1. Thermochemistry! AP Chapter 5

2. Temperature vs. Heat Temperature is the average kinetic energy of the particles in a substance. Heat is the energy that is transferred from one object to another. Heat always flows from the hotter object to the colder object.

3. Energy! Energy is the ability to do work. Kinetic Energy - the energy of motion Potential Energy – the energy that an object has as a result of its composition or its position with respect to another object.

4. Units of Energy 1 Joule = 1 kg m 2 /s 2 (1 kJ is 1000J) Used to calculate the energies associated with chemical reactions. Calorie – Amount of energy required to raise the temperature of 1 gram of substance 1 ° C. (This is specific heat!) 1 calorie will raise the temperature of 1 g of H2O from 14.5 ° C to 15.5 ° C. 1 calorie is equal to 4.184 Joules (exactly!)

5. Systems and Surroundings System – the portion used in a study. –It can be an open system or a closed system. Open system – matter and energy can interact with the surroundings. Closed system – the matter cannot interact with the surroundings.

6. First Law of Thermodynamics Energy Is Conserved!

7. Internal Energy Internal Energy is the sum of all the kinetic and potential energies of all its components. ΔE = E final - E initial

8. ΔEΔE A positive value for ΔE is when E final > E initial If energy has been absorbed from its surroundings, it is endothermic. If energy is given off to the surroundings, it is exothermic.

9. Initial state refers to the reactants, while final state refers to the products.

10. Endothermic reactionExothermic reaction

11. A system composed of H 2 ( g ) and O 2 ( g ) has greater internal energy than a system composed of H 2 O ( l ). Gases have greater kinetic energy and must lose some of that energy to change states back to the liquid state.

12. Internal energy is a function of state.

13. a)If a battery is shorted out and loses energy to the environment only as heat, no work is done. b)If a battery is discharged and loses energy as work (to make the fan run) it also loses heat energy. c)The value of ∆E is the same.

14. Enthalpy The change in enthalpy for a reaction (∆H) is the overall measure of energy that is absorbed to break bonds and the energy that is released when new bonds form. A reaction is said to be spontaneous if it occurs without being driven by an outside force. (driving forces are enthalpy & entropy) ∆H = ΣH (products) - ΣH (reactants)

15. In an endothermic system where it absorbs heat, ∆H will be positive ( ∆H > 0). In an exothermic system, where heat is given off, ∆H will be negative ( ∆H < 0).

16. Enthalpy Diagrams Enthalpy is an extensive property – it depends on how much you have. If 1mol of CH 4 and 2 mol O 2 yield -890 kJ, then 2 mol CH 4 and 4 mol O 2 would yield double that. The enthalpy change for a reaction is equal in magnitude, but opposite sign, for a reverse reaction.

18. Calorimetry This is a measure of the amount of energy that is needed or lost when a certain mass of a substance changes temperature. q = mC ∆T q is the amount of energy (J) m is the mass of the substance (g) C is the specific heat capacity of the substance ∆T is the change in temperature

19. Calorimeters Calorimeters are devices that measure the transfer of heat from one object to another.

20. Heat of Formation ( ∆H ° f ) The heat change that occurs when one mole of a compound is formed from its elements at 1 atm pressure. Generally, the standard enthalpy of formation for any element in its most stable form is 0. (i.e. O 2 gas would have a standard enthalpy of 0.) Remember Appendix C!

21. Standard Enthalpy Changes The standard enthalpy change can be calculated from the standard enthalpies of formation of the reactants and products in the reaction (see Appendix C for values.) ∆H° rxn = Σ n∆H° f (products) - Σm ∆H° f (reactants) The n and m refer to the molar coefficients in the chemical equation.

22. Also refer to Appendix C!

23. Hess’s Law If you can break a chemical reaction into several steps, add up all of the ∆H’s for each step to get the overall ∆H for the reaction.

25. Entropy Entropy is a measure of randomness or disorder of a system. The greater the disorder, the greater the entropy. In terms of entropy, gases>liquids>solids. When pure substance dissolves in a liquid, its entropy increases. When gas molecules escape a solvent, entropy increases. Entropy increases with molecular complexity. Reactions that increase the number of moles of particles often increase the entropy of the system.

26. Predict! Na + (aq) + Cl - (aq) → NaCl (s) ∆S is negative NH 4 Cl (s) → NH 3 (g) + HCl (g) ∆S is positive