
Energy

• Units
• SI Unit Joule (J)
• Additional units
• Calorie (Cal) food calorie
• calorie (cal) scientific calorie
• Conversions
• 1 cal 4.184 J
• 1000 cal 1 Cal

5
Energy and Chemical Bonds

• Chapter 6
• Kept a careful accounting of atoms as they
  rearranged themselves
• Reactions also involve a transfer of energy

6
Energy and Chemical Bonds

• A chemical
• Potential - attractive forces in an ionic
  compound or sharing of electrons covalent
  compound
• Kinetic (often in form of heat) occurs when
  bonds are broken and particles allowed to move
• To determine the energy of a reaction it is
  necessary to keep track of the energy changes
  that occur during the reaction

7
Internal Energy and State Functions

• In an experiment Reactants and products are the
  system everything else is the surroundings.
• Energy flow from the system to the surroundings
  has a negative sign (loss of energy).
• Energy flow from the surroundings to the system
  has a positive sign (gain of energy).

8
Internal Energy and State Functions

• Tracking energy changes
• Energy changes are measured from the point of
  view of the system (Internal Energy - IE)
• Change in Energy of the system ?E
• ?E Efinal - Einitial

9
Internal Energy and State Functions

• IE depends on
• Chemical identity, sample size, temperature, etc.
• Does not depend on the systems history
• Internal Energy is a state function
• A function or property whose value depends only
  on the present state (condition) of the system,
  not on the path used to arrive at that condition

10
Expansion Work

• E q w
• In physics w force (F) x distance (d)
• Force energy that produces movement of an
  object
• In chemistry w expansion work
• Force - the pressure that the reaction exerts on
  its container against atmospheric pressure hence
  it is negative
• Distance change in volume of the reaction
• w -P?V

11
Energy and Enthalpy

• ?E q P?V
• The amount of heat exchanged between the system
  and the surroundings is given the symbol q.
• q DE PDV
• At constant volume (DV 0) qv DE
• At constant pressure Energy due to heat and work
  but work minimal compared to heat energy
• qp DE PDV DH
• Enthalpy change (heat of reaction) DH
  Hproducts Hreactants

12
The Thermodynamic Standard State

• ?H amount of energy absorbed or released in the
  form of heat
• DH Hproducts Hreactants
• Important factors
• States of matter
• Thermodynamic standard state most stable form
  of a substance at 1 atm and at a specified
  temperature, usually 25oC and 1 M concentration
  for all substances in solution
• DH valid for the reaction as written including
  exact of moles of substances
• N2H4(g) H2(g) ? 2 NH3(g) heat (188 kJ)

13
Enthalpies of Physical and Chemical Change
14
Enthalpies of Physical and Chemical Changes

• Enthalpies of Chemical Change Often called heats
  of reaction (DHreaction).
• Endothermic Heat flows into the system from the
  surroundings and DH has a positive sign.
  Unfavorable Process
• Exothermic Heat flows out of the system into the
  surroundings and DH has a negative sign.
  Favorable process

15
Enthalpies of Physical and Chemical Changes

• Reversing a reaction changes the sign of DH for a
  reaction.
• C3H8(g) 5 O2(g) ? 3 CO2(g) 4 H2O(l) DH
  2219 kJ
• 3 CO2(g) 4 H2O(l) ? C3H8(g) 5 O2(g) DH
  2219 kJ
• Multiplying a reaction increases DH by the same
  factor.
• 3 C3H8(g) 15 O2(g) ? 9 CO2(g) 12 H2O(l) DH
  6657 kJ

16
Problems

• How much heat (in kilojoules) is evolved or
  absorbed in each of the following reactions?
• Burning of 15.5 g of propane C3H8(g) 5 O2(g)
  ? 3 CO2(g) 4 H2O(l) DH 2219 kJ
• Reaction of 4.88 g of barium hydroxide
  octahydrate with ammonium chloride Ba(OH)28
  H2O(s) 2 NH4Cl(s) ? BaCl2(aq) 2 NH3(aq) 10
  H2O(l) DH 80.3 kJ

17
Determination of Heats of Reaction

• Experimentally calorimetry
• Hesss Law
• Standard Heats of Formation
• Bond Dissociation Energies

18
Calorimetry and Heat Capacity

• Calorimetry is the science of measuring heat
  changes (q) for chemical reactions. There are
  two types of calorimeters
• Bomb Calorimetry A bomb calorimeter measures
  the heat change at constant volume such that q
  DE.
• Constant Pressure Calorimetry A constant
  pressure calorimeter measures the heat change at
  constant pressure such that q DH.

19
Calorimetry and Heat Capacity
20
Calorimetry and Heat Capacity

• Heat capacity (C) is the amount of heat required
  to raise the temperature of an object or
  substance a given amount.
• Specific Heat The amount of heat required to
  raise the temperature of 1.00 g of substance by
  1.00C.
• Molar Heat The amount of heat required to raise
  the temperature of 1.00 mole of substance by
  1.00C.

21
Problems

• What is the specific heat of lead if it takes 96
  J to raise the temperature of a 75 g block by
  10.0C?
• When 25.0 mL of 1.0 M H2SO4 is added to 50.0 mL
  of 1.0 M NaOH at 25.0C in a calorimeter, the
  temperature of the solution increases to 33.9C.
  Assume specific heat of solution is 4.184
  J/(g1C1), and the density is 1.00 g/mL1,
  calculate the heat absorbed or released for this
  reaction.

