Enthalpies of Formation

The enthalpy of formation, Hf, or heat of formation, is defined as the change in enthalpy when one mole of a compound is formed from its stable elements.

The standard enthalpy of formation (Hfo) of a compound is

defined as the enthalpy change for the reaction that forms 1 mole of compound from its elements, with all substances in their standard states.

2C(s) + 1/2 O2(g) + 3 H2 (g) --> C2H5OH(l) Hfo = -277.69 kJ

The standard enthalpy of formation of the most stable form of an element under standard conditions is ZERO.

O2 (g) --> O2 (g) H = 0

1/2 N2 (g) + 3/2 H2 (g) --> NH3 (g) Hof = -46.19 kJ/mol

Using Enthalpies of Formation to calculate Standard Reaction Enthalpies

Combustion of propane (C3H8) gas to form CO2(g) and H2O(l)

C3H8 (g) + 5 O2 (g) --> 3CO2 (g) + 4H2O(l)

Looking up the standard heats of formation for each equation

Horxn = -(-103.85) + 3(-393.5) + 4(-285.8)) = -2220 kJ

This equation can be written as the sum of the following three equations

C3H8(g) --> 3C(s) + 4H2(g) H1 = - Hfo (C3H8(g) )

+ 3C(s) + 3O2(g) --> 3CO2(g) H2 = 3 x Hfo (CO2(g) )

+ 4H2(g) + 2O2(g) --> 4H2O(l) H3 = 4 x Hfo (H2O (l) )

C3H8 (g) + 5 O2 (g) --> 3CO2 (g) + 4H2O(l)

Horxn = H1 + H2+ H3

In general,

Horxn = n Hf

o (products) - n Hfo (reactants)

n is the stoichiometric coefficients in the reaction

Calculate the standard enthalpy change for the combustion of 1 mole of benzene (C6H6 (l)) to CO2(g) and H2O(l). Compare the quantity of heat produced by the combustion of 1.00 g of propane (C3H8(g)) to that produced by 1.00 g of C6H6 (l)

First write a balanced equation for the combustion of 1 mole of C6H6 (l)

C6H6 (l) + O2 (g) --> 6CO2 (g) + 3H2O(l)152

Horxn = [6 Hf

o(CO2) + 3Hfo(H2O)] - [1Hf

o(C6H6) + (15/2)Hf

o(O2)]

= 6(-393.5 kJ) + 3(285.8 kJ) - 49.0 kJ - 7.5(0 kJ)

= -3267 kJ

For the combustion of 1 mole of propane Horxn = -2220 kJ

Hence for 1.00g propane, which corresponds to 0.0227 mol

propane, Horxn = 0.0227mol x -2220 kJ/mol = - 50.3 kJ/g

For C6H6 (l) => Horxn = - 41.8 kJ/g

Bond Enthalpies

Strength of a chemical bond is measured by the bond enthalpy, HB

Bond enthalpies are positive, because heat must be supplied to break a bond.

Bond breaking is endothermic

Bond formation is exothermic.

H2(g) --> 2 H Ho = +436 kJ

HB = 436 kJ/mol

Mean bond enthalpy: average molar enthalpy change accompanying the dissociation of a given type of bond.

Estimate the enthalpy change of the reaction between gaseous iodoethane and water vapor.

CH3CH2I(g) + H2O(g) --> CH3CH2OH(g) + HI(g)

Reactant: break a C-I bond and an O-H bond

Ho = 238 kJ + 463 kJ = 701 kJ

Product: to form a C-O bond and an H-I bond

Ho = -360 kJ + -299 kJ = -659 kJ

Overall enthalpy change = 701 kJ - 659 kJ = 42 kJ

Fuels

During the complete combustion of fuels, carbon is completely converted to CO2 and hydrogen to H2O.

C3H8 (g) + 5 O2 (g) --> 3CO2 (g) + 4H2O(l)

Standard heats of formation of CO2(g) and H2O(l)

Hfo (CO2(g)) = -393.5 kJ/mol

Hfo(H2O(l)) = -286 kJ/mol

The greater the percentage of carbon and hydrogen in a fuel, the higher its fuel value.

