Gas and condensed matter


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Gas and condensed matter

  2. 2. Bonding model for covalent molecular substances  1. 2. Bonding for covalent molecular substances falls into two categories The strong forces of attraction which holds atoms together within molecules The weak forces of attraction between molecules
  3. 3. Forces between molecules (intermolecular forces) will learn about the forces between molecules or compounds are called intermolecular forces  Inter means between or among  Internet, interstate, international  What would Interstellar travel be?  we
  4. 4. Intermolecular Forces  Forces that occur between molecules.
  5. 5. Intramolecular forces  What would intramolecular forces be?  Forces within molecules e.g covalent, metallic or ionic  intra means within  Intrastate, intranet, intracellular  Most of the intermolecular forces we look at occur between covalently bonded molecules or covalent molecular substances  Intramolecular bonds are stronger than intermolecular forces.
  6. 6. Overview  All matter is held together by force.  The force between atoms within a molecule is a chemical or intramolecular force.  The force between molecules are a physical or intermolecular force.  These physical forces are what we overcome when a chemical changes its state (e.g. gas liquid).
  7. 7. What causes intermolecular forces?  Molecules are made up of charged particles: positive nuclei and negative electrons.  When one molecule approaches another there is a multitude of forces between the particles in the two molecules.  Each electron in one molecule is attracted to the nuclei in the other molecule but also repelled by the electrons in the other molecule.  The same applies for nuclei
  8. 8. Types of Intermolecular forces  The three main types of intermolecular forces are: 1. Dipole-dipole attraction occur only btw polar molecules 2. H bonding – only with Hydrogen and Oxygen, Fluorine and Nitrogen) 3. Dispersion forces (London Dispersion Forces)
  9. 9. Intermolecular Forces  Forces that occur between molecules.  Dipole–dipole forces   Hydrogen bonding London dispersion forces
  10. 10. Dipole–Dipole Attraction
  11. 11. Dipole-Dipole Forces  Dipole moment – molecules with polar bonds often behave in an electric field as if they had a center of positive charge and a center of negative charge.  Molecules with dipole moments can attract each other electrostatically. They line up so that the positive and negative ends are close to each other.  Only about 1% as strong as covalent or ionic bonds.
  12. 12. Hydrogen Bonding Strong dipole-dipole forces. Hydrogen is bound to a highly electronegative atom – nitrogen, oxygen, or fluorine.
  13. 13. Hydrogen Bonding in Water  Blue dotted lines are the intermolecula r forces between the water molecules.
  14. 14. Hydrogen Bonding Affects  physical properties Boiling point
  15. 15. London Dispersion Forces    Instantaneous dipole that occurs accidentally in a given atom induces a similar dipole in a neighboring atom. Significant in large atoms/molecules. Occurs in all molecules, including nonpolar ones.
  16. 16. London Dispersion Forces Nonpolar Molecules
  17. 17. London Dispersion Forces  Become stronger as the sizes of atoms or molecules increase.
  18. 18. Melting and Boiling Points  In general, the stronger the intermolecular forces, the higher the melting and boiling points.
  19. 19. Strength of Intermolecular Interactions Hydrogen Bonding ↑ Dipole – Dipole ↑ London Dispersion Forces
  20. 20. Kinetic Molecular Theory The kinetic theory of matter is based on the following postulates: 1. Matter is composed of small particles called molecules 2. The particles are in constant random motion They possess kinetic energy due to their motion 3. There are repulsive and attractive forces between particles. They posses potential energy due to these forces 4. Average particle speed increases with temperature 5. No energy is lost when the particles collide, called elastic collision
  21. 21. Kinetic Molecular Theory The kinetic energy of a particle is given by the equation: 1 KE = mv 2 Where: 2 m = particle mass in kg v = particle velocity in m/s KE = kg-m2/s2 = j (joule) According to postulate 4 of our kinetic theory particle velocity increases with temperature. This means as temperature increases then kinetic energy increases.
  22. 22. Potential Energy Potential energy is the sum of the attractive and repulsive forces between particles. Examples of these types of forces are the gravitational attractive forces between objects and the repulsive forces between the same poles of magnets. Alternatively we can say forces between particles may be either cohesive or disruptive.
  23. 23. Interparticle Forces Cohesive forces include dipole-dipole interactions, dispersion forces, attraction between oppositely charged ions. Cohesive forces are largely temperature independent. e.g. magnets and gravity function the same way at different temperature.
  24. 24. Interparticle Forces Disruptive forces are those forces that make particles move away from each other. These forces result predominately from the particle motion. Disruptive forces increase with temperature in agreement with postulate 4. We can conclude that as we increase the temperature particles will become further apart from each other.
