Gases Chapter 11

Gases Gases are made of _______________________. These particles move in _______________ at varying speeds until they collide with another particle or a barrier.

Gases A gas has ______________________ The Gas Laws pertain to an _____________, an imaginary gas that serves as a model for gas behaviors.

Is it real or ideal? Ideal Conditions/Ideal Gas - imaginary gas that fits the all assumptions of the kinetic-molecular theory Real Gas – particles have size and intermolecular attraction to other particles, so do not conform to gas laws under very _____ pressure or very _____ temperatures.

Kinetic Theory based on idea that particles are always in ________ At _______ temperature, all gases have the ________ kinetic energy (KE) But gases do not all have same ______

Rearrange to fastest to slowest KE = __________ m = mass v = velocity Rearrange to fastest to slowest O2 CO2 H2 N2 _____________________ Why?? Well who’s faster, a 350 lb football player or a 120 lb runner?

ALL gases at the SAME temperature have the same average KE If the KE = 20, note the speed for the gases . H2 KE = ½ m v2 20 = ½ 2g v2 20 = 1 v2 20/1 = v2 20 = v2 v2 = /20 v = 4.472 Ar KE = ½ m v2 20 = ½ 40g v2 20 = 20 v2 20/20 = v2 1 = v2 v2 = /1 v = 1

Kinetic molecular theory assumptions: 1.   Gases consist of large # of ___ molecules that are ___ apart relative to their size. Thus they have ___________ and lots of empty space between them.

Kinetic molecular theory assumptions: 2.    Lots of _________________ - don’t stick together after collisions. No net loss of _____________ 3.    NO _________ or __________ for each other (ideal gases do not condense to a liquid or a solid as no attraction)

Kinetic molecular theory assumptions: (cont) 4.    ____________ = energy of motion Gas particles are in constant, rapid motion that is random. Therefore, passes “kinetic energy” = energy of motion. 5.    Average kinetic energy depends on the _______________ of the gas.

The Nature of Gases for Ideal Gases 1. Expand to fit container - no definite ___________ or __________ 2. ______ - slide past each other – fluid as attraction forces not significant 3. ____________ - particles far apart 4. _______________ – particles can be pressed close together.

Nature cont. Rate of Diffusion depends on: a. ____________ 5. ___________ - random mixing caused by random motion Rate of Diffusion depends on: a.   ____________ b.   ______________ c. ______________ between particles O2 vs H2O

Nature cont. Smaller molecule → __________________________ 6. _____________ - gas particles under pressure pass thru tiny openings Smaller molecule → __________________________ (Does a balloon stay full forever? Why not?)

To describe what a gas is doing, 4 measurable qualitites are used           ________________          ________________        ________________         ________________

Pressure Pressure = _______ (Force is in _____________ Force is caused by gas molecules hitting wall - their collisions Air around us exerts a pressure on us. What would happen to your eyeballs if you were to go out into outer space?? Where is the most pressure - sea level, ocean bottom, mountain top?

Barometer Note air pushing down Vacuum has ________ inside, so is not pushing down

A simple manometer.

Standard Temperature & Pressure (STP) -- must be standard to compare things STP = ______________________________________________________ pressure = __________________________________________________________

Conversions Standard Pressure with different units 1 atmosphere (atm) =760 mm Hg (millimeter of mercury) =760 torr =14.7 lb/in2=76 cm Hg =29.9 in Hg = 1013 millibar (mbar) = 101,325 Pascal = 101.325 KPa (SI unit) Pascal - the SI unit = 1N/ m2

These all equal each other, so can use as conversion factors Given 700 mm Hg, how many atm?