US crude oil production

Hubbert’s Peak, K. S. Deffeyes

Global Energy Reserves (1988) (units of Q = 1021 J)

Fuel Type Proven Reserves Est. Reserves

Coal 25Q 118Q

Oil 5Q 9Q

Natural Gas 4Q 10Q

Total amount of commercially energy currently consumed by humans ~ 0.5Q annually

“Non-renewable” sources of energy

Alternate Fuels

Natural Gas and Propane

C(s) + O2(g) --> CO2(g) H = -393.5 kJ/mol

CH4(g) + 2 O2(g) --> CO2(g) + 2 H2O(l) H = -890 kJ/mol

C3H8(g) + 5 O2(g) --> 3CO2 + 4 H2O H = -2213 kJ/mol

Natural gas, primarily methane with small amounts of ethane and propane used for cooking and heating.

Highly compressed natural gas (CNG) - commercial vehicles.

Liquid petroleum gas (LPG) - propane - also used as a fuel for vehicles

Name Heat released per gramC(s) 34 kJCH4(g) 55.6 kJC3H8(g) 50.3 kJ

Name Heat released per mole of CO2(g) releasedC(s) 393.5 kJCH4(g) 890 kJC3H8(g) 738 kJ

CH4(g) and C3H8(g) release more energy per gram and can be considered to be “cleaner” fuels.

Disadvantages: leakage of CH4 from pipes, storage and transportation, need to be compressed

Methanol & Ethanol

Alcohols have the advantage over natural gas in that they are liquids at atmospheric pressure and temperature.

Compound Hcombustion (kJ/g)

CH3OH(l) -22.7

C2H5OH (l) -29.7

CH4(g) -55.6

C(s) -34

Hydrogen

H2(g) + 1/2O2(g) -------> H2O(l) H = -286 kJ/molspark

Advantages of using H2 as a fuel:

energy released per gram

low polluting

Disadvantage: gas at room temperature

H2/O2 Fuel cells: Electrical energy is produced during the redox reaction

Methane (CH4), Ethanol (C2H5OH), hydrogen (H2) are “renewable” fuels.

CH4: bacterial digestion of waste

H2 : electrolysis of ocean water

C2H5OH: biological fermentation of starches (e.g. in corn)

Combustion of CH4 and C2H5OH produce CO2, but they produce less CO2 per gram than gasoline. And they are renewable.

Compound Hoc Specific Enthalpy

Enthalpy density kJ/mol kJ/g kJ/LHydrogen (H2(g)) -286 -142 -13Methane (CH4(g)) -890 -55 -40Octane (C8H18(l)) -5471 -48 -3.8 x 104

Methanol (CH3OH(l)) -726 -23 -1.8 x 104

Spontaneous Change

A spontaneous change is one that occurs without external intervention and has definite direction.

Spontaneous for

T > 0oC

Spontaneous for

T < 0oC

A spontaneous process need not be fast

The change in enthalpy during a reaction is an important factor in determining whether a reaction is favored in the forward or reverse direction.

Are exothermic reaction more likely to be spontaneous than an endothermic reaction?

Not necessarily. The endothermic dissolution of ammonium nitrate, NH4NO3, occurs spontaneously.

Entropy

Both endothermic and exothermic reactions can be spontaneous

Are there additional factors which determine spontaneity?

Energy and matter tend to become more disordered.

A measure of disorder is ENTROPY.

When the valve is open, there are four possible arrangements or STATES for both particles.

Note: these arrangements are all equal in energy.

Opening the valve allows a higher degree of disorder.

The reverse process of the two gas particles occupying only one flask is not spontaneous.

As the number of particles increases in the system, the number of possible arrangements that the system can be in increases

Processes in which the disorder of the system increases tend to occur spontaneously.

Ice melts spontaneously at T>0oC even though it is an endothermic process.

The molecules of water that make up the ice crystal lattice are held rigidly in place.

When the ice melts the water molecules are free to move around, and hence more disordered than in the solid lattice.

Melting increases the disorder of the system.