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THERMOCHEMISTRY HEAT AND CHEMICAL CHANGE

Explore energy transformations, endothermic and exothermic processes, calorimetry, Hess’s Law, and heat change in state transitions.

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THERMOCHEMISTRY HEAT AND CHEMICAL CHANGE

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  1. THERMOCHEMISTRYHEAT AND CHEMICAL CHANGE CHAPTER 11

  2. 11.1 ENERGY TRANSFORMATIONS • Thermochemistry- • Heat gain or loss that occur during chemical changes • Physical change (boiling, evaporation, condensing) • Chemical reaction (combustion, combination)

  3. Energy- • Ability to do work or supply heatEnergy can not be observed, only detected because of its effects

  4. Chemical potential energy-energy stored in chemical bonds • During reactions as bonds are broken and rearranged, energy is released or absorbed

  5. H-heat energy that is transferred from one object to another (book uses q, not H) • ∆H is the change in heat energy • Heat always flows from the warmer object to the cooler one

  6. ENDOTHERMIC AND EXOTHERMIC PROCESS • To study heat, must define the system and surroundings • Together the system and the surroundings constitute the universe • Law of Conservation of Energy • Energy lost equals energy gained • Hence energy is neither created nor destroyed

  7. EXOTHERMIC SURROUNDINGS • Δ H = - kJ/mol • Initially-steam • Energyreleased by steam into surroundings • System (steam)losesheat energy • Finally-water • Surroundings arewarmer SYSTEM

  8. ENDOTHERMIC SURROUNDINGS • Initially-liquid water • Energyaddedfrom surroundings to water • System (water) gains heat energy to vaporize • Finally-vapor • Surroundings arecooler • Δ H =+kJ/mol SYSTEM SYSTEM

  9. H2O vapor EXOTHERMIC Δ H = - kJ/mol ENDOTHERMIC Δ H = + kJ/mol H20 liquid

  10. Heat Capacity & Specific Heat • Calorie:the quantity of heat needed to raise the temperature of 1 g of pure water 1 degree C. 1 Calorie = 1 kilocalorie = 1000 calories 1 J = 0.2390 cal 4.184 J = 1 cal

  11. Specific Heat Capacity • The amount of heat it takes to raise the temperature of 1 g of the substance 1 degree C. • Water vs. Iron • Which one has a higher specific heat capacity?

  12. Enthalpy •  H = m x C x T m = mass C = specific heat capacity (table) T = T(final) – T (initial) 

  13. Thermochemical Equations • Includes heat change in the equation CaO (s) + H2O (l) Ca(OH)2 (s) + 65.2 kJ 2NaHCO3 (s) + 129 kJ Na2CO3 (s) + H2O (g) + CO2 (g) Which one is endothermic? Exothermic?

  14. Is the following reaction endothermic or exothermic? CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l)  H = -890 kJ

  15. CH. 11.3: HEAT IN CHANGES OF STATE • Heat of Fusion (Hfus): solid  liquid 6.01 kJ/mol or 334 J/g • Heat of Solidification (Hsolid): liquid  solid -6.01 kJ/mol or –334 J/g

  16. How many grams of ice at 0C and 101.3 kPa could be melted by the addition of 2.25 kJ of heat?

  17. Heat of Vaporization (Hvap): liquid  gas 40.7 kJ/mol or 2.26 kJ/g • Heat of Condensation (Hcond): gas  liquid -40.7 kJ/mol or –2.26 kJ/g

  18. How much heat (in kJ) is absorbed when 24.8 g H2O (l) at 100C is converted to steam at 100C?

  19. ·Heat of Solution (Hsoln): heat change caused by dissolution • How much heat (in kJ) is released when 2.500 mol NaOH (s) is dissolved in water? NaOH (s)  Na+ (aq) + OH- (aq) Hsoln = -445.1 kJ/mol

  20. CH. 11.4: CALCULATING HEAT CHANGES Hess’s Law: If you add 2 or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction.

  21. If you reverse an equation, then reverse the sign of H. • If you multiply an equation by a coefficient, then multiply the value of H.

  22. C (diamond) C (graphite) C (graphite) + O2 (g)  CO2 (g) H = -393.5 kJ C (diamond) + O2 (g)  CO2 (g) H = -395.4 kJ

  23. Standard Heats of Formation oHf of a free element in its standard state is 0. (example: diatomic molecules) oH = Hf (products) - Hf (reactants)

  24. What is the standard heat of reaction (H) for the following reaction? 2CO (g) + O2 (g)  2CO2 (g)

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