1. 10-4 Enthalpy (Section 10.6)
And you

2. Enthalpy, symbolized by H, can be thought of as the potential energy stored in the bonds of molecules. Chemists use the change in enthalpy ∆H to measure the heat content of a system (when the pressure is constant).
We define the “system” to be the chemicals and everything else is termed the “surroundings”.
Applying the First Law of Thermodynamics (Conservation of Energy), any heat lost by the system will equal the heat gained by the surroundings (and vice versa).
Exothermic (“exo” means released or “exits” and therm refers to heat)

4. Exothermic reactions characteristics
Reactants: high E (H), less stable, weak bonds
Products: low E (H), more stable, strong bonds
System releases PE from bonds to KE of surroundings (which feel hot).
∆H = P – R = negative value (heat released)

5. Endothermic (“endo” means absorbed or “going in” and therm refers to heat)

6. Characteristics: Reactants: low E (H), more stable, strong bonds
Products: high E (H), less stable, weak bonds
System absorbs KE from surroundings as PE in the bonds. Surroundings will feel cold.
∆H = P – R = positive value (heat absorbed)

7. Bond Energies and ∆H:
It requires energy to break the bonds of the reactants. It releases energy when new bonds of the products form. The difference between these two energies is the ∆H. Note, though, if:
Energy absorbed to break reactants > Energy released forming products Endo ∆H = + 
Energy absorbed to break reactants < Energy released forming products Exo ∆H = -

8. Example ½ H2 + ½ Cl2  HCl Think of bond energies as KE entering or
leaving the system.
 The change in KE = the
change in PE
 91
H = -91kJ/mol Exothermic

9. 10-5 Enthalpy of Formation
The enthalpy of formation, ∆Hf, is defined as the heat absorbed or released when making 1 mole of a compound from its elements (at 25oC and 1 atm = standard state). Note that the conditions are important!
By convention, the Hf of any element at this temperature and pressure is zero.
ex: O2 (g)

10. what does the balanced reaction look like? N2 (g) + 3H2 (g) = 2NH3 (g)
Example reaction 1: Nitrogen gas reacts with hydrogen gas to form ammonia (NH3)
what does the balanced reaction look like?
 N2 (g) + 3H2 (g) = 2NH3 (g)
How would you write this to show 1 mole of product?
½ N2 (g) + 3/2 H2 (g)  NH3 (g) ▲Hf = -46 kJ.mol
What is more stable the reagents or the product?
The product!!!
E is given off – product has less H

11. Example reaction 2: Nitrogen gas reacts with oxygen gas to form nitrogen dioxide.
Bal. Rxn. w/1 mole product???
½ N2 (g) + O2 (g)  NO2 (g) ▲Hf = KJ/mol
What is more stable the reagents or the products?
The reagents
E goes into syst – prod. have greater H

12. Reaction #3: Aluminum solid reacts with oxygen gas to form aluminum oxide.
? rxn. w/1 mol. product
2Al (s) + 3/2 O2 (g)  Al2O3 (s) ▲H = KJ/mol
What is more stable?
The product

13. Summary Reaction Hf (kJ/mol) Stability of Product
½ N2(g) + 3/2 H2(g) → NH stable
½ N2(g) + O2 → NO unstable
2 Al + 3/2 O2 → Al2O very stable

14. 10-6 Enthalpy of a Reaction
H = ∑Hf (products) – ∑Hf (reactants)
Hc = enthalpy of combustion
~ defined for the combustion of 1 mole of a fuel
 CH4(g) O2(g) → CO2(g) H2O(l)
[ (-285.8)] – [ ] =
kJ/mol
Burning fuels is always exothermic!!!