1. Chapter 6 – Thermochemistry

2. A.The chemistry related to heat change in chemical reactions 1.Energy – ability to do work or produce heat work = force x distance w = F x d a. Kinetic – motion b. Potential – stored c. Radiant – travels through space

3. What is Energy? the resources for producing usable power that which is needed to oppose natural attractions the capacity to do work or to produce heat

4. Forms of Energy Potential – Energy due to position or composition Kinetic – Energy due to the motion of an object E K = ½ mv 2

5. Energy is the capacity to do work Thermal energy is the energy associated with the random motion of atoms and molecules Chemical energy is the energy stored within the bonds of chemical substances Nuclear energy is the energy stored within the collection of neutrons and protons in the atom Electrical energy is the energy associated with the flow of electrons Potential energy is the energy available by virtue of an object’s position

6. 2. Units – Joule – J E K = ½ mv 2 = ½ (2kg) (1 m/s) 2 = 1 kg m 2 s2s2 = 1 J(joule) 4.184 J = 1 calories 1 CAL = 1000 cal = 1 kcal 4.184 kJ = 1 CAL = 1 kcal mass = 2 kg height = 1 meter

7. 3.System vs. Surroundings System : That part of the Universe whose change we are going to measure. Surroundings : Every thing else that is relevant to the change is defined as the “surroundings”.

8. Heat energy transferred between two objects as a result of the temperature difference between them.

9. Temperature A measure of kinetic energy

10. Heat is the transfer of thermal energy between two bodies that are at different temperatures. Energy Changes in Chemical Reactions Temperature is a measure of the thermal energy. Temperature = Thermal Energy 90 0 C 40 0 C greater thermal energy

11. B. 1st Law of Thermodynamics The energy of the universe is constant. Law of Conservation of Energy 1. Internal Energy – Energy of a system is the sum of all the kinetic and potential parts  E = E final  E initial   E if energy leaves system +  E if energy enters system

13. 2. State Function A function or property whose value depends only on the present state (condition) of the system, not on the path used to arrive at that condition. Note the E of a system doesn’t depend on how system got there

14. Thermodynamics State functions are properties that are determined by the state of the system, regardless of how that condition was achieved. Potential energy of hiker 1 and hiker 2 is the same even though they took different paths. energy, pressure, volume, temperature

15. E is a state function, q and w are not.

16. 3. Heat and Work  E = q + w heat gain or loss work done = -P  V REMEMBER: E in is + (endo) E out is – (exo)

17. Example #1 : Heat and Work A system performs 50 kJ of work on its surroundings and absorbs 20 kJ of heat from its surroundings. What is the change in internal energy of the system?  E = q + w = 20 kJ + (-50 kJ) = -30 kJ

18. Determining Energy Change in a System Problem: In the internal combustion engine, the heat produced by the combustion of the fuel causes the carbon dioxide and water that is produced during the combustion to expand, pushing the pistons. Excess heat is removed by the cooling system. Determine the change in energy ( E) in J, kJ, and kcal if the expanding gases do 515 J of work on the pistons, and and the system loses 407 J of heat to the cooling system. q = - 407 J w = -515 J  E = q + w = - 407 J + ( - 515 J) = - 922 J

19. Thermodynamics  E = q + w  E is the change in internal energy of a system q is the heat exchange between the system and the surroundings w is the work done on (or by) the system w = -P  V when a gas expands against a constant external pressure