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Second Year Chemistry

Second Year Chemistry. 1 st semester: Organic 1 st semester: Physical (2005-2006) December exams 2 nd : Analytical & Environmental 2 nd : Inorganic Summer exams Physical: 3 lecturers Þ ~ 8 topics Dónal Leech: four topics Thermodynamics Gases , Laws, Phases, Equilibrium.

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Second Year Chemistry

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  1. Second Year Chemistry • 1st semester: Organic • 1st semester: Physical (2005-2006) • December exams • 2nd: Analytical & Environmental • 2nd: Inorganic • Summer exams • Physical: 3 lecturers Þ ~8 topics • Dónal Leech: four topics • Thermodynamics • Gases, Laws, Phases, Equilibrium

  2. Course Director • Dónal Leech • Room C205 (in Physical Chemistry) • E-mail: donal.leech@nuigalway.ie • Phone: 493563 (from outside), ext 3563 (internal phones) Web-site: http://www.nuigalway.ie/chem/Donal/home.htm

  3. Introduction Energetics and Equilibria What makes reactions “go”! This area of science is called THERMODYNAMICS Thermodynamics is expressed in a mathematical language BUT Don’t, initially anyway, get bogged down in the detail of the equations: try to picture the physical principle expressed in the equations We will develop ideas leading to one important Law, and explore practical applications along the way The Second Law of Thermodynamics

  4. Lecture Resources 12 lectures leading to four exam questions (section A, you must answer two from this section) • Main Text: “Elements of Physical Chemistry” Atkins & de Paula, 4th Edition (Desk reserve) http://www.oup.com/uk/booksites/content/0199271836/ OTHERS. “Physical Chemistry” Atkins & de Paula, 7th Edition or any other PChem textbook These notes available on NUI Galway web pages at http://www.nuigalway.ie/chem/degrees.htm See also excellent lecture notes from James Keeler, Cambridge, although topics are treated in a different running order than here.

  5. Course Structure • Revision of gases • Energy, heat and expansion work • 1st Law of thermodynamics • Thermochemistry and phase diagrams • Entropy • 2nd Law of thermodynamics • Chemical equilibria

  6. Revision States of Matter (bulk) Gas: fluid form that fills container Liquid: fluid form with well-defined surface, fills bottom of container (in gravitational field) Solid: retains its own inherent shape Difference between these states related to freedom of particles (molecules) to move past each other. We describe the macroscopic physical state of matter under conditions of volume, pressure, temperature and amount present.

  7. Blank-to be presented in Lecture

  8. Pressure (revision) • Pressure is the force that acts on a given area (P=F/A). • Gravity on earth exerts a pressure on the atmosphere: atmospheric pressure. • We can evaluate this by calculating the force due to acceleration (by gravity) of a 1m2 column of air extending through the atmosphere (this has a mass of ~10,000kg). This unit is a Newton (N) This unit is a Pascal (Pa)

  9. Units of Pressure S.I. unit of pressure is the N/m2, given the name Pascal (Pa). A related unit is the bar (1x105 Pa) used because atmospheric pressure is ~ 1x105 Pa (100 kPa, or 1bar). Torricelli (a student of Galileo) was the first to recognise that the atmosphere had weight, and measured pressure using a barometer Standard atmospheric pressure was thus defined as the pressure sufficient to support a mercury column of 760mm (units of mmHg, or torr). Another popular unit was thus introduced to simplify things, the atmosphere (atm = 760mmHg).

  10. Pressure • Atmospheric pressure and relationship between units 1 atm = 760 mmHg = 760 torr = 101.325kPa = 1.01325 bar) Measuring Pressure: the manometer Exercise: On a certain day a barometer gives the atmospheric pressure as 764.7 torr. If a metre stick is used to measure a height of 136.4mm in the open arm, and 103.8mm in the gas arm of a manometer, what is the pressure of the gas sample? (give in torr, atm, kPa and bar).

