Gases. T boil, K. 11 gaseous elements. Rare gases. He, Ne, Ar, Kr, Xe, Rn Diatomic gaseous elements H 2, N 2, O 2, F 2, Cl 2
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1 Gases Gas T boil, K Rare gases 11 gaseous elements He, Ne, Ar, Kr, Xe, Rn 165 Rn 211 N 2 O 2 77 F Diatomic gaseous elements Cl H 2, N 2, O 2, F 2, Cl 2 H 2 He Ne Ar Kr Xe
2 Gas Large part of chemical and physical theories was developed on experiments with gases. Different kinds of air studies on gases, concept of gas Gas sylvestre = CO 2 CO 2 is formed: Burning coal with KNO 3 (salpeter) Fermentation of beer and wine Action of vinegar on limestone Johann Baptista van Helmont Grotto del Cane ( ) 2
3 Collisions of gas molecules with container walls Pressure Pa 760 mm Hg 760 torr (Torricelli) 1 atm F = force, N A = area, m 2 p = F A Evangelista Torricelli ( ) barometer
4 F = m g g = m s 2 Atmospheric Pressure Column of air 1 m 2, from ground to stratosphere m = 10 4 kg Atmospheric Pressure 1 atm 4
5 p = h ρ g Hydrostatic Pressure 5
6 1662 Boyle s Law Product of gas pressure and its volume is a constant for a given amount of gas at constant temperature Isothermic p V = const. Robert Boyle ( ) Does not depend on the kind of gas, mixture of gases Exceptions NO 2 6
7 Gas Compression p = V p V at constant temperature p = h ρ g tlak roste 7
8 Boyle s Law V = const. / p Objem, ml 8
9 1/V = const. p 9
10 p = konst V Isotherms T = const. p = V p V
11 Air in a tank for a 60 minute dive at surface Scuba Diving Depth, m Pressure, atm How much air? X minutes in 30 m 11
12 1787 Charles-Gay-Lussac Law p = const. Isobaric Different gases expand by the same fraction of volume on heating by the same degree Jacques A. C. Charles ( ) The first solo flight in ballon The first H 2 ballon Joseph Louis Gay-Lussac ( ) 12
13 Charles-Gay-Lussac Law V, cm 3 V = a t + b Temp, o C p = const. Isobaric 13
14 Charles-Gay-Lussac Law V = a t + b p = const. Isobaric V = a t + b V = a (t + b/a) b/a = 273 C absolute temperature scale V = k T T = absolute temperature [K] Concept of absolute zero temperature 14
15 Charles-Gay-Lussac Law V, cm 3 V = k T p = const. Isobaric Temp, K 15
16 Isobars V = a t + b p = const. Temp, o C 16
17 Charles-Gay-Lussac Law V = V 0 (1 + α t) α = 1/273 coefficient of thermal expansion t = temp. in C V 1 = T 1 V T 2 2 const. n and p 17
18 Amonton s Law p = p 0 (1 + α t) α = 1/273 coefficient of thermal expansion p 1 = T 1 p T 2 2 const. n and V isochore 18
19 Law of Constant Volumes (Gay-Lussac) 1809 Gases combine in simple volume ratios 2 volumes of hydrogen + 1 volume of oxygen 2 volumes of water vapor H? H O H + O H O Joseph Louis Gay-Lussac ( ) 19
20 1811 A. Avogadro deduced from Dalton s atomic theory and from Gay-Lussac s law: At the same temperature and pressure, the same volumes of different gases contain the same number of particles. Gases are diatomic molecules H 2, N 2, O 2 Avogadro s Hypothesis Not accepted till 1858, Cannizzaro Water till that time was considered as OH, M(O) = 8 after 1858 H 2 O, M(O) = 16 Amadeo Avogadro ( ) 20
21 Law of Constant Volumes At the same temperature and pressure, the same volumes of different gases contain the same number of particles. Gases are diatomic molecules H 2 H 2 O + O 2 H 2 H 2 O 21
22 Avogadro s Law 1811 The same volumes of different gases contain the same number of particles (at const. p, T) V = n const. Volume of 1 mole of gas = 22.4 liter at 0 C and Pa (STP) V M = 22.4 l mol 1 molar volume of ideal gas V/n = const. (at 0 C and Pa (1 bar) V M = l mol 1 ) Pressure depends on the number of molecules, temperature and volume p V = f (n, T) 22
23 Ideal Gas Composed of small particles (atoms, molecules) that are in constant motion along linear trajectories in random directions with high velocities. Dimensions of particles are very small in comparison to their distances They do not exert attractive or repulsive forces. Collisions are elastic, no loss of energy. Kinetic energy of a particle depends on temperature (but not on pressure). E kin = ½ m v 2 E kin = m v = 3 2 k B T 23
24 V = n const. V/n = const. 1 mole of a Real Gas V m = l mol 1 At standard temperature and pressure (STP) p = kpa = 1 atm = 760 torr t = 0 C 24
25 Ideal Gas Equation Ideal gas Molecular volume is zero (very small in comparison to gas volume) No intermolecular forces p V = n R T n = amount of substance V = (n R T) / p p= (nrt)/v n/v = p / RT R = gas constant R = J K 1 mol 1 25
