1 / 43

Chemistry 434

Chemistry 434. A Brief Review of Thermodynamics. Internal Energy and the First Law. The infinitesimal change in the internal energy . For a general process. The First Law of Thermodynamics. The Constant Volume Heat Capacity. Define the constant volume heat capacity, C V. Enthalpy.

nika
Télécharger la présentation

Chemistry 434

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemistry 434 A Brief Review of Thermodynamics

  2. Internal Energy and the First Law • The infinitesimal change in the internal energy • For a general process The First Law of Thermodynamics

  3. The Constant Volume Heat Capacity • Define the constant volume heat capacity, CV

  4. Enthalpy • We define the enthalpy of the system, H

  5. The Constant Pressure Heat Capacity • Define the constant pressure heat capacity, CP

  6. Thermodynamic Definition • Spontaneous Process – the process occurs without outside work being done on the system.

  7. Mathematical Definition of Entropy • The entropy of the system is defined as follows

  8. The Fundamental Equation of Thermodynamics • Combine the first law of thermodynamics with the definition of entropy.

  9. The Temperature dependence of the Entropy • Under isochoric conditions, the entropy dependence on temperature is related to CV

  10. Entropy changes Under Constant Volume Conditions • For a system undergoing an isochoric temperature change • For a macroscopic system

  11. The Temperature dependence of the Entropy • Under isobaric conditions, the entropy dependence on temperature is related to CP

  12. Entropy changes Under Constant Pressure Conditions • For a system undergoing an isobaric temperature change • For a macroscopic system

  13. The Second Law of Thermodynamics • The second law of thermodynamics concerns itself with the entropy of the universe (univS). • univS unchanged in a reversible process • univS always increases for an irreversible process

  14. The Third Law of Thermodynamics • The Third Law - the entropy of any perfect crystal is 0 J /(K mole) at 0 K (absolute 0!) • Due to the Third Law, we are able to calculate absolute entropy values.

  15. Combining the First and Second Laws • From the first law

  16. Pressure Volume and Other Types of Work • Many types of work can be done on or by chemical systems. • Electrical work. • Surface expansion. • Stress-strain work. dw=-Pext dV+dwa where dwa includes all other types of work

  17. The General Condition of Equilibrium and Spontaneity • For a general system

  18. Isothermal Processes • For a systems where the temperature is constant and equal to Tsurr

  19. The Helmholtz Energy • Define the Helmholtz energy A A(T,V) =U – TS • Note that for an isothermal process dA  dw A  w • For an isochoric, isothermal process A  0

  20. The Properties of A • The Helmholtz energy is a function of the temperature and volume

  21. Isothermal Volume Changes • For an ideal gas undergoing an isothermal volume change

  22. Isothermal Processes at Constant Pressure • For an isothermal, isobaric transformation

  23. The Gibbs Energy • Define the Gibbs energy G G(T,P) =U – TS+PV • Note that for an isothermal process dG  dwa G wa • For an isothermal, isobaric process G  0

  24. The Properties of G • The Gibbs energy is a function of temperature and pressure

  25. Isothermal Pressure Changes • For an ideal gas undergoing an isothermal pressure change

  26. Temperature Dependence of A • Under isochoric conditions

  27. Gibbs Energy Changes As a Function of Temperature • The Gibbs energy changes can be calculated at various temperatures

  28. The Chemical Potential • Define the chemical potential  = G/n

  29. Gibbs Energy and Spontaneity sysG < 0 - spontaneous process sysG > 0 - non-spontaneous process (note that this process would be spontaneous in the reverse direction) sysG = 0 - system is in equilibrium

  30. Applications of the Gibbs Energy • The Gibbs energy is used to determine the spontaneous direction of a process. • Two contributions to the Gibbs energy change (G) • Entropy (S) • Enthalpy (H) G = H - TS

  31. Thermodynamics of Ions in Solutions • Electrolyte solutions – deviations from ideal behaviour occur at molalities as low as 0.01 mole/kg. • How do we obtain thermodynamic properties of ionic species in solution?

  32. For the H+(aq) ion, we define • fH = 0 kJ/mole at all temperatures • S = 0 J/(K mole) at all temperatures • fG = 0 kJ/mole at all temperatures

  33. Activities in Electrolyte Solutions • For the following discussion • Solvent “s” • Cation “+” • Anion “=“ • Consider 1 mole of an electrolyte dissociating into + cations and - anions G = nss + n  = nss + n+ + +n- - • Note – since  = + + -   = + + +- -

  34. The Mean Ionic Chemical Potential • We define =  /  • We now proceed to define the activities  =  + RT ln a + = + + RT ln a+ - = - + RT ln a-  =  + RT ln a

  35. The Relationship Between a and a • Since =  /   =  + RT ln a =  ( + RT ln a) Since =  /  • This gives us the relationship between the electrolyte activity and the mean activity (a)= a

  36. The Relationship Between a , a- and a+ • We note that  = + + +- - and =  /  • This gives us the following relationship ( + RT ln a) = n+ (+ + RT ln a+) + - ( - + RT ln a-) • Since  = +++ -- (a) = (a+)+ (a-)-

  37. Activities in Electrolyte Solutions • The activities of various components in an electrolyte solution are defined as follows a+ = + m+ a- = - m- a+ = + m+ • As with the activities () = (+)+ (-)- (m) = (m+)+ (m-)-

  38. The Chemical Potential Expression • This can be factored into two parts Deviations from ideal behaviour The ideal part

  39. Activity Coefficients As a Function of Molality Data obtained from Glasstone et al., Introduction to Electrochemistry, Van Nostrand (1942). CRC Handbook of Chemistry and Physics, 63rd ed.; R.C. Weast Ed.; CRC Press, Boca Raton, Fl (1982). CaCl2 HCl LaCl3 KCl H2SO4

  40. Estimates of Activity Coefficients in Electrolyte Solutions • The are a number of theories that have been proposed to allow the theoretical estimation of the mean activity coefficients of an electrolyte. • Each has a limited range of applicability.

  41. The Debye Hűckel Limiting Law • This is valid in the up to a concentration of 0.010 molal! Z+ = charge of cation; z- = charge of anion

  42. Debye Hűckel Extended Law • This equation can reliably estimate the activity coefficients up to a concentration of 0.10 mole/kg. B = 1.00 (kg/mole)1/2

  43. The Davies Equation • This equation can reliably estimate the activity coefficients up to a concentration of 1.00 mole/kg. k = 0.30 (kg/mole)

More Related