  11. Ideal Gas Law • Can specify state of sample by giving V, P, T and n. • These are however interdependent Equation of state of low-pressure gas is known (from combination of Boyle’s, Charles’s Laws and Avogadro’s principle) PV = nRT R = 8.314 J K-1 mol-1 (= NAk) (or L kPa K-1 mol-1 or m3 Pa K-1 mol-1)

  12. Boyle’s Law Living Graph of Boyle's Law

  13. Charles’s Law

  14. Avogadro principle • At a given T and p, equal volumes of gases contain the same number of molecules, Vm = V/n • Table below presents the molar volumes of selected gases at standard ambient temperature (298.15 K) and pressure (1 atm)

  15. Blank-to be presented in Lecture

  16. Gas mixtures • Dalton’s Law of partial pressures The total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone (partial pressure) PT=P1+P2+P3+….Pn Mole fractions: xi = ni/n Q: If dry air is composed of N2, O2, Ar at sea level in mass percent of 75.5: 23.2: 1.3. What is partial pressure for each when total pressure is 1.0 bar (100 kPa)?

  17. Kinetic model of gases • Based on 3 assumptions • Molecules are in ceaseless random motion • Size of molecules is negligible • Molecules do not interact Can derive: (see further information 1.1 in textbook) Where c is the root-mean-squared (rms) speed

  18. Kinetic Molecular Theory • Compare KMT to Ideal Gas Law

  19. Maxwell Distribution of Speeds • Not all molecules travel at the same speed • Distribution of speeds derived by James Clerk Maxwell

  20. Diffusion and Effusion Thomas Graham proposed a Law (1883) to summarize experimental observations on effusion Rate of Effusion  1/√M Relative rates of effusion

  21. Blank-to be presented in Lecture

  22. Critical Point • point at which surface separating two phases no longer appears: interface between vapour and liquid phases disappears, their densities become equal-supercritical fluid

  23. Compression factor • Small difference between real and perfect behaviour at high T, low p (see CO2 isotherms) • Model using virial equation of state • pVm = RT(1 + B’p + C’p2 + …) • More convenient expression • pVm = RT(1 + B/Vm + C/Vm2 + …) • In each case Z = expression in parentheses • B factor is most important, and is positive for H2, negative for others in the figure Z=1 for perfect gas. Deviations from this measure how far gases depart from ideal behaviour.

  24. Virial Coefficients and Boyle Temperature • Virial coefficients depend upon T • T at which Z  0 is called the Boyle Temperature (most like perfect gas) • pVm = RTB Although the virial equation of state is the most reliable, it does not provide much insight into the behaviour of gases Johannes van der Waals (Dutch physicist) proposed in 1873 an alternate approximate equation of state

  25. Van der Waals equation of state • Actual volume reduced in proportion to number of molecules present (repulsions) • Attractive forces reduce frequency of collisions and their strength • Parameters depend on the gas, but are taken to be independent of T. • a is large when attractions are large, b scales in proportion to molecular size (note units)

  26. Features of vdW equation • Reduces to perfect gas equation at high T and V • Liquids and gases coexist when attractions ≈ repulsions • Critical constants are related to coefficients. Flat inflexion of curve when T=Tc. • Can derive (by setting 1st and 2nd derivatives of equation to zero) expression for critical constants • Vc = 3b, pc = a/27b2, Tc =8a/27Rb • Can derive expression for the Boyle Temperature • TB = a/Rb

  27. Maxwell Construction Below Tc calculated vderW isotherms have oscillations that are unphysical. In the Maxwell construction these are replaced with horizontal lines, with equal areas above and below, to generate the isotherms.

  28. Blank-to be presented in Lecture

  29. Liquefaction-Irish Links! • Refrigeration developed by Carl von Linde in 19th Century, in response to a request from Guinness in Dublin for a new cooling technique. • Based upon the fact that gases cool as they expand: Joule-Thomson effect (William Thomson, later Lord Kelvin, born in Belfast), The Linde refrigerator combines the JT process with a counter-flow heat exchanger. The gas is re-circulated and it cools on expansion through the throttle. The cooled gas cools the high-pressure gas, which cools still further as it expands. Eventually liquefied gas drips from the throttle.

  30. Summary • Simplest state of matter is that of a gas • We can assemble an equation of state for an idealised gas from experimental results (Boyle, Charles, Avogadro) • Kinetic Molecular Theory can help explain the molecular basis for these Laws • Real gases differ from ideal gases because of inter-molecular interactions.

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