Introduction to Thermodynamics and Kinetics

1. Introduction to Thermodynamics and Kinetics

2. Topics • Types of energy and units of energy • exothermic vs. endothermic reactions • heat capacity and specific heat • energy transfers in changes of state • calorimetry • enthalpy and entropy • Kinetics and catalysts • equilibrium and LeChatelier’s Principle

3. Energy is defined as the capacity to do work. • There are two catagories of energy: • Kinetic energy. • Potential energy • Kinetic Energy is the energy of motion, like climbing steps, or gears moving, etc. • Potential energy is energy of position, like a rock sitting at the edge of a cliff, or a tightly wound up spring. It can convert into Kinetic energy.

4. Chemical energy is a form of energy stored in the structure of a chemical substance.

5. Consider the potential energy of the rock up on the cliff. If one rock is up on a 1000 ft. cliff and another is up on a 10 ft. hill, the rock that is 1000ft. up will hit the ground with much more energy than the other if it falls.

6. Chemical energy has similar differences in that the structure of one substance may contain much more potential energy than another substance.

7. Thermochemistry is the study of heat transfers occurring in chemical and physical changes of substances. • Thermal energy is the energy associated with the random motion of atoms and molecules. • Heat is the transfer of thermal energy between two bodies that are different temperatures.

8. The Universe is made up of the system and the surroundings The system is the part of the universe that is of interest to us. This could be, for example, a beaker in which a chemical reaction is taking place. The surroundings is everything outside of the system.

9. Exothermic vs. Endothermic Processes • Any process that gives off heat to the surroundings is an Exothermic process. When bonds are made between particles, the process is exothermic.

10. When a process absorbs heat from the surroundings it is an Endothermic process. When ice melts it absorbs heat from the surroundings, thus it is an endothermic process.

11. The symbol DH is used to represent the change in heat into or out of the system. It is defined as the change in enthalpy.

12. The enthalpy of reaction is the difference between the enthalpies of the products and reactants. DH(rxn) = DH(products) - DH(reactants)

13. The diagram illustrates energy givenoff when bonds are made between H2 and O2 , (a). Energy is absorbed when bonds are broken in HgO , (b).

14. The units of energy are Kilojoules and Calories. Enthalpy of reaction will be expressed as Kilojoules (Kj). A conversion factor to go between these two units is 1 cal = 4.184 J

15. It is important to note the following when calculating the enthalpy of reaction: • The enthalpy of a substance in its standard state is equal to zero. • Enthalpy is dependent on the quantity of the substance therefore the enthalpy of a substance must be multiplied by its coefficient in the chemical reaction.

16. Enthalpy values of compounds are found in tables of thermodynamic data. Ex. 1 Determine the enthalpy for the following chemical reaction: CO2 (g) + 2 H2O (l) -----> 2 O2 (g) + CH4 (g) continued...

17. Use table 7.2 in your text to find the individual enthalpy values. CO2 (g) = -393.509, H2O (l) = -285.83 CH4 (g) = -74.81 , O2(g)= 0 CO2 + 2H2O ----> 2 O2 + CH4 DH = DH(products) – DH(reactants) DH = (0 x -74.81)-(-393.509 + 2(- 285.83)) = -74.81 + 965.17 = 890.36 Kj

18. Specific heat and heat Capacity Heat capacity is the amount of heat required to raise the temperature of a given quantity of a substance by 1o C. Specific Heat is the amount of heat required to raise the temperature of 1 gram of a substance by 1o C.

19. These can be thought of as a substance’s ability to absorb heat and to store heat. For example, a metal does not require much energy to heat it up and it does not hold the heat for a long time. A liquid like water (which contains strong intermolecular forces, hydrogen bonds), requires much more energy to heat it up and it does hold the heat for a longer period of time.

20. The metal has a low heat capacity and the water has a high heat capacity. The specific heat capacity of water is an important and easy number to remember ; 1 cal / g oC , (or 4.184 J / goC

21. Calorimetry is a technique used in the lab to measure the enthalpy of a reaction. One apparatus used is the Bomb Calorimeter. It is a Heavy walled, steel container which has a known specific heat.

22. The heat of the reaction (q) is equal to the negative of the heat absorbed by the (bomb plus the water) surrounding the bomb.

23. q is used to represent the quantity and direction of heat transferred, using a calorimeter.

24. Necessary Equations: q = (specific heat)(mass)(change in Temp.) q = (J/gK)(g)(DT) q(rxn) = - (q of the water + q of the bomb)

25. Subliminal message..... Wake up !!!

26. Ex. 2 Calorimetry A 466g sample of water is heated from 8.50oC to 74.60oC. Calculate the amount of heat absorbed by the water. q = (sp. heat)(mass)(DT) q = (4.184 J/goC)(466g)(74.60-8.50oC) q = 1.29 x 105 J = 129 KJ

27. Ex. 3 1.435 g of Naphthalene (molar mass =128.2) was burned in a bomb calorimeter. The temp. rose from 20.17oC to 25.84 oC. The mass of the water surrounding the calorimeter was 2000.g and the heat capacity of the bomb was 1.80 KJ/oC. Calculate the heat of combustion of Naphthalene.

28. q(rxn) = -(q water + q bomb) q(water) = (2000g)(4.184 J/goC)(5.67oC) = 4.74 x 104 q(bomb) = (1.80 x 103J / oC)(5.67oC) = 1.02 x 104J q(rxn) = -(4.74 x 104 J + 1.02 x 104 J) = -5.76 x 104 J