1. Thermochemistry 1

2.  Thermodynamics = the study of heat and its transformations.  Thermochemistry = the part of thermodynamics that deals with changes in heat that take place during chemical reactions as in most chemical reactions, energy is absorbed or released. 2

3.  Internal Energy (E)- the combined kinetic and potential energy of all particles in a system.  The total internal energy of a system is not usually known. Energy changes (  E) can be measured. 3

4.  Potential Energy- Intermolecular forces of attraction or repulsion Intramolecular forces of attraction or repulsion  Kinetic Energy Rotational nature of molecules Vibrational nature of molecules 4 interactions between particles of matter down to the subatomic level (p, n, and e - ) but we are generally concerned about interactions at the molecular level.

5.  What is the system? The object or substance undergoing a change of physical state or reaction. What we are studying.  Surroundings are everything around the system that can exchange energy with the system. 5

6. 1. Open system can exchange mass and energy (usually in the form of heat) with the surroundings. 2. Closed system can exchange energy but not mass. 3. Isolated system – no transfer of energy or mass. 6

7.  State functions are path independent. They depend only on present state and are independent of history of the system.  The change in a state function does not depend on how the process is carried out. 7

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9. E total = mgh + ½ mv 2 OR Potential Energy + Kinetic Energy  Internal energy of the system = sum of kinetic and potential energies making up a substance down to the subatomic level. We often don’t know the total energy of a system.  Chemists aren’t so interested in total energy in a system- they’re more interested in the energy changes that happen during chemical reactions. 9

10.  The total energy in the universe is constant (Law of Conservation of Energy).  “You can’t get something for nothing”  There are essentially two ways to change the energy of a system- heat(q) and work(w).  E = q + w 10

11. Energy change in a system can be determined:  E = Energy can be transferred as heat + work  E = E f – E i = q + w Where work = force X distance OR pressure X volume change w = p  V Note: Expansion volume = -p  V. Work is being done by the system on the surroundings. 11

12.  Often times in chemical reactions, there isn’t a volume change or a force applied over a distance.  Therefore, for the chemist, heat changes are the primary concern when considering the energy changes that exist in a chemical reaction.  E = q (for many, not all, chemical reactions) 12

13.  Heat = thermal energy (motion of atoms)  Something is “hot” because it has thermal energy and its atoms are moving rapidly.  The more energy, the faster they move.  The total thermal energy of a substance is the sum of the individual thermal energies of the atoms and molecules making up a substance. 13

14.  Heat is a form of energy transfer.  Heat transfers from hot object to cooler object.  q>0 Heat is transferred from surroundings to system. ENDOTHERMIC  q<0 Heat is transferred from system to surroundings. EXOTHERMIC 14

15.  A measure of heat and relates to the average kinetic energy of atoms in a sample.  The higher the temperature, the greater the thermal motion. Thermal energy depends upon sample size. Heat ≠ Temperature  Thermometers display changes in thermal energy. 15

16.  By convention in an exothermic process, heat is transferred FROM the system TO the surroundings. Heat is a product. 2H 2 (g) + O 2 (g) → 2H 2 O (l) + energy  Combustion reactions are exothermic!  Chemical (potential) energy is converted to heat (kinetic energy). 16

17.  Heat is transferred TO the system FROM the surroundings. Heat is a reactant. Energy + 2HgO (s) → 2Hg(l) + O 2 (g)  Thermal energy (kinetic) is converted to chemical energy (potential). 17

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19.  Specific heat = the heat needed to raise the temperature of one gram of a substance by 1.0 0 C. Units J/g 0 C. Intensive property- not dependent on how much is present.  Heat Capacity = the measure of an overall effect of heat transfer on the temperature of an object. Units J/g. Extensive property (depends on mass). 19

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21.  The specific heat of water is one of the highest known specific heats.  That’s why temperature fluctuations are smaller near large bodies of water like Puget Sound than someplace like the desert in Arizona. 21

22.  The laboratory technique used to measure the heat released or absorbed during a chemical or physical change.  In a calorimeter, there is no volume change in the system, therefore,  E = q q = mc  T 22

23.  Calorimetry depends on the assumption that all the heat involved changes the contents of the calorimeter and its contents  no heat is gained or lost through the environment.  For many chemical reactions, a styrofoam cup is a convenient calorimeter because it has a very low heat capacity (for most purposes negligible) and it is a good insulator. 23

24.  A piece of iron with a mass of 72.4 grams is heated to 100. o C and plunged into 100.g of water that is initially at 10.0 o C. Calculate the final temperature that is reached assuming no heat loss to the surroundings. C water = 4.18 J/g o C C iron =0.449 J/g o C NOTE: the heat gained by the cooler body + the heat lost by the warmer body = 0 Answer: T final = 16.5 o C 24

25.  A 1.5886 g sample of glucose was ignited in a bomb calorimeter. The temperature increased by 3.682 o C. The heat capacity of the calorimeter was 3.52 kJ/ o C, and the calorimeter contained 1.000 kg of water. Find the total heat released and the molar heat of reaction. Remember: Heat energy goes into heating up the water and the calorimeter itself. Answer: = -28.35 kJ -3211 kJ/mol 25