1. 1.2.2 Heat of Formation

2.  Standard Heat of Formation Δ H o f  the amount of energy gained or lost when 1 mole of the substance is formed from its elements under standard conditions (25°C, 1 atm = 101.3 kPa)

3.  Ex. the formation reaction for liquid water is described by the following equation:  H 2 (g) + ½O 2 (g) → H 2 O(l) + 285.8 kJ  The standard heat of formation is: 285.8 kJ. Since the reaction is exothermic: Δ H o f -285.8 kJ.

4.  A heat of formation is a type of reaction where one mole of the compound forms from its elements  The heat of formation for pure elements, such as H 2 (g), O 2 (g), Al(s), etc. is 0 kJ/mole. You'll find it useful to remember this.

5.  Look over the Heat of formation table  See values in outline

6.  Keep the following points in mind: 1. Balance the equation so that one mole of the compound is produced. 2. Remember the diatomic (7) molecules and write them correctly (H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 ). 3. The reactants must be elements, not polyatomic ions. Examples of polyatomic ions are hydroxide, OH -, carbonate, CO 3 2-, and ammonium, NH 4 +.

7.  H2(g) + ½O2(g) → H2O(l) + 285 kJ.  If 285.8 kJ of energy are released during the formation of one mole of H2O(l), how much energy do you imagine would be released if two moles of water were produced?

8.  If you predicted 571.6 kJ of energy you're right! Our formation reaction tells us that 285.8 kJ of energy are released for every one mole of H 2 O. This can be written mathematically as: 285.8 kJ/mol H 2 O×2 mol H 2 O =571.6 kJ  Our new equation looks like this:  2 H2(g) + O2(g) → 2 H2O(l) + 571.6 kJ

9.  1. Write heats of formation reactions for each of the following compounds. Be sure to include the energy term with the equation, either as part of the equation or separately as Δ H. You will need to refer to a Table of Thermochemical Data.  CO 2 (g),  CuCl 2 (g),  CuCl (g),  N 2 H 4 (l),  NH 4 Cl (s).

10.  C (s) + O 2 (g) → CO 2 (g) + 393.5 kJ orC (s) + O 2 (g) → CO 2 (g) Δ H = -393.5 kJ  Cu (s) + Cl 2 (g) → CuCl 2 (s) + 220.1 kJ or Cu (s) + Cl 2 (g) → CuCl 2 (s) Δ H = -220.1 kJ  Cu (s) + ½ Cl 2 (g) → CuCl 2 (s) + 137.2kJ or Cu (s) + ½ Cl 2 (g) → CuCl 2 (s) Δ H = -137.2 kJ  N 2 (g) + 2H 2 (g) + 50.6 kJ → N 2 H 4 (l) or N 2 (g) + 2H 2 (g) → N 2 H 4 (l) Δ H = +50.6 kJ  ½N 2 (g) + 2H 2 (g) + ½Cl 2 (g) → NH 4 Cl (s)+ 314.4 kJ or ½N 2 (g) + 2H 2 (g) + ½Cl 2 (g) → NH 4 Cl (s) Δ H = -314.4 kJ

11. The standard heat of formation, Δ H o f, for sulfur dioxide (SO 2 ) is -297 kJ/mol. How many kJ of energy are given off when 25.0 g of SO 2 (g) is produced from its elements?

12.  Step 1: Calculate moles SO2 a. Find the molar mass of SO2 = 32.1 + 2(16.0) = 64.1 g/mol b. moles SO2 = Step 2: Determine kJ for 0.390 mol We know from the question that 297 kJ of energy is released for 1 mole of SO 2 Determine how much energy will be released for 0.390 mol of SO 2 :

13. The heat of reaction for the combustion of 1 mol of ethyl alcohol is -9.50 × 10 2 kJ: C 2 H 5 OH (l) + 3 O 2 (g) → 2 CO 2 (g) + 3 H 2 O (l) + 9.5 × 10 2 kJ How much heat is produced when 11.5 g of alcohol is burned?

15.  First determine how many moles of ethyl alcohol are being combusted. You need to begin by finding the molar mass of C 2 H 5 OH, which is 46.0 g/mol.  moles C 2 H 5 OH =  From our balanced equation we see that 9.50 × 10 2 kJ of energy are released for every 1 mole of C 2 H 5 OH. Now we determine how much energy will be released for 0.250 mol:

16. Δ H for the complete combustion of 1 mol of propane is -2.22 × 10 3 kJ: C 3 H 8 (g) + 5 O 2 (g) → 3 CO 2 (g) + 4 H 2 O (l) Calculate the heat of reaction for the combustion of 33.0 g of propane.

18.  4. determine moles of propane actually used. The molar mass of C 3 H 8 is: moles C 3 H 8 = From our balanced equation we see that 2.22 × 10 3 kJ of energy are released for every 1 mole of C 3 H 8. Now determine how much energy will be released for 0.750 mol:

19.  Assignment- practice problems