22
Hesss Law

• Allows the enthalpy to be determined for
• Reactions that occur too quickly or take too long
  to use calorimetry
• Reactions that are too dangerous
• Works like the Haber process in chapter 6
• Take reactions for which the heat is known and
  manipulate them to give the desired reaction

23
Standard Heats of Formation

• Standard Heats of Formation (DHf) The enthalpy
  change for the formation of 1 mole of substance
  in its standard state from its constituent
  elements in their standard states.
• The standard heat of formation for any element in
  its standard state is defined as being ZERO.
• DHf 0 for an element in its standard state

24
Standard Heats of Formation

• H2(g) 1/2 O2(g) ? H2O(l) DHf 286 kJ/mol
• 3/2 H2(g) 1/2 N2(g) ? NH3(g) DHf 46 kJ/mol
• 2 C(s) H2(g) ? C2H2(g) DHf 227 kJ/mol
• 2 C(s) 3 H2(g) 1/2 O2(g) ? C2H5OH(g) DHf
  235 kJ/mol

25
Standard Heats of Formation

• Calculating DH for a reaction
• DH SDHf (products) x moles SDHf
  (Reactants) x moles
• For a balanced equation, each heat of formation
  must be multiplied by the stoichiometric
  coefficient.
• aA bB cC dD
• DH cDHf (C) dDHf (D) aDHf (A)
  bDHf (B)

26
Problems

• Calculate DH (in kilojoules) for the reaction of
  ammonia with O2 to yield nitric oxide (NO) and
  H2O(g), a step in the Ostwald process for the
  commercial production of nitric acid.
• Calculate DH (in kilojoules) for the
  photosynthesis of glucose from CO2 and liquid
  water, a reaction carried out by all green
  plants.

27
Energy Calculations

• Other methods for calculating enthalpies
• Bond dissociation energies measures the energy
  given off by the formation of bonds in the
  products and substracts the energy required to
  break bonds in the reactants

28
Why do chemical reactions occur?

• A chemical reaction will move from less stability
  to greater stability.
• Achieved by giving off more energy than is
  absorbed by the reactants
• This indicates that exothermic reactions occur by
  why do endothermic reactions occur?
• Gibbs Free Energy
• DG DH TDS
• DH enthalpy, T temperature, DS - entropy

29
An Introduction to Entropy

• Second Law of Thermodynamics Reactions proceed
  in the direction that increases the entropy of
  the system plus surroundings. (increases the
  degree of disorder)
• A spontaneous process is one that proceeds on its
  own without any continuous external influence.
• A nonspontaneous process takes place only in the
  presence of a continuous external influence.

30
An Introduction to Entropy
31
An Introduction to Entropy
32
An Introduction to Entropy

• The measure of molecular disorder in a system is
  called the systems entropy this is denoted S.
• Entropy has units of J/K (Joules per Kelvin).
• DS Sfinal Sinitial
• Positive value of DS indicates increased disorder
  (favorable).
• Negative value of DS indicates decreased disorder
  (unfavorable).

33
Problems

• Predict whether DS is likely to be positive or
  negative for each of the following reactions.
  Using tabulated values, calculate DS for each
• a. 2 CO(g) O2(g) ? 2 CO2(g)b. 2 NaHCO3(s) ?
  Na2CO3(s) H2O(l) CO2(g)c. C2H4(g) Br2(g) ?
  CH2BrCH2Br(l)d. 2 C2H6(g) 7 O2(g) ? 4 CO2(g)
  6 H2O(g)

34
An Introduction to Free Energy

• To decide whether a process is spontaneous, both
  enthalpy and entropy changes must be considered
• Spontaneous process Decrease in enthalpy
  (DH). Increase in entropy (DS).
• Nonspontaneous process Increase in enthalpy
  (DH). Decrease in entropy (DS).

35
An Introduction to Free Energy

• Gibbs Free Energy Change (DG) Weighs the
  relative contributions of enthalpy and entropy to
  the overall spontaneity of a process.
• DG DH TDS
• DG lt 0 Process is spontaneous (favorable)
• DG 0 Process is at equilibrium
• DG gt 0 Process is nonspontaneous (unfavorable)

36
Problems

• Which of the following reactions are spontaneous
  under standard conditions at 25C?
• a. AgNO3(aq) NaCl(aq) ? AgCl(s) NaNO3(aq)
  DG 55.7 kJ
• b. 2 C(s) 2 H2(g) ? C2H4(g) DG 68.1 kJ
• c. N2(g) 3 H2(g) ? 2 NH3(g) DH 92 kJ
  DS 199 J/K

37
An Introduction to Free Energy

• Equilibrium (DG 0) Estimate the temperature
  at which the following reaction will be at
  equilibrium. Is the reaction spontaneous at room
  temperature?
• N2(g) 3 H2(g) ? 2 NH3(g) DH 92.0 kJ
  DS 199 J/K
• Equilibrium is the point where DG DH TDS
  0

38
Problem

• Benzene, C6H6, has an enthalpy of vaporization,
  DHvap, equal to 30.8 kJ/mol and boils at 80.1C.
  What is the entropy of vaporization, DSvap, for
  benzene?

39
Optional Homework

• Text - 8.28, 8.32, 8.50, 8.52, 8.56, 8.58, 8.66,
  8.70, 8.74, 8.82, 8.88, 8.90
• Chapter 8 Homework from website

40
Required Homework

• Chapter 8 Assignment