Covert the following 860 mm Hg = ______ atm 720 torr = ______ in Hg 0.89 atm = ______ mm Hg 1.2 atm = ______ in Hg

The Gas Laws ______________ -the volume of a fixed mass of gas varies ___________ with the pressure at constant temperature. So as P __________  V ________ ! http://preparatorychemistry.com/Bishop_Boyles_frames.htm

As pressure increases, the volume decreases Temperature and number of moles are constant

The effects of decreasing the volume of a sample of gas at constant temperature. Molecules hit more often, so Pressure ____________

_____________________ Boyle’s Law P1V1 = P2V2 initial = final _____________________

Boyle’s law Ex. A sample of O2 has a volume of 100 ml and a pressure of 200 torr. What will it’s volume be if the pressure is increased to 300 torr (T is kept constant) First think – what is the relationship? P ___________, then V ___________

Boyle’s law P1 x V1 = P2 x V2 V2 =

Plotting Boyle's data from Table 5. 1 Plotting Boyle's data from Table 5.1. (a) A plot of P versus V shows that the volume doubles as the pressure is halved. (b) A plot of V versus 1/P gives a straight line. The slope of this line equals the value of the constant k.

A plot of PV versus P for several gases at pressures below 1 atm.

deals with ________ and _____________ (Pressure constant). Charles Law deals with ________ and _____________ (Pressure constant). V and T (in Kevin) are ________________ As you _______________ the temperature - molecules MOVE __________ KE (Kinetic energy) _____________ as molecules jump around more http://preparatorychemistry.com/Bishop_Charles_frames.htm

Charles Law Found gases had “zero” volume at -273 oC Temperature MUST be in ____________ Found gases had “zero” volume at -273 oC So named this _________________, which equals 0 Kelvin = 0 K   K = 273 + ____ oC oC = ___ K - 273

Charles Law V1 = T1 _________________ V2 T2 Can rewrite: V1T2 = V2T1 (Watch 1's & 2's)

The effects of increasing the temperature of a sample of gas at constant pressure. Molecules are moving ____________ so they hit the container harder, thus size must _________ to keep P same

Charles Ex. A sample of neon gas occupies a volume of 752 ml at 25 oC. What volume will it occupy at 50 oC. Pressure is constant. Think first - as Temperature increases, Volume will ____________

Charles MUST change oC to Kelvin Given : V1 = 752 ml T1 = 25  oC + 273 = 298 K V2 = ? T2 = 50 oC + 273 = 323 K

Charles V1 x T2 = V2 x T1 V2 =

Gay- Lussac’s Law the pressure of a fixed mass of gas at constant volume varies ___________ with the Kelvin temperature P1 = T1 P2 T2 _________________ Or P1 x T2 = P2 x T1 http://preparatorychemistry.com/Bishop_Gay_Lussac_frames.htm

The effects of increasing the temperature of a sample of gas at constant volume. Molecules _____ _______, so pressure __________

Ex. A gas in an aerosol can is at a pressure of 3. 00 atm at 25 ° C Ex. A gas in an aerosol can is at a pressure of 3.00 atm at 25 ° C. Directions warn the user not to keep the can in a place where temperature exceeds 52 ° C. What would the pressure in the can be at 52 oC? Given P1 = 3.00 atm P2 = ? T1 = 25 oC or 298 K T2 = 52 oC or 325 K

Combined Gas Law P1V1 = P2V2 T1 T2 or P1V1 T2 = P2V2 T1 Expresses the relationship between ________, _______, and __________ of a fixed amount of gas. P1V1 = P2V2 T1 T2 or P1V1 T2 = P2V2 T1

Ex. A helium balloon has a volume of 50. 0 L at 25 ° C and 1. 08 atm Ex. A helium balloon has a volume of 50.0 L at 25 ° C and 1.08 atm. What volume will it have at 10 oC and 0.855 atm? Given: V1 = 50.0L V2 = ? P1 = 1.08 atm P2 = 0.855 atm V2 = T1 = 25 ° C = 298 K T2 = 10 ° C = 283 K

Ideal Gas Equation _____________ - equal volumes of gases at the same temperature and pressure contain equal numbers of molecules http://preparatorychemistry.com/Bishop_Avogadro_frames.htm Thus, the volume occupied by one mole of gas at ______ = standard molar volume of a gas = ________

Ideal Gas Equation Ex. A chemical reaction produces 0.0680 mol of O2. What volume will it occupy at STP? Given: 0.0680 mol O2 1 mol = 22.4 L at STP

