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Me cchapter 5


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Me cchapter 5

  1. 1. Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 5 States of Matter Gases, Liquids, and SolidsDennistonToppingCaret7th Edition
  2. 2. Changes in State• Changes in state are considered to be physical changes• During a change of physical state many other physical properties may also change• This chapter focuses on the important differences in physical properties among – Gases – Liquids – Solids
  3. 3. Comparison of Physical Properties of Gases, Liquids, and Solids
  4. 4. 5.1 The Gaseous State Ideal Gas Concept• Ideal gas - a model of the way that particles of a gas behave at the microscopic level• We can measure the following of a gas: – temperature – volume We can systematically change one of the properties and see – pressure the effect on the others – mass
  5. 5. 5.1 The Gaseous State Measurement of Gases • Gas laws involve the relationship between: – number of moles (n) of gas – volume (V) – temperature (T) – pressure (P) • Pressure - force per unit area • Gas pressure is a result of force exerted by the collision of particles with the walls of the container
  6. 6. 5.1 The Gaseous State Barometer • Measures atmospheric pressure – Invented by Evangelista Torricelli • Common units of pressure – atmosphere (atm) – torr (in Torricelli’s honor) – pascal (Pa) (in honor of Blaise Pascal) • 1 atm is equal to: – 760 mmHg – 760 torr – 76 cmHg
  7. 7. 5.1 The Gaseous State Kinetic Molecular Theory of Gases 1. Gases are made up of small atoms or molecules that are in constant, random motion 2. The distance of separation is very large compared to the size of the individual atoms or molecules – Gas is mostly empty space 1. All gas particles behave independently – No attractive or repulsive forces exist between them
  8. 8. 5.1 The Gaseous State Kinetic Molecular Theory of Gases 4. Gas particles collide with each other and with the walls of the container without losing energy – The energy is transferred from one atom or molecule to another 4. The average kinetic energy of the atoms or molecules increases or decreases in proportion to absolute temperature – As temperature goes up, particle speed goes up
  9. 9. Kinetic Molecular Theory of Gases5.1 The Gaseous State Explains the following statements: • Gases are easily compressible – gas is mostly empty space, room for particles to be pushed together • Gases will expand to fill any available volume – move freely with sufficient energy to overcome attractive forces • Gases have low density – being mostly empty space; gases have low mass per unit volume
  10. 10. • Gases readily diffuse through each other – they are5.1 The Gaseous State in continuous motion with paths readily available due to large space between adjacent particles • Gases exert pressure on their containers – pressure results from collisions of gas particles with the container walls • Gases behave most ideally at low pressure and high temperature – Low pressure, average distance of separation is greatest, minimizing interactive forces – High temperature, rapid motion overcomes interactive forces more easily
  11. 11. Ideal Gases vs. Real Gases5.1 The Gaseous State • In reality there is no such thing as an ideal gas – It is a useful model to explain gas behavior • Nonpolar gases behave more ideally than polar gases because attractive forces are present in polar gases
  12. 12. 5.1 The Gaseous State Gas Diffusion Ammonia (17.0 g/mol) Hydrogen chloride (36.5 g/mol) Ammonia diffused farther in same time, lighter moves faster
  13. 13. 5.1 The Gaseous State Boyle’s Law • Boyle’s law - volume of a gas varies inversely with the pressure exerted by the gas if the temperature and number of moles are held constant • The product of pressure (P) and volume (V) is a constant PV = k 1 • Used to calculate – Volume resulting from pressure change – Pressure resulting from volume change PiVi = PfVf
  14. 14. 5.1 The Gaseous State Application of Boyle’s Law • Gas occupies 10.0 L at 1.00 atm pressure • Product, PV = (10.0 L) (1.00 atm) = k1 • Double the pressure to 2.0 atm, decreases the volume to 5.0 L – (2.0 atm)(Vx) = (10.0 L)(1.00 atm) – Vx = 5.0 L
  15. 15. 5.1 The Gaseous State Boyle’s Law Practice 1. A 5.0 L sample of a gas at 25oC and 3.0 atm is compressed at constant temperature to a volume of 1.0 L. What is the new pressure? 2. A 3.5 L sample of a gas at 1.0 atm is expanded at constant temperature until the pressure is 0.10 atm. What is the volume of the gas?
