2 5.1 Gases Temperature vs. Intermolecular attraction Atomic Gases Noble Gases, H2, N2, O2, F2, Cl2Molecular GasesUsually light molecules with weak attraction forcesEg: HCl, CO2, NH3, H2S, NO2Ionic CompoundsStrong forces; not normally gases
3 5.2 Pressure Force per unit area P = F/A N/m2 unit defined as Pascal (Pa)Standard air pressure = kPaAlso called 1 atmosphere (atm)Measured by unequal mercury levelsManometers and barometersCommon unit called mmHg (or Torr)Standard air pressure = 760 mmHg
4 5.2 Dimensional Analysis Convert 75.0 kPa to mmHg 75.0 kPa mmHg x = 563 mmHg kPaTry this oneConvert 1.25 atm to kPa
5 5.3 Boyle’s Law Pressure is inversely proportional to volume Hold temperature and amount of gas constantV a 1/PV = k x (1/P)PV = kBest used with changing conditionsP1V1 = P2V2
6 5.3 Boyle’s Law ProblemsA 175 mL sample of methane is stored at 125 kPa. What pressure is needed to compress the gas to a volume of 50.0 mL?P1V1 = P2V2(125 kPa) (175 mL) = P2 (50.0 mL)P2 = 438 kPaTry this oneA sample of argon occupies 476 mL at 650 Torr. Find the volume at 975 Torr.
7 5.3 Charles’ Law Also credited to Gay-Lussac Volume is directly proportional to temperatureHold pressure and amount of gas constantV a TV = kTLinear relationshipMust use Kelvins!V V = --- T T2
8 5.3 Charles’ Law ProblemsA 5.00 L helium balloon is heated from 20oC to 75oC. Find its new volume.V1/T1 = V2/T25.00 L V = K KV2 = 5.94 LTry this oneA 670 mL sample of chlorine is stored at 50oC. At what temperature will its volume be 450 mL?
9 5.3 More Gas Laws Another form of Charles’ Law Avogadro’s Law Pressure is directly proportional to temperatureP = kTP1/T1 = P2/T2Avogadro’s LawVolume is directly proportional to the amount of gas presentV a nVolume relationships in chemical reactions
10 5.3 Avogadro’s Law Problems How many liters of hydrogen are needed to completely react with 1 liter of oxygen?2 H2 + O2 2 H2O2 mol hydrogen react with 1 mol oxygenV a n, so….2 L hydrogen react with 1 L oxygenTry this oneHow many liters of ammonia are formed when 1 L of hydrogen reacts with excess nitrogen?
11 5.4 The Ideal Gas Equation Ideal Gas No intermolecular attraction forcesParticles have no volumeCombine Boyle’s, Charles, and Avogadro’s LawsPV = nRTSTP = 1 atm, 273 KMolar volume of a gas = L at STPR = atm L / mol K
12 5.4 Ideal Gas Equation Problems A sample of fluorine occupies 3.65 L at 45oC and 2.50 atm. How many moles of fluorine are present?PV = nRT(2.50 atm)(3.65 L) = n (0.0821)(318 K)n = molTry this oneA mol sample of propane occupies 2.15 L. If the temperature is 28oC, find the pressure.
13 5.4 Gas Density Since n = m/M…. PV = (m/M) RT MPV = mRT Divide by V to get density (m/V)MP = rRTGas density expressed in g/L
14 5.4 Gas Density Problems Find the density of nitrous oxide at STP. First, find molecular mass of N2OMP = rRT(44.0 g/mol)(1.00 atm) = r (0.0821)(273 K)r = 1.96 g/LTry this one….A gas is found to have a density of 2.54 g/L at 15oC and 1.50 atm. Find its molecular mass.
15 5.5 Gas Stoichiometry Mass-Mass problems (review) Volume-Volume problemsVolume is proportional to moles, so….Mol relationship from reaction can be used directlyNo conversions needed!
16 5.5 Volume-Volume Problem 2 H2 + O2 2 H2OIf 3.25 L of oxygen react, how many liters of water vapor are formed?3.25 L O L H2O x = 6.50 L H2O L O2Volume-Volume is just Avogadro’s Law!
17 5.5 Mass-Volume Problems Key step – get to moles! Mass conversion – use molecular massVolume conversion – use gas equationNeed to know temperature and pressure conditions
18 5.5 Mass-Volume Problems25.0 g of sodium react with excess water at STP. How many liters of hydrogen are produced?2 Na + 2 H2O 2 NaOH + H225.0 g Na 1 mol Na 1 mol H x x = mol g Na 2 mol NaNow use ideal gas equation to get volume
19 5.5 Mass-Volume Problems PV = nRT (1.00 atm) V = (0.543 mol)(0.0821)(273 K)V = 12.2 LTry this onePotassium chlorate decomposes into potassium chloride and oxygen gas. How many grams of KClO3 are needed to produce 5.00 L of oxygen at atm and 18oC?Hint: This one is backwards!
20 5.6 Dalton’s LawPartial pressure – the pressure of an individual gas in a mixture of gasesTotal pressure of a mixture equals the sum of the partial pressures of each gasPt = P1 + P2 + PPartial pressure is proportional to the mol fraction (X1 = n1 / nt)P1 = X1 Pt
21 5.6 Dalton’s Law2.00 mol He is mixed with 1.00 mol Ar. Find the partial pressure of each at 1.75 atm pressure.XHe = 2.00 mol / 3.00 mol = 0.667XAr = 1.00 mol / 3.00 mol = 0.333PHe = (0.667) (1.75 atm) = 1.17 atmPAr = (0.333) (1.75 atm) = atmTry this...Find the partial pressure of oxygen in air if it makes up 21% of the Earth’s atmosphere by volume. (Note: The volume gives you the mole ratio because of Avogadro’s law.)
22 5.7 Kinetic Molecular Theory Explains gas behavior in terms of molecular motionEnergyWork done by a moving objectMeasured in SI unit Joule (J)Kinetic energyEnergy due to motionK = ½ mv2KMT is a simplification of reality (ideal gas)
23 5.7 Kinetic Molecular Theory Gas molecules are separated by great distancesThey can be treated as “point masses”Gas molecules are in constant random motionFrequent elastic collisions (no energy lost)No attractive or repulsive forcesAverage K is proportional to Temperature
24 5.7 Distribution of Molecular Speeds Maxwell-Boltzmann DistributionMolecular speeds distributed around averagePeak velocity depends on temperature and on molec. massRoot Mean Square Speed_____ vrms = √3RT/MRate of diffusion
25 5.8 Deviations from Ideal Behavior We made approximations!Point massesNo intermolecular forcesThese approximations become bad at...High pressureLow temperatureLiquefactionvan der Waals Equation(P + an2/V2) (V – nb) = nRT