  25. 25. Tutorial 1  State and describe briefly three (3) main types of intermolecular forces.  State five (5) assumption in the kinetic molecular theory
  26. 26. A Gas  Has neither a definite volume nor shape.  Uniformly fills any container.  Mixes completely with any other gas  Exerts pressure on its surroundings.
  27. 27. Earth-like Atmosphere Composition of Earth’s Atmosphere Compound %(Volume) Mole Fractiona Nitrogen Oxygen Argon Carbon dioxide 78.08 20.95 0.934 0.033 0.7808 0.2095 0.00934 0.00033 Methane Hydrogen 2 x 10-4 5 x 10-5 2 x 10-6 5 x 10-7 a. mole fraction = mol component/total mol in mixture.
  28. 28. A mercury barometer The column height is proportional to the atmospheric pressure. Atmospheric pressure results from the mass of the atmosphere and gravitational forces. The pressure is the force per unit area. P = F/A 1 atm = 760 mmHg 1 atm = 1.01325 E5 Pa 1 mmHg = 1 torr
  29. 29. Units for Expressing Pressure Unit Atmosphere Pascal (Pa) Kilopascal (kPa) mmHg Torr Bar mbar psi Value 1 atm 1 atm = 1.01325 x 105 Pa 1 atm = 101.325 kPa 1 atm = 760 mmHg 1 atm = 760 torr 1 atm = 1.01325 bar 1 atm = 1013.25 mbar 1 atm = 14.7 psi
  30. 30. Pressure is equal to force/unit area  SI units = Newton/meter2 = 1 Pascal (Pa)  1 standard atmosphere = 101,325 Pa (100,000 Pa = 1 bar)  1 standard atmosphere = 1 atm =  760 mm Hg = 760 torr = 1013.25 hPa = 14.695 psi Meteorologists often report pressure in millibar; 1 mbar =0.001bar =0.1 kPa = 1hPa
  31. 31. Variables Affecting Gases  Pressure (P)  Volume (V)  Temperature  Number (T) of Moles (n)
  32. 32. Elevation and Atmospheric Pressure
  33. 33. Manometer Manometers are used to measure gas pressure in closed systems. For instance in a reaction vessel.
  34. 34. The Gas Laws and Absolute Temperature The relationship between the volume, pressure, temperature, and mass of a gas is called an equation of state. We will deal here with gases that are not too dense. Boyle’s Law: the volume of a given amount of gas is inversely proportional to the pressure as long as the temperature is constant.
  35. 35. Boyle Law…pressure is inversely proportional to volume (at constant T and moles, n).
  36. 36. The Gas Laws and Absolute Temperature The volume is linearly proportional to the temperature, as long as the temperature is somewhat above the condensation point and the pressure is constant: Extrapolating, the volume becomes zero at −273.15°C; this temperature is called absolute zero.
  37. 37. Avogadro’s Law  For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas (at low pressures). V α n V1 = V2 n1 n2
  38. 38. The Gas Laws and Absolute Temperature The concept of absolute zero allows us to define a third temperature scale – the absolute, or Kelvin, scale. This scale starts with 0 K at absolute zero, but otherwise is the same as the Celsius scale. Therefore, the freezing point of water is 273.15 K, and the boiling point is 373.15 K. Finally, when the volume is constant, the pressure is directly proportional to the temperature:
  39. 39. Combined Gas Law  Combining the gas laws the relationship P α T(n/V) can be obtained.  If n (number of moles) is held constant, then PV/T = constant. P1 V P2 V 1 2 = T T 1 2
  40. 40. Ideal Gas Law PV = nRT R = universal gas constant = 0.08206 L atm K-1 mol-1 P = pressure in atm V = volume in liters n = moles T = temperature in Kelvin
  41. 41. Standard Temperature and Pressure (for gases) “STP”     P = 1 atmosphere T = 0°C The molar volume of an ideal gas is 22.42 liters at STP (put 1 mole, 1 atm, R, and 273 K in the ideal gas law and calculate V) Note STP is different for other phases, e.g. solutions or enthalpies of formation.
  42. 42. The Ideal Gas Law A mole (mol) is defined as the number of grams of a substance that is numerically equal to the molecular mass of the substance: 1 mol H2 has a mass of 2 g 1 mol Ne has a mass of 20 g 1 mol CO2 has a mass of 44 g The number of moles in a certain mass of material:
  43. 43. The Ideal Gas Equation • Charles’s Law: 1 V ∝ (constant n, T ) P V ∝ T (constant n, P ) • Avogadro’s Law: V ∝ n (constant P, T ) • Boyle’s Law: • We can combine these into a general gas law: nT V∝ P
  44. 44. The Ideal Gas Equation • R = gas constant, then  nT  V = R   P  • The ideal gas equation is: PV = nRT • R = 0.08206 L·atm/mol·K = 8.3145 J/mol·K • J = kPa·L = kPa·dm3 = Pa·m3 • Real Gases behave ideally at low P and high T.