  16. 16. 5.1 The Gaseous State Charles’s Law • It is possible to relate gas volume and temperature • Charles’s law - volume of a gas varies directly with the absolute temperature (K) if pressure and number of moles of gas are constant • Ratio of volume (V) and temperature (T) is a constant V V = k2 Vi f = T Ti T f
  17. 17. 5.1 The Gaseous State Application of Charles’s Law • If a gas occupies 10.0 L at 273 K with V/T = k2 • Doubling temperature to 546 K, increases volume to 20.0 L 10.0 L / 273 K = Vf / 546 K
  18. 18. 5.1 The Gaseous State Practice with Charles’s Law 1. A 2.5 L sample of gas at 25oC is heated to 50oC at constant pressure. Will the volume double? 2. What would be the volume? 3. What temperature would be required to double the volume?
  19. 19. 5.1 The Gaseous State Combined Gas Law • If a sample of gas undergoes change involving volume, pressure, and temperature simultaneously, use the combined gas law • Derived from a combination of Boyle’s law and Charles’s law PV PiVi f f = Ti Tf
  20. 20. 5.1 The Gaseous State Using the Combined Gas Law • Calculate the volume of N2 resulting when 0.100 L of the gas is heated from 300. K to 350. K at 1.00 atm • What do we know? PiVi Pf V f = – Pi = 1.00 atm Pf = 1.00 atm Ti Tf – Vi = 0.100 L Vf = ? L – Ti = 300. K Tf = 350. K • Vf = ViTf / Ti this is valid as Pi = Pf • Vf = (0.100 L)(350. K) / 300. K = 0.117 L • Note the decimal point in the temperature to indicate significance
  21. 21. 5.1 The Gaseous State Practice With the Combined Gas Law Calculate the temperature when a 0.50 L sample of gas at 1.0 atm and 25oC is compressed to 0.05 L of gas at 5.0 atm.
  22. 22. 5.1 The Gaseous State Avogadro’s Law • Avogadro’s law - equal volumes of any ideal gas contain the same number of moles if measured under the same conditions of temperature and pressure V = k3 n • Changes in conditions can be calculated by rewriting the equation Vi V f = ni n f
  23. 23. 5.1 The Gaseous State Using Avogadro’s Law • If 5.50 mol of CO occupy 20.6 L, how many liters will 16.5 mol of CO occupy at the same temperature and pressure? • What do we know? – Vi = 20.6 L Vf = ? L – ni = 5.50 mol nf = 16.5 mol – Vf = Vinf / ni = (20.6 L)(16.5 mol) (5.50 mol) = 61.8 L CO
  24. 24. 5.1 The Gaseous State Molar Volume of a Gas • Molar volume - the volume occupied by 1 mol of any gas • STP – Standard Temperature and Pressure – T = 273 K (or 0oC) – P = 1 atm • At STP the molar volume of any gas is 22.4 L
  25. 25. Gas Densities5.1 The Gaseous State • Density = mass / volume • Calculate the density of 4.00 g He – What is the mass of 1 mol of H2? 4.00 g DensityHe = 4.00g / 22.4L = 0.178 g/L at STP
  26. 26. 5.1 The Gaseous State The Ideal Gas Law • Combining: – Boyle’s law (relating volume and pressure) – Charles’s law (relating volume and temperature) – Avogadro’s law (relating volume to the number of moles) gives the Ideal Gas Law PV=nRT • R is a constant, ideal gas constant • R = 0.0821 L.Atm/mol.K If units are P in atm, V in L, n in number of moles, T in K
  27. 27. 5.1 The Gaseous State Calculating a Molar Volume • Demonstrate molar volume of O2 gas at STP L ⋅ atm 1mol(0.08206 )273 K nRT mol ⋅ K V= = = 22.4 L P 1 atm
  28. 28. Practice Using the Ideal Gas Law5.1 The Gaseous State 1. What is the volume of gas occupied by 5.0 g CH4 at 25oC and 1 atm? 2. What is the mass of N2 required to occupy 3.0 L at 100oC and 700 mmHg?
