Thermodynamics - The First Law

1. Thermodynamics- The First Law Atkins Chapter 7

2. Introduction • The laws of thermodynamics govern chemistry and life. • They explain why reactions take place and let us predict how much heat reactions release and how much work they can do. • Thermodynamics plays a role in every part of our lives. • Key concerts of this part – heat and work, internal energy, enthalpy, heat capacity

3. Systems Systems and Surroundings System: part of the universe we are interested in. Surroundings: the rest of the universe.

4. Three typical types of Systems Open Closed Isolated

5. Work and energy • Work (w) is the transfer of energy to a system by a process • Work = opposing force X distance moved • Total energy stored in a system is called internal energy U of the system, cannot measure absolute energy, only measure relative energy ∆U • The work done to a system is converted to the internal energy of the system if no other source • ∆U = w

6. Expansion and nonexpansion work • Two types of work: • Expansion – involves change of volume • Nonexpansion – no change of volume

7. Expansion work The minus sign meansthat the internalenergy of the system decreases because some energy is lost as work. The work is proportional to the external pressure For a given external pressure, more work is done when the systemexpands by a large amount

8. Reversible process In thermodynamics, a reversible process is one that can be reversed by an infinitely small change in a variable (an “infinitesimal” change). All processes are irreversible, reversible is an approximation Reversible is a process in which each step is infinitely close to equilibrium, there is no friction, there is no heat dispersion (entropy lost)

9. Reversible process For example, consider a cylinder without any friction If the external pressure exactly matches the pressure of the gas in the cylinder, then the piston moves in neither direction; If the external pressure is increased infinitesimally, the piston moves in. If the external pressure is reduced infinitesimally, the piston moves out. This is a quasi-equilibrium process, close to reversible

10. Work by reversible process The simplest kind of reversible change to consider is the reversible, isothermal (constant temperature) expansion of an ideal gas.

11. Reversible process According to the first law: ∆U = w + q, because ∆U=0 (ideal gas, isothermal), then we have: –w = q, which means work by the system equals the heat transferred in. In a reversible process, there is no friction, no heat lost, therefore: The work a system can do is greatest in a reversible process. To prove this statement, we need the second law (Carnort theorem)

12. Reversible and irreversible processes A piston confines 0.100 mol Ar(g) in a volume of 1.00 L at 25°C. Two experiments are performed. (a) The gas is allowed to expand through an additional 1.00 L against a constant pressure of 1.00 atm. (b) The gas is allowed to expand reversibly and isothermally to the same final volume. Which process does more work?

13. Heat In thermodynamics, heat is the energy transferred as a result of a temperature difference. Heat is not “thermal energy” which is a sum of potential and kinetic energy arising from the chaotic thermal motion When energy is transferred as heat and no other processes occur: When energy enters a system as heat, q is positive; when energy leaves a system as heat, q is negative.

14. The measurement of heat If system does not lose energy as work, energy taken as heat will increase temperature Heat capacity measures how much energy as heat is transferred to increase one degree of temperature.

16. The first law ∆U = q + w is a complete statement of how change of internal energy can be achieved Energy can be transferred either as work or as heat. not possible to make a “perpetual motion machine” The difference between these two ways of transfer is how atoms are moved. The internal energy of an isolated system is constant

17. State Function Property that only depends on state Internal energies, pressure, temperature, volume, density, enthalpy, entropy, etc. Work and Heat are not state functions

18. Ideal gas expansion • For isotherm expansion of ideal gas: ∆U=0 • Because • No interaction – potential energy does not change • Same temperature – kinetic energy is the same

19. Ideal gas expansion For ideal gas ∆U=0, reversible and irreversible processes

20. The origin of internal energy • Internal energy is stored as molecular kinetic and potential energy. • The equipartition theorem says each degree of freedom contributes kT/2 • For ideal gas, only consider intra-molecular contributions • The vibrational contribution is insignificant at ordinary temperatures.

21. Heat capacity of molecular origin • Heat capacity is due to energy stored in molecules • Each degree of freedom takes R/2 when it is in full strength • Different motions become important at different temperatures • Cv for I2 is shown in the figure

22. Enthalpy At constant volume, if no non-expansion work, then ∆U = q Most experiments are done at constant pressure. Need to define another statefunciton enthalpy H = U + PV If there is no non-expansion work ∆H = q

23. Proof

24. Heat capacities C = q / ∆T At constant pressure ∆H = q Cp = ∆H / ∆T For ideal gas, Cp = Cv + nR Without vibration (at low temperature)

25. Enthalpy of physical changes - water How energies are transferred as heat?

26. Enthalpy of chemical changes Calorimetry shows that burning 1.00 mol CH4(g) produces 890. kJ of heat at298 K and 1 bar. How are energies transferred?

27. Enthalpy of ideal gas expansion Consider gas expansion isothermally to fill an evacuated flask. If it is ideal gas, how is energy changed? If it is not ideal gas, what happens?

28. Homework Atkins book: 7.6, 7.10, 7.20, 7.26, 7.32, 7.34, 7.42,7.48, 7.56, 7.70, 7.84, 7.90, 7.98, 7.106, 7.108, 7.112