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CHEMISTRY 161 Chapter 5

CHEMISTRY 161 Chapter 5. Classification of Matter. solid liquid gas. 1. Gases. substances that exist in the gaseous phase under normal atmospheric conditions. T = 25 o C p = 1 atm. HF, HCl, HBr, HI CO, CO 2 CH 4 , NH 3 , H 2 S, PH 3 NO, NO 2 , N 2 O

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CHEMISTRY 161 Chapter 5

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  1. CHEMISTRY 161 Chapter 5

  2. Classification of Matter solid liquid gas

  3. 1. Gases substances that exist in the gaseous phase under normal atmospheric conditions T = 25oC p = 1 atm

  4. HF, HCl, HBr, HI CO, CO2 CH4, NH3, H2S, PH3 NO, NO2, N2O SO2

  5. Jupiter (H2, He) Io (SO2)

  6. Helix Nebula Orion Nebula

  7. 2. Pressure Ar EXP I molecules/atoms of gas are constantly in motion

  8. Atmospheric Pressure

  9. Standard Atmospheric Pressure 760 mm at 273 K at sea level Torricelli barometer 1 atm = 760 mm Hg = 760 torr pressure of the atmosphere is balanced by pressure exerted by mercury

  10. SI units force pressure = area p = F / A [p] = Nm-2 = kg m-1 s-2 = Pa

  11. pressure measurement manometer

  12. 3. Gas Laws 3.1. pressure p versus volume V 3.2. temperature T versus volume V 3.3. volume V versus amount n p, V, T, n

  13. 3.1. Boyle’s Law pressure – volume relationship (temperature is constant) Boyle (1627-1691)

  14. p ∞ 1/V EXP II

  15. p ∞ 1/V p = const/V p × V = const p1× V1 = const p2× V2 = const p1× V1 = p2× V2

  16. 3.2. Gay-Lussac’s Law temperature – volume relationship (pressure is constant) Gay-Lussac (1778-1850)

  17. V ∞ T EXP III

  18. V ∞ T V = const’ ×T V/T = const’ V1 / T1 = const’ V2 / T2 = const V1 / T1 = V2 / T2

  19. 3.3. Avogadro’s Law amount – volume relationship (pressure and temperature are constant) Avogadro (1776-1856)

  20. 2 H2(g) + O2(g) → 2 H2O(l) n ∞ V

  21. n ∞ V n = const’’ × V n/V = const’’ n1 / V1 = const’’ n2 / V2 = const’’ n1 / V1 = n2 / V2

  22. SUMMARY 3.1. Boyle’s Law 3.2. Gay-Lussac’s Law 3.3. Avogadro’s Law p ∞ 1/V V ∞ T n ∞ V

  23. 1. IDEAL GAS EQUATION (1) p ∞ 1/V V ∞ 1/p (2) V ∞ T V ∞ T (3) n ∞ V V ∞ n V ∞ T × n / p p × V = const × n × T

  24. p × V = const × n × T p × V = R × n × T p × V = n × R × T ideal gas equation

  25. p × V = n × R × T [R] = [p] × [V] / [n] / [T] m3 mol K Pa = N/m2 [R] = N × m / mol / K [R] = J / mol / K

  26. [R] = J / mol / K R = 8.314 J / mol / K ideal gas constant

  27. 2. MOLAR VOLUME What is the volume of 1 mol of a gas at 273.15 K (0oC) and 1 atm (101,325 Pa)? standard temperature and pressure (STP) p × V = n × R × T V = 22.4 l EXP IV

  28. p × V = n × R × T V = 22.4 l Vm = 22.4 l the molar volume at standard pressure and temperature is independent on the gas type

  29. 3. STOICHIOMETRY NaN3(s) → Na(s) + N2(g) How many liters of nitrogen gas are produced in the decomposition of 60.0 g sodium azide at 80oC and 823 torr? • Balancing • Mole ratios • Convert grams into moles • Convert moles into liters

  30. 4. DENSITY CALCULATION ς = m / V p × V = n × R × T V = n × R × T / p relate the moles (n) to the mass (m) via the molecular weight (M) m = n × M n = m / M ς = p × M / (R × T)

  31. 5. DALTON’S LAW pure gases gas mixtures (atmospheres) Dalton (1801)

  32. DALTON’S LAW the total pressure of a gas mixture, p, is the sum of the pressures of the individual gases (partial pressures) at a constant temperature and volume p = pA + pB + pC + …. EXP V

  33. p × V = n × R × T pA× V = nA× R × T pA = nA× R × T / V pB× V = nB× R × T pB = nB× R × T / V p = pA + pB p = (nA + nB) × R × T / V p × V = n × R × T

  34. p × V = (nA + nB) × R × T pA = nA× R × T / V pA / p = nA /(nA + nB) = xA mole fraction x < 1 pA = xA × p

  35. 2 KClO3→ 2 KCl + 3 O2 EXP VI/VII

  36. SUMMARY 1. ideal gas equation p × V = n × R × T R = 8.314 J / mol / K 2. molar volume Vm = 22.4 l

  37. 3. Density of gases ς = p × M / (R × T) 4. Dalton’s Law n p = Σ pi i=1

  38. 1. Kinetic Molecular Theory of Gases macroscopic (gas cylinder) microscopic (atoms/molecules) Maxwell (1831-1879) Boltzmann (1844-1906)

  39. Kinetic Energy of Gases physical properties of gases can be described by motion of individual gas atoms/molecules each macroscopic and microscopic particle in motion holds an energy (kinetic energy)

  40. Assumptions of the Kinetic Theory of Gases • gases are composed of atoms/molecules which are • separated from each other by a distance l much more than their • own diameter d d = 10-10 m l = 10-3 m….. few m molecules are mass points with negligible volume l

  41. 2. gases are constantly in motion in random reactions and hold a kinetic energy gases collide and transfer energy (billiard ball model)

  42. 3. gases atoms/molecules do not exert forces on each other (absence of intermolecular interactions) F(inter) = 0 p(inter) = 0

  43. Gas Diffusion

  44. 2.Distribution of Molecular Speeds Maxwell-Boltzmann distribution

  45. 3.Real Gases p × V = n × R × T (n = 1) deviation of ideal gas law at high pressures p ≈ 90 atm

  46. North America Nebula p << 10-10 atm

  47. ideal gas law p V = n R T real gas law (van der Waals equation) (p + (a n2 / V2) ) (V – n b) = n R T corrected volume (volume occupied by molecules) corrected pressure (additional pressure/force from attraction)

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