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Chemistry 231

Chemistry 231 . Introduction and Gases . Physical Chemistry. Physics - study of the properties of matter that are shared by all substances Chemistry - the study of the properties of the substances that make up the universe and the changes that these substances undergo

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Chemistry 231

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  1. Chemistry 231 Introduction and Gases

  2. Physical Chemistry Physics - study of the properties of matter that are shared by all substances Chemistry - the study of the properties of the substances that make up the universe and the changes that these substances undergo Physical Chemistry - the best of both worlds!

  3. Thermodynamics and Thermochemistry Thermodynamics – the study of energy and its transformations Thermochemical changes – energy changes associated with chemical reactions

  4. Studying Systems • Interested in the numerical values of the state variables (defined later) that quantify the systems at that point in time. • Systems can be either • macroscopic • microscopic

  5. State of a System • Described by variables such as • temperature (T) • pressure (P) • volume (V) • energy (U) • enthalpy (H) • Gibbs energy (G)

  6. State and Path Functions • State Variables • system quantity whose values are fixed at constant temperature, pressure, composition • State Function • a system property whose values depends only on the initial and final states of the system. • Path Functions • system quantity whose value is dependent on the manner in which the transformation is carried out.

  7. State and Path Functions (Continued) • Examples of state functions • H •  G •  V •  T • Examples of path functions • work (w) • heat (q)

  8. Equilibrium vs. Metastable Metastable - the progress towards the equilibrium state is slow Equilibrium state - state of the system is invariant with time

  9. Reversible and Irreversible Reversible transformation - the direction of the transformation can be reversed at any time by some infinitesimal change in the surroundings Irreversible transformation - the system does not attain equilibrium at each step of the process

  10. The Definition of a Gas Gas - a substance that is characterised by widely separated molecules in rapid motion Mixtures of gases are uniform. Gases will expand to fill containers.

  11. Examples of Gaseous Substances • Common gases include - O2 and N2, the major components of "air" • Other common gases - F2, Cl2, H2, He, and N2O (laughing gas)

  12. The Definition of Pressure The pressure of a gas is best defined as the forces exerted by gas on the walls of the container Define P = force/area The SI unit of pressure is the Pascal 1 Pa = N/m2 = (kg m/s2)/m2

  13. The Measurement of Pressure • How do we measure gas pressure? • We use an instrument called the barometer - invented by Torricelli • Gas pressure conversion factors • 1 atm = 760 mm Hg = 760 torr • 1 atm = 101.325 kPa = 1.01325 bar • 1 bar = 1 x 105 Pa (exactly)

  14. The Gas Laws • Experiments with a wide variety of gases revealed that four variables were sufficient to fully describe the state of a gas • Pressure (P) • Volume (V) • Temperature (T) • The amount of the gas in moles (n)

  15. Boyle's Law • The gas volume/pressure relationship • The volume occupied by the gas is inversely proportional to the pressure • V 1/P • note temperature and the amount of the gas are fixed

  16. Boyle's Law V V 1/P P

  17. Charles and Gay-Lussac's Law Defines the gas volume/temperature relationship. V  T (constant pressure and amount of gas) Note T represents the temperature on the absolute (Kelvin) temperature scale

  18. Charles and Gay-Lussac's Law V t / C Absolute Zero (-273C = 0 K)

  19. The Kelvin temperature scale Lord Kelvin – all temperature/volume plots intercepted the tc axis at -273.15°C). Kelvin termed this absolute 0 – the temperature where the volume of an ideal gas is 0 and all thermal motion ceases!

  20. The Temperatures Scales • T (K) = [ tc (°C) + 273.15°C] K/°C • Freezing point of water: tc = 0 °C; T = 273.15 K • Boiling point of water: tc = 100 °C; T = 373.15 K • Room temperature: tc = 25 °C; T = 298 K • NOTE tc = °C; T (K) = K NO DEGREE SIGN

  21. Amonton’s Law The pressure/temperature relationship For a given quantity of gas at a fixed volume, P  T, i.e., if we heat a gas cylinder, P increases!

  22. Avogadro’s Law The volume of a gas at constant T and P is directly proportional to the number of moles of gas V  n => n = number of moles of gas

  23. The Ideal Gas Equation of State • We have four relationships • V  1/P; Boyle’s law • V  T; Charles’ and Gay-Lussac's law • V  n; Avogadro’s law • P  T; Amonton’s law

  24. The Ideal Gas Law • Combine these relationships into a single fundamental equation of state - the ideal gas equation of state

  25. The Definition of an Ideal Gas An ideal gas is a gas that obeys totally the ideal gas law over its entire P-V-T range Ideal gases – molecules have negligible intermolecular attractive forces and they occupy a negligible volume compared with the container volume

  26. Standard Temperature and Pressure • Define: STP (Standard Temperature and Pressure) • Temperature - 0.00 °C = 273.15 K • Pressure - 1.000 atm • The volume occupied by 1.000 mole of an ideal gas at STP is 22.41 L!

  27. Standard Ambient Temperature and Pressure • Define: SATP (Standard Ambient Temperature and Pressure) • Temperature - 25.00 °C = 273.15 K • Pressure - 1.000 bar (105 Pa) • The volume occupied by 1.000 mole of an ideal gas at SATP is 24.78 L!

  28. Partial Pressures 2 1 2 2 1 1 2 1 2 2 1 1 1 2 Let's consider two ideal gases (gas 1 and gas 2) in a container of volume V.

  29. Partial Pressures • The pressure exerted by gas #1 • P1 = n1 RT / V • The pressure exerted by gas #2 • P2= n2 RT / V • The total pressure of the gases • pT = nT RT / V • nT represents the total number of moles of gas present in the mixture

  30. Partial Pressures (continued) • P1 and P2 are the partial pressures of gas 1 and gas 2, respectively. • PT = P1 + P2 = nT (RT/V) • PT = P1 + P2 + P3 = j PJ • note Pj is known as the partial pressure of gas j

  31. Dalton's Law of Partial Pressure • Gaseous mixtures - gases exert the same pressure as if they were alone and occupied the same volume. • The partial pressure of each gas, Pi, is related to the total pressure by Pi = Xi PT • Xj is the mole fraction of gas i. • Xj= nj / nT

  32. Ideal Gas Temperature Scale In the limit of low pressures

  33. The Isothermal Compressibility The Isothermal Compressibility

  34. Coefficient of Thermal Expansion The coefficient of thermal expansion

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