Energy and Chemical Reactions

1. Chapter 6 Energy and Chemical Reactions Germain Henri Hess 1802-1850. Hess’s Law. James P. Joule 1818-1889. Discovered mechanical equivalent of heat, which led to the First Law of Thermodynamics.

2. The Nature of Energy Kinetic and Potential Energy Kinetic energy is the energy of motion. Potential energy is energy which is available, such as an object that can fall a distance, potentially releasing kinetic energy when it hits the ground.

3. The Nature of Energy Energy Units We sometimes use the calorie instead of the joule: 1 cal = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal SI Unit for energy is the joule, J: For example, energy of motion (kinetic energy): (in J)

4. The Nature of Energy All objects and substances have an internal energy, E (this is a form of potential energy, and in this course is commonly called “chemical energy”). In chemistry, we’re interested in the internal energies of reactants and products and how it changes (ΔE) in going from: reactantsproducts We usually can only measure ΔE, not E “Δ” means change in

5. First Law of Thermodynamics Total energy is conserved Energy cannot be created or destroyed. In chemistry, we are concerned with heat and chemical energy. The first law of thermodynamics tells us that the sum of heat and chemical energy must be constant (at constant pressure). Heat at constant pressure is called enthalpy.

6. Enthalpy Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure. Can only measure the change in enthalpy: DH = Hfinal - Hinitial = qP We will use enthalpy (H) and energy (E) interchangeably

7. Endothermic and Exothermic Processes Endothermic: system absorbs heat from surroundings. (ΔH = +) Exothermic: system transfers heat to the surroundings. (ΔH = -) An endothermic reaction feels cold. An exothermic reaction feels hot.

8. Enthalpies of Reaction For a reaction DHrxn = H(products) - H (reactants) Enthalpy is an extensive property (magnitude DH is directly proportional to amount): CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) DH = -1604 kJ When we reverse a reaction, we change the sign of DH: CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ Change in enthalpy depends on state (gas, liquid, solid): H2O(g)  H2O(l) DH = -44 kJ

9. Calorimetry Heat Capacity and Specific Heat Calorimetry = measurement of heat flow. Calorimeter = apparatus that measures heat flow. Heat capacity = the amount of energy required to raise the temperature of a substance (by 1o). Molar heat capacity = heat capacity of 1 mol of a substance (energy required to raise temp. of onemole of a substance (by 1o) Specific heat = specific heat capacity = heat capacity of 1 g (energy required to raise temp of onegram by 1o. Heat absorbed by system, q, is given by: q = (specific heat)  (grams of substance) T (Be careful of the sign of q)

10. Calorimetry Some specific heats (in J/g-K)… NOTE the very high specific heat for liquid water. N2(g): 1.04 Al(s): 0.90 Fe(s): 0.45 CH4(g): 2.20 CaCO3(s): 0.82 H2O(l): 4.18 This has a moderating influence on the weather and the temperature of living organisms.

11. Calorimetry Constant-Pressure Calorimetry Atmospheric pressure is constant! DH = qP Heat emitted by reaction (qrxn) is absorbed by solution (qsoln) So, qrxn = -qsoln qrxn = -qsoln = -(specific heat of soln)  (grams of solution) DT.

12. Calorimetry 9.55g sample of solid NaOH dissolves in 100.0 g water in coffee-cup calorimeter. Temp rises from 23.6 to 47.4oC. Calculate ΔH (kJ/mol) for the soln process: NaOH(s) Na+ (aq) + OH- (aq) Assume specific heat of soln is same as that of pure water (which is 4.18 J/g-K).

13. Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry) Used to study combustion processes qrxn = -Ccal ΔT where Ccal is the heat capacity of the calorimeter (to be determined experimentally)

14. Hess’ Law • Hess’ law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step. Note that: H1 = H2 + H3

15. Hess’ Law • A consequence of Hess’ Law is that reactions can be added just like algebraic equations. • For example: CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ 2H2O(g)  2H2O(l) H = -88 kJ CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ

16. Enthalpies of Formation • If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof . This is also called the heat of formation. Standard conditions (standard state): 1 atm and 25 oC (298 K).

17. Enthalpies of Formation • Standard enthalpy of formation of the most stable form of an element is zero. For example, it is zero for C(s), H2(g), Pb(s), O2(g), etc

18. Enthalpies of Formation Using Enthalpies of Formation to Calculate Enthalpies of Reaction. We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation Example: Combustion of 1 mol of propane gas, C3H8 (g) Can construct alternate path of three steps shown as 1, 2 and 3 Hrxn =H1 + H2 +H3

19. Enthalpies of Formation Using Enthalpies of Formation to Calculate Enthalpies of Reaction For a reaction: (a) 2 SO2(g) + O2(g) 2 SO3 (g) -- what is ΔH? (c) 4 FeO (s) + O2 (g) 2 Fe2O3 (s) -- what is ΔH?

20. Foods and Fuels 1 Nutritional Calorie = kcal Foods

21. Foods and Fuels Fuels U.S.: 1.0 x 106 kJ of fuel per day. Most from petroleum and natural gas. Remainder from coal, nuclear, and hydroelectric. Fossil fuels are not renewable.