1. Thermodynamics The study of the transformation of energy from one form into another. Chemical Thermodynamics The energy changes taking place during a chemical reaction.

2. Thermodynamics Energy is a measure of the capacity of a system to do work or produce heat. Potential Energy is the energy of position. Kinetic Energy is the energy of motion. Heat is a form of kinetic energy.

3. Thermodynamics Heat is the transfer of energy between objects due to a difference in temperature. In any spontaneous process energy flows from the object at higher temperature to the colder one. Work is the energy required to move something.

4. Thermodynamics A State function is independent of path and depends only on the initial and final positions, or states, of the system. Energy is a state function but work and heat are not. The system is the part of the universe we are studying. The surroundings is the rest of the universe

5. Thermodynamics Chemical energy is a form of potential energy. E.g. the reaction of aluminum and iron(III) oxide (the thermite process). 2Al (S) + Fe 2 O 3(S)  Al 2 O 3(S) + 2Fe (L) + heat This is an example of an exothermic reaction. Heat flows out of the reaction and q, the heat produced, is negative. In an endothermic process heat would flow into the system and q would be positive.

6. Thermodynamics Molar heat capacity at constant pressure, C P, is the amount of heat required to raise the temperature of 1 mole of a substance 1 K. Molar heat capacity at constant volume, C V, is the amount of heat required to raise the temperature of 1 mole of a substance 1 K. For an ideal monoatomic gas C V = 3R/2 and C P = C V +R & ΔE=nC V ΔT

7. Thermodynamics “q” is the amount of heat flowing into (+q) or out of (-q) a system during a change in the system. “w” is the amount of work done on (+w) or by (-w) the system during a change in the system. We will look only at PΔV work, the kind of work done when a gas expands (-w) or is compressed (+w). The change in energy of a system may be calculated using: ΔE = q + w

8. Thermodynamics ΔE = q + w (Between any two equilibrium states, the change in internal energy is equal to the sum of the heat transferred into the system and the work done by the system.) or: Energy is conserved or: You can’t win These are all statements of the First Law of Thermodynamics

9. Thermodynamics We will consider only PΔV work. (ΔP=0) If the system does work on the surroundings w is negative, but V Final > V Initial, or PΔV > 0 &w = -PΔV if ΔE > 0, the process will be endothermic. if ΔE < 0, the process will be exothermic & ΔE = q P +w

10. Thermodynamics We have a hot air balloon. We change the volume from 4.00*10 6 L to 4.50*10 6 L by adding 1.3810 8 Joules of energy as heat. If we do this at sea level, find ΔE for the process. BTW: 1 Joule = 1 kg-meter 2 /sec 2 or 0.239 calorie or 9.87*10 -3 L-atm

11. Thermodynamics Enthalpy Enthalpy is equal to the internal energy of a system plus the PV product of the system. The change in enthalpy of a system, ΔH can be found using the relationship: ΔH = ΔE + Δ(PV), and if ΔP = 0 (constant pressure), ΔH = ΔE + PΔV, therefore: ΔH = q P for a system at constant pressure and the only work is PΔV work.

12. Thermodynamics & Ideal Gases For an ideal monatomic gas C V = 3R/2 and C P = C V +R & ΔE=nC V ΔT (all the energy going into heating the gas, no work is being done because ΔV = 0) & ΔH = ΔE + Δ(PV) or ΔH = ΔE + nRΔT or ΔH = nC V ΔT + nRΔT = n(C V +R) ΔT or ΔH = nC P ΔT

13. Thermodynamics & Ideal Gases Note: C V was defined for ideal monatomic gases. For polyatomic gases C V is always larger because heat can be stored internally as vibrations and rotations, not just as linear motion. For example, gasH, Ne, Ar H 2, N 2 CO 2 C2H6C2H6 CVCV 12.520.629.044.6