THERMOCHEMISTRY

ENERGY CHANGES ASSOCIATED WITH CHEMICAL

REACTION

ENERGY

• Capacity to do work or supply heat• Kinetic Energy: KE = 1/2 mv2 = energy due to

motion (v ≠ 0), Joule is the unit• Potential Energy: PE = stored energy due to

position, energy in a chemical bond (recall endo and exo Expt 1), Joule

• Energy is conserved • SI unit: Joule = kg (m/s)2; 1 calorie = 4.184 J

HEAT• Heat is the energy transfer between system

(chem rxn of reactants and products = focus of study) and surroundings (everything else) due to temperature difference, Joule

• q > 0 if heat absorbed by chem rxn; endothermic. Fig 6.3

• q < 0 if heat given off by chem rxn; exothermic. Fig 6.2

• Heat is a path function

WORK

• Work is the energy transferred between system and surroundings, Joule

• w = F · d = force that moves object a distance d • Consider work associated with gas expansion or

contraction: w = -P ΔV where P = external pressure• If w < 0, system does work on surroundings and system

loses energy; e.g. gas expands• If w > 0, surroundings does work on system and system

gains energy; eg. gas is compressed• Work is a path function• Note that 1.00 (L-atm) = 101.3 J

Figure 6.4 The Piston, Moving a Distance Against a Pressure P, Does Work On the Surroundings

FIRST LAW OF THERMODYNAMICS

• The energy of the universe is constant; in a physical or chemical change, energy is exchanged between system and surroundings, but not created nor destroyed.

• ΔE = internal energy = q + w = Efinal - Einitial

• If ΔV = 0, then ΔE = qV

• ΔE < 0, energy lost by system• ΔE > 0, energy gained by system

STATE FUNCTIONPATH FUNCTION

• State Function: A property of the system which depends only on the present state of the system and not the path used to get there; E, V, T

• Path Function; a property that depends on path taken during the change; w and q.

• Note ΔE = w + q is a constant for specific initial and final states even though q and w are path functions.

ENTHALPY

• If a chem rxn occurs at constant pressure (ΔP = 0) and only PV work occurs, then the heat associated with this rxn is called enthalpy, Joule

• H = enthalpy = state function, tabulated in Appendix 4

• H = E + PV; ΔH = ΔE + PΔV = qP

• ΔH = Hfinal - Hinitial = HP - HR

ENTHALPY (2)

• ΔH < 0 energy lost by system, exothermic• ΔH > 0 energy gained by system,

endothermic• Enthalpy depends on amount of substance

(I.e. #mol, #g); extensive property.• Chemical rxns are accompanied by enthalpy

changes (ΔH can be > 0 and < 0) that are measurable and unique.

Figure 6.2 Exothermic Process

Figure 6.3 Endothermic Process

Problems

• 24, 28, 30, 34, 36

THERMOCHEMICAL EQUATION

• Balanced chemical equation at a specific T and P includes reactants, products, phases and ΔH .

• Basis for stoichiometric problems that focus on ΔH associated with the chemical rxn.

• ΔH for reverse rxn = - ΔH for forward rxn

• If amount of reactants or products changes, then ΔH changes

CALORIMETRY

• Experimental method of determining heat (q) absorbed or released during a chem. rxn.

• Expts are either done at constant P (qP = ΔH) or constant V (qV = ΔE).

• This heat is proportional to the temp. change during the rxn: q = C ΔT where C is a constant and ΔT = Tfinal - Tinitial.

• C = heat capacity of the calorimeter; J/oC

CALORIMETRY (2)

• Here are two expressions of heat capacity• s = specific heat (capacity) = amount of

energy needed to raise the temp. of 1 g of material 1 oC; (units = J/oC-g) Table 6.1

• Cm = Molar Heat Capacity = amt of energy needed to raise temp. of 1 mol of sample 1 oC; (units = J/mol-oC)

• q = s m ΔT or q = Cm n ΔT

Table 6.1 The Specific Heat Capacities of Some Common Substances

Figure 6.5 A Coffee-Cup Calorimeter Made of Two Styrofoam Cups

Figure 6.6 A Bomb Calorimeter.

Problems

• 42, 46, 48, 54

THERMODYNAMIC STANDARD STATE

• The standard or reference state of a pure compound is its state at T = 25oC and – P = 1.00 atm for a gas or

– 1.00 M concentration for a solution.

• For an element, the std state is 1 atm and 25oC.• ΔHo = standard enthalpy of rxn or heat of rxn

when products and reactants are in their standard states.

PHYSICAL CHANGES

• There are ΔH values associated with phase or physical changes – Melting/freezing solid / liquid– Boiling/condensing liquid /

vapor– Subliming/condensing solid / vapor

• The former changes are endothermic; the latter are exothermic.

• Note that these changes are reversible.

HESS’S LAW: Law of Heat Summation

• Given a specific chem rxn at a stated T and P values, ΔH for the chem rxn is – constant and not dependent on intermediate chem

rxns.– the sum of the enthalpy changes for the

intermediate rxns. (Chem eqns are additive and their associated rxn ΔH values are additive).

• Hess’s Law facilitates the determination of rxn enthalpies for numerous rxns. (p 246)

The Principle of Hess’s Law

Stoichiometry and Thermochemical Equations

• Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

ΔH = -23 kJ

• 2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g)

ΔH = +23 kJ

• 2Fe2O3(s) + 6CO(g) 4Fe(s) + 6CO2(g)

ΔH = (2) -23 kJ = -46 kJ

Stoichiometry and Thermochemical Equations (2)

• Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

ΔH = -23 kJ per one mol Fe2O3(s) reacting

• Calculate the heat given off if 500 g of Fe2O3(s) reacts with excess CO.

g Fe2O3(s) mol Fe2O3(s) heat given off

STANDARD ENTHALPY OF FORMATION

• Enthalpy change for the formation of one mole of a substance in its standard state from its elements in their standard states

• ΔHof (1 atm and 25 oC) values are tabulated

in App. 4; note elements have ΔHof = 0.

• Combine ΔHof to calculate heat of rxn.

• ΔHorxn = ∑nPΔHo

f (prod.) - ∑nRΔHof (react.)

Table 6.2 Standard Enthalpies of Formation for Several Compounds at

25°C

Problems

• 58, 60, 66, 72

ENERGY SOURCES

• Variety of and emerging sources of energy and preparation of fuels

• Impact on the environment

• Combustion = type of reaction in which substance burns in oxygen.