Ex. A chemical reaction produces 98. 0 mL of SO2 at STP Ex. A chemical reaction produces 98.0 mL of SO2 at STP. What mass in grams was produced? Given: Volume SO2 at STP 98.0 mL = 0.098 L 0.098 L SO2 1 mol = .004375 mol SO2 22.4L .004375 mol SO2 64.07 g = 0.28 g SO2 1 mole OR 0.098 L SO2 1 mole 64.07 g = 0.28g SO2 22.4 L 1 mole

the constant, R, is the ideal gas constant has a value of .0821 L atm Ideal Gas Law the mathematical relationship of pressure, volume, temperature, and the number of moles of gas. ___________________ pressure x volume = # of moles x constant x temperature the constant, R, is the ideal gas constant has a value of .0821 L atm mole K WATCH units !!

The effects of increasing the number of moles of gas particles at constant temperature and pressure.

Plots of PV/nRT versus P for several gases (200 K).

Plots of PV/nRT versus P for nitrogen gas at three temperatures.

Ex. What is the pressure in atm exerted by 0 Ex. What is the pressure in atm exerted by 0.50 mol sample of N2 in a 10.0L container at 298 K? Given: V of N2 = 10.0L n of N2 = 0.50 mole T = 298 K P = ? PV = nRT therefore, P = nRT V

Combining the ideal gas law with density - to find molar mass or gas density. Ex. What is the density of ammonia gas at 63 oC and 705 mm Hg? Given: P = 705 mmHg V = ? n = assume 1 mole R = 0.0821 L atm /mole K T = 63 oC

Combining the ideal gas law with density - to find molar mass or gas density. V = 1mole(0.0821 L atm / mole K) ( 336 K) .928atm = D = mass mass of 1 mol NH3 = 17g volume (from Periodic table) D = 17g 29.73L =

Gas Stoichiometry Stoichiometry – the mass relationship between reactants and products in a chemical reaction. In general, follow these rules: a. Use stoichiometry to do mole or mass conversions (or if STP, 1 mole = 22.4 L). b. Use the Ideal Gas Law (PV=nRT) to convert between moles and volume.

Gas Stoichiometry Practice CaCO3 → CaO + CO2 How many grams of calcium carbonate must be decomposed to produce 5.00 L of carbon dioxide at STP? Do this two ways: 1st way: Use PV = nRT first, solve for n, then stoichiometry. 2nd way: Use stoich & 1 mole = 22.4 liters

Gas Stoichiometry Practice Calculate the volume of H2 gas that can be obtained under laboratory conditions of temperature and pressure of 25 C and 0.900 atm when 5.00 g of sodium is reacted with water: 2 Na + H2O  H2 + 2 NaOH

Gas Mixtures and Partial Pressures Many gases are a mixture of gases; air being a prime example. We define the ______________________ of a gas which is the pressure a component of a gas mixture would have if it were all by itself in the same container.

Dalton’s Law John Dalton (atomic theory) extended the gas laws simply as: Ptotal = P1 + P2 + P3 + … This simply says that the sum of the ____________________ of the gases will add up to the total pressure. This way we can treat a mixture of gases just like a pure gas.

Partial Pressure Practice What would be the final pressure of the mixture of gases for the processes depicted in each of the following illustration?

Effusion and Diffusion __________________________ - the rates of effusion of gases at the same temperature and pressure are inversely proportional to the square roots of their molar masses. Rate of effusion A = Mb Rate of effusion B Ma

The effusion of a gas into an evacuated chamber.

Ex. Compare the rates of effusion of H2 and O2 at same temperature and pressure. Given: H2 mol weight 2 O2 mol weight 32 Rate of effusion H2 = √ Mass O2 Rate of effusion O2 √ Mass H2

Important Gas relationships As volume increases, pressure ____________ at constant temperature As temperature increases, pressure _________ at constant volume As temperature increases, volume _________ at constant pressure Standard Temperature and Pressure - to compare gases use (standard temperature & Pressure _____) Standard temperature = _____ C Standard Pressure = _____ mm Hg = ___ atm average barometric pressure at sea level