Enthalpies of Formation The enthalpy of formation,  H f, 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 (  H f o ) 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 O 2 (g) + 3 H 2 (g) --> C 2 H 5 OH(l)  H f o = kJ

The standard enthalpy of formation of the most stable form of an element under standard conditions is ZERO. O 2 (g) --> O 2 (g)  H = 0 1/2 N 2 (g) + 3/2 H 2 (g) --> NH 3 (g)  H o f = kJ/mol

Using Enthalpies of Formation to calculate Standard Reaction Enthalpies

Combustion of propane (C 3 H 8 ) gas to form CO 2 (g) and H 2 O(l) C 3 H 8 (g) + 5 O 2 (g) --> 3CO 2 (g) + 4H 2 O(l)

Looking up the standard heats of formation for each equation  H o rxn = -( ) + 3(-393.5) + 4(-285.8)) = kJ This equation can be written as the sum of the following three equations C 3 H 8 (g) --> 3C(s) + 4H 2 (g)  H 1 = -  H f o (C 3 H 8 (g) ) + 3C(s) + 3O 2 (g) --> 3CO 2 (g)  H 2 = 3 x  H f o (CO 2 (g) ) + 4H 2 (g) + 2O 2 (g) --> 4H 2 O(l)  H 3 = 4 x  H f o (H 2 O (l) ) C 3 H 8 (g) + 5 O 2 (g) --> 3CO 2 (g) + 4H 2 O(l)  H o rxn =  H 1 +  H 2 +  H 3

In general,  H o rxn =  n  H f o (products) -  n  H f o (reactants) n is the stoichiometric coefficients in the reaction

Calculate the standard enthalpy change for the combustion of 1 mole of benzene (C 6 H 6 (l)) to CO 2 (g) and H 2 O(l). Compare the quantity of heat produced by the combustion of 1.00 g of propane (C 3 H 8 (g)) to that produced by 1.00 g of C 6 H 6 (l) First write a balanced equation for the combustion of 1 mole of C 6 H 6 (l) C 6 H 6 (l) + O 2 (g) --> 6CO 2 (g) + 3H 2 O(l) 15 2  H o rxn = [6  H f o (CO 2 ) + 3  H f o (H 2 O)] - [1  H f o (C 6 H 6 ) + (15/2)  H f o (O 2 )] = 6( kJ) + 3(285.8 kJ) kJ - 7.5(0 kJ) = kJ

For the combustion of 1 mole of propane  H o rxn = kJ Hence for 1.00g propane, which corresponds to mol propane,  H o rxn = mol x kJ/mol = kJ/g For C 6 H 6 (l) =>  H o rxn = kJ/g

Bond Enthalpies Strength of a chemical bond is measured by the bond enthalpy,  H B Bond enthalpies are positive, because heat must be supplied to break a bond. Bond breaking is endothermic Bond formation is exothermic. H 2 (g) --> 2 H  H o = +436 kJ  H B = 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. CH 3 CH 2 I(g) + H 2 O(g) --> CH 3 CH 2 OH(g) + HI(g) Reactant: break a C-I bond and an O-H bond  H o = 238 kJ kJ = 701 kJ Product: to form a C-O bond and an H-I bond  H o = -360 kJ kJ = -659 kJ Overall enthalpy change = 701 kJ kJ = 42 kJ

Fuels During the complete combustion of fuels, carbon is completely converted to CO 2 and hydrogen to H 2 O. C 3 H 8 (g) + 5 O 2 (g) --> 3CO 2 (g) + 4H 2 O(l) Standard heats of formation of CO 2 (g) and H 2 O(l)  H f o (CO 2 (g)) = kJ/mol  H f o (H 2 O(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 = J) Fuel TypeProven ReservesEst. Reserves Coal25Q118Q Oil5Q9Q Natural Gas4Q10Q 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) + O 2 (g) --> CO 2 (g)  H = kJ/mol CH 4 (g) + 2 O 2 (g) --> CO 2 (g) + 2 H 2 O(l)  H = -890 kJ/mol C 3 H 8 (g) + 5 O 2 (g) --> 3CO H 2 O  H = 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

NameHeat released per gram C(s)34 kJ CH 4 (g) 55.6 kJ C 3 H 8 (g)50.3 kJ NameHeat released per mole of CO 2 (g) released C(s)393.5 kJ CH 4 (g) 890 kJ C 3 H 8 (g)738 kJ CH 4 (g) and C 3 H 8 (g) release more energy per gram and can be considered to be “cleaner” fuels. Disadvantages: leakage of CH 4 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  H combustion (kJ/g) CH 3 OH(l)-22.7 C 2 H 5 OH (l)-29.7 CH 4 (g)-55.6 C(s)-34

Hydrogen H 2 (g) + 1/2O 2 (g) > H 2 O(l)  H = -286 kJ/mol spark Advantages of using H 2 as a fuel: energy released per gram low polluting Disadvantage: gas at room temperature H 2 /O 2 Fuel cells: Electrical energy is produced during the redox reaction

Methane (CH 4 ), Ethanol (C 2 H 5 OH), hydrogen (H 2 ) are “renewable” fuels. CH 4 : bacterial digestion of waste H 2 : electrolysis of ocean water C 2 H 5 OH: biological fermentation of starches (e.g. in corn) Combustion of CH 4 and C 2 H 5 OH produce CO 2, but they produce less CO 2 per gram than gasoline. And they are renewable. Compound  H o c Specific Enthalpy Enthalpy density kJ/mol kJ/g kJ/L Hydrogen (H 2 (g)) Methane (CH 4 (g)) Octane (C 8 H 18 (l)) x 10 4 Methanol (CH 3 OH(l)) x 10 4

Spontaneous Change A spontaneous change is one that occurs without external intervention and has definite direction. Spontaneous for T > 0 o C Spontaneous for T < 0 o C

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, NH 4 NO 3, 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>0 o C 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.