  45. 45. Ideal Gas Law in Terms of Molecules: Avogadro’s Number Since the gas constant is universal, the number of molecules in one mole is the same for all gases. That number is called Avogadro’s number: The number of molecules in a gas is the number of moles times Avogadro’s number:
  46. 46. Ideal Gas Law in Terms of Molecules: Avogadro’s Number Therefore we can write: where k is called Boltzmann’s constant. (13-4)
  47. 47. The Ideal Gas Equation  Calculate the pressure exerted by 84.0 g of ammonia, NH3, in a 5.00 L container at 200. oC using the ideal gas law. PV = nRT P = nRT/V n = 84.0g * 1mol/17 g T = 200 + 273 P = (4.94mol)(0.08206 L atm mol-1 K-1)(473K) (5 L) P = 38.3 atm
  48. 48. Tutorial 2.  The pressure on a sample of an ideal gas was increased from 715 torr to 3.55 atm at constant temperature. If the initial volume of the gas was 485. mL, what would be the final volume? A 7.9 L sample of gas was cooled from 79°C to a temperature at which the volume of the gas was 4.3 L. Assuming the pressure remains constant, calculate the final temperature.  Calculate the pressure in atmospheres and pascals of a 1.2 mol sample of methane gas in a 3.3 L container at 25°C.
  49. 49. Real Gases: Deviations from Ideality    Real gases behave ideally at ordinary temperatures and pressures. At low temperatures and high pressures real gases do not behave ideally. The reasons for the deviations from ideality are: 1. The molecules are very close to one another, thus their volume is important. 2. The molecular interactions also become important. J. van der Waals, 1837-1923, Professor of Physics, Amsterdam. Nobel Prize 1910.
  50. 50. Real Gases: Deviations from Ideality  van der Waals’ equation accounts for the behavior of real gases at low temperatures and high pressures.  n 2a  V − nb) = nRT P + 2 ( V   • The van der Waals constants a and b take into account two things: 1. a accounts for intermolecular attraction a. b. For nonpolar gases the attractive forces are London Forces For polar gases the attractive forces are dipole-dipole attractions or hydrogen bonds. 2. b accounts for volume of gas molecules At large volumes a and b are relatively small and van der Waal’s equation reduces to ideal gas law at high temperatures and low pressures.
  51. 51. Real Gases: Deviations from Ideal Behavior The van der Waals Equation nRT n 2a P= − 2 V − nb V Corrects for Corrects for molecular molecular volume attraction • General form of the van der Waals equation:  n 2a  P + ( V − nb ) = nRT  V2   
  52. 52. Example
  53. 53. Condensed matter : the three states of matter.
  54. 54. Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior gas liquid solid assumes the shape and volume of its container particles can move past one another assumes the shape of the part of the container which it occupies particles can move/slide past one another retains a fixed volume and shape rigid - particles locked into place compressible lots of free space between particles not easily compressible little free space between particles not easily compressible little free space between particles flows easily particles can move past one another flows easily particles can move/slide past one another does not flow easily rigid - particles cannot move/slide past one another
  55. 55. Clearly, a theory used to describe the condensed states of matter must include an attraction between the particles in the substance  Condensed of Matter: .  Liquids  Solids States
  56. 56. Kinetic Theory Description of the Liquid State.  Like gases, the condensed states of matter can consist of atoms, ions, or molecules.  What separates the three states of matter is the proximity of the particles in the substance.  For the condensed states of matter the particles are close enough to interact.
  57. 57. Phase Changes
  58. 58. Triple Point Diagram of Water  Regions: Each region corresponds to one phase which is stable for any combination of P and T within its region  Lines Between Region: Lines separating the regions representing phase-transition curves  Triple Point: The triple point represents the P and T at which all 3 phases coexist in equilibrium  Critical Point: At the critical point the vapor pressure cannot be condensed to liquid no matter what pressure is applied.
  59. 59. Tutorial 3  Van der Waals, realized that two of the assumptions mentioned above were questionable. He then developed the Van der Waals equation of state which predicted the formation of liquid phase. Write the equation and state two corrections that he made.  (a) Calculate the pressure exerted by 1.00 mol of CO2 in a 1.00 L vessel at 300 K, assuming that the gas behaves ideally. (b) Repeat the calculation by using the van der Waals equation.  Sketch and label the liquid region, gas region solid region and triple point in water phase diagram