  29. 29. 5.1 The Gaseous State Dalton’s Law of Partial Pressures • Dalton’s law – a mixture of gases exerts a pressure that is the sum of the pressures that each gas would exert if it were present alone under the same conditions Pt=p1+p2+p3+... • Total pressure of our atmosphere is equal to the sum of the pressures of N2 and O2 – (principal components of air) Pair = p N + pO 2 2
  30. 30. 5.2 The Liquid State• Liquids are practically incompressible – Enables brake fluid to work in your car• Viscosity - a measure of a liquid’s resistance to flow – A function of both attractive forces between molecules and molecular geometry – Flow occurs because the molecules can easily slide past each other • Glycerol - example of a very viscous liquid – Viscosity decreases with increased temperature
  31. 31. 5.2 The Liquid State Surface Tension • Surface tension - a measure of the attractive forces exerted among molecules at the surface of a liquid – Surface molecules are surrounded and attracted by fewer liquid molecules than those below – Net attractive forces on surface molecules pull them downward • Results in “beading” • Surfactant - substance added which decreases the surface tension, for example – soap
  32. 32. 5.2 The Liquid State Vapor Pressure of a Liquid • Place water in a sealed container – Both liquid water and water vapor will exist in the container • How does this happen below the boiling point? – Temperature is too low for boiling conversion • Kinetic theory - liquid molecules are in continuous motion, with their average kinetic energy directly proportional to the Kelvin temperature
  33. 33. Temperature Dependence of Liquid Vapor Pressure5.2 The Liquid State energy + H2O(l) → H2O(g) • Average molecular kinetic energy increases as does temperature • Some high energy molecules have sufficient energy to escape from the liquid phase • Even at cold temperatures, some molecules can be converted
  34. 34. Movement From Gas Back to Liquid5.2 The Liquid State H2O(g) → H2O(l) + energy • Molecules in the vapor phase can lose energy and be converted back to the liquid phase • Evaporation - the process of conversion of liquid to gas at a temperature too low to boil • Condensation - conversion of gas to the liquid state
  35. 35. Liquid Water in Equilibrium With Water Vapor5.2 The Liquid State • When the rate of evaporation equals the rate of condensation, the system is at equilibrium • Vapor pressure of a liquid - the pressure exerted by the vapor at equilibrium
  36. 36. 5.2 The Liquid State Boiling Point • Boiling point - the temperature at which the vapor pressure of the liquid becomes equal to the atmospheric pressure • Normal boiling point - temperature at which the vapor pressure of the liquid is equal to 1 atm • What happens when you go to a mountain where the atmospheric pressure is lower than 1 atm? – The boiling point lowers • Boiling point is dependant on the intermolecular forces – Polar molecules have higher b.p. than nonpolar molecules
  37. 37. 5.2 The Liquid State Van der Waals Forces • Physical properties of liquids are explained in terms of their intermolecular forces • Van der Waals forces are intermolecular forces having 2 subtypes – Dipole-dipole interactions – Attractive forces between polar molecules – London forces – As electrons are in continuous motion, a nonpolar molecule could have an instantaneous dipole
  38. 38. 5.2 The Liquid State London Forces • Exist between all molecules • The only attractive force between nonpolar atoms or molecules • Electrons are in constant motion • Electrons can be, in an instant, arranged in such a way that they have a dipole (Instantaneous dipole) • The temporary dipole interacts with other temporary dipoles to cause attraction
  39. 39. 5.2 The Liquid State Hydrogen Bonding • Hydrogen bonding: – not considered a Van der Waals force – is a special type of dipole-dipole attraction – is a very strong intermolecular attraction causing higher than expected b.p. and m.p. • Requirement for hydrogen bonding: – molecules have hydrogen directly bonded to O, N, or F
  40. 40. Examples of Hydrogen Bonding5.2 The Liquid State • Hydrogen bonding has an extremely important influence on the behavior of many biological systems • H2O • NH3 • HF
  41. 41. 5.3 The Solid State• Particles highly organized, in a defined fashion• Fixed shape and volume• Properties of solids: – incompressible – m.p. depends on strength of attractive force between particles – crystalline solid - regular repeating structure – amorphous solid - no organized structure
  42. 42. Types of Crystalline Solids5.3 The Solid State 1. Ionic solids • held together by electrostatic forces • high m.p. and b.p. • hard and brittle • if dissolves in water, electrolytes • NaCl 2. Covalent solids • held together entirely by covalent bonds • high m.p. and b.p. • extremely hard • diamond
  43. 43. 3.Molecular solids5.3 The Solid State • molecules are held together with intermolecular forces • often soft • low m.p. • often volatile • ice 4.Metallic solids • metal atoms held together with metal bonds • metal bonds – overlap of orbitals of metal atoms – overlap causes regions of high electron density where electrons are extremely mobile - conducts electricity
  44. 44. Four Types of Crystalline Solids5.3 The Solid State