26 Ideal Gas Equation p V = n R T p = const. isobaric V = const. isochoric T = const. isothermal 26
27 Calculation of Gas Density and M r p V = n R T = (m/m) R T r = m/v = p M / R T M = r RT / p = r V m Gas Density Gas molar mass V m = R T / p 27
28 Partial Pressure, p i p i = Pressure of a component of a mixture if it was alone in a given volume. Molární zlomek x i = n i / Σ n i Σ x i = 1 x i = n i Σn i Pressure of gas trapped above a liquid p = p(gas) + vapor pressure 28
29 Dalton s Law of Partial Pressures p tot = p 1 + p 2 + p p n = Σ p i p(air) = p(o 2 ) + p(n 2 ) + p(ar) + p(co 2 ) + p(other) Partial Pressure Pressure of a component of a mixture if it was alone in the given volume. P He = x He P tot P Ne = x Ne P tot P tot = P He + P Ne 29
30 Nonideal (Real) Gas Nonideal gas behavior approaches ideal gas behavior at high temp and low pressure Z = compressibility factor Z = Z = 1 for ideal gas At low temp and high press, intermolecular forces and molecular volume have to be considered 30
31 Z = compressibility factor Nonideal (Real) Gas Z > 1 Molar volume of nonideal gas is bigger than ideal gas Repulsive intermolecular interactions prevail Z < 1 Molar volume of nonideal gas is smaller than ideal gas Attractive intermolecular interactions prevail 31
32 Van der Waals State Equation of Real Gases p a + V 2 m ( V b) = RT m V m = molar volume of gas b = volume of molecules of gas (must be substracted) a = intermolecular attraction (must be added to p) J. D. van der Waals ( ) NP in Chemistry
33 Van der Waals State Equation of Real Gases 2 an P + )( V nb) = V ( 2 nrt P = nrt ( V nb) an ( V 2 2 ) Gas a (l 2 bar mol -2 ) b (l mol -1 ) Helium Hydrogen Nitrogen Oxygen Benzene
34 34
35 Liquification of Gases Condensation requires vdw forces Low T, high p, decrease of E kin, close approach of molecules Ideal gas cannot be liquified Critical temperature = above critical temperature, gas cannot be liquified by any high pressure 35
36 Joule-Thompson Effect Joule-Thompson effect = change of temperature during adiabatic expansion of compressed gas thru an orifice (pressure drops dp < 0) μ = dt/dp Joule-Thompson coeficient μ> 0 cooling (dt < 0) breaking of vdw bonds, required energy is taken from E kin, T drops. Below J-T inversion temp. O 2, N 2, NH 3, CO 2, freons N 2 (348 C) O 2 (491 C) μ = 0 ideal gas, real gas at J-T inversion temp. 36
37 Joule-Thompson Effect Joule-Thompson effect = change of temperature during adiabatic expansion of compressed gas thru an orifice (pressure drops dp < 0) μ = dt/dp Joule-Thompson coeficient μ < 0 heating (dt > 0) Above J-T inversion temp. H 2, He, Ne. He (-222 C) There are repulsive interactions in compressed gas that are removed upon expansion, energy is released = heating 37
38 Liquification of Gases Heat exchanger Jet Compresor 38
39 Kinetic-Molecular Theory of Gases 1738 Daniel Bernoulli ( ) Atoms and molecules are in perpetual motion, temperature is a measure of intensity of this motion Statistical mechanics, Clausius, Maxwell, Boltzmann Average velocity of molecules of H 2 at 0 C <v> = m s 1 = 6624 km h 1 39
40 Kinetic-Molecular Theory of Gases Average kinetic energy of gas molecules E kin = 1 / 2 m<v 2 > m = gas molecule mass <v> = Root-mean-square speed Average kinetic energy of all gases at a given temperature is the same E kin = 3 2 k B T 40
41 Maxwell-Boltzmann Distribution kt Speed ~ m = RT M 41
42 Maxwell-Boltzmann Distribution dn = 4πN (m / 2 π kt) 3 /2 exp( ½mv 2 /kt) v 2 dv Most probable speed Average speed Root-mean-square speed v mp = (2kT / m) ½ v av = (8kT / π m) ½ v rms = (3kT / m) ½ kt Speed ~ m = RT M 42
43 Kinetic-Molecular Theory of Gases Number of molecules Area under curves is constant because total number of molecules is constant Speed, m s 1 No molecule has zero speed Maximum speed The higher the speed, the less molecules 43
44 44
45 Maxwell-Boltzmann Distribution kt Speed ~ m = RT M 45
46 Mean free path, l, average distance between collisions Depends on p and T Diffusion l = const T/ p = const /n π (2r) 2 n = number of particles in m 3 r = molecular radius l = Å At normal p,t Viscosity, thermal conductivity 46
47 Effusion Graham s Law v 1 /v 2 = (ρ 2 /ρ 1 ) ½ = (M 2 /M 1 ) ½ 47
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