Chapter Eleven Gases

Section 11.1 Properties of Gases

Properties of Gases Gas vs. Vapor Vapor is a term referring to the gas state of something that is liquid/solid at room temperature Few elements exist as gases at room temp. H,N,O,F, Cl, and the noble gases Molecular compounds with LOW molar masses tend to be gases at room temp (i.e. CO2, HCl)

Gases versus Solids & Liquids Gases expand and take shape of container No definite shape or volume Gases are compressible Gases have MUCH lower densities Density is also HIGHLY variable depending on temperature and pressure Gases form homogeneous mixtures with one another

Section 11.2 Kinetic Molecular Theory of Gases

Kinetic Molecular Theory (KMT) Four basic assumptions: Gases are tiny particles with large amounts of space in between (volume of gas particle is negligible). Gases in constant, random, straight-line motion (all collisions perfectly elastic) No attractive/repulsive forces on each other Average kinetic energy is proportional to the absolute temperature

Other Properties of Gases As molecular weight (mass) increases, speed decreases of gas particles Diffusion: mixing of gases as the result of random motion and collisions Effusion: escape of gas molecules from container to a region of vacuum

Section 11.3 Gas Pressure

Gas Pressure Pressure = Force/Area SI unit for force is Newtons (N) SI unit for pressure is pascals (Pa) Gases exert pressure on everything they touch Internal pressure & External pressure Atmospheric Pressure: pressure exerted by Earth’s atmosphere

Units & Conversions Atmospheres (atm) Pascals (Pa) and kilopascals (kPa) Bar (bar) Millimeters of Mercury (mmHg) Torr (torr) Pounds per Square Inch (psi) 1.00 atm = 101,325 Pa = 101.325 kPa = 1.01325 bar = 760. mmHg = 760. torr = 14.7 psi

Measuring Pressure Barometer: measures atmospheric pressure

Measuring Pressure Manometer: measures pressures other than atmospheric pressure

Examples of Manometers

Group Quiz #1 Convert the following: 12.4 psi  kPa 509 bar  atm 98,320. Pa  mmHg Assuming atmospheric pressure is 765.0 mmHg, what is the pressure of the gas in the bulb?

Section 11.4 The Gas Laws

The Gas Laws Gases and their physical state can be described with the following four parameters: Pressure (P) Volume (V) Temperature (T) Number of moles (n)

Boyle’s Law Relationship between Pressure & Volume Constant # of moles and temperature P1V1 = P2V2 = k (k is a constant) As pressure increases, volume decreases (inverse) If pressure doubles, volume was halved Ex: What will be the pressure of 21 mL of a gas if its pressure was 731 mmHg at 43 mL?

Charles’ s Law Relationship between Volume & Temperature Constant pressure & number of moles V1 = V2 As temperature increases, so T1 T2 does volume (direct relationship) Based on absolute (Kelvin) scale, so temp MUST be in Kelvin! ( °C K, add 273.15) Ex: If a 1.45 L balloon is cooled from 25.0°C to 15.0°C, what is the new volume?

Avogadro’s Law Relationship between Volume & # of moles Constant pressure and temperature V1 = V2 As # of moles increases, volume n1 n2 increases (direct relationship) Avogadro proposed equal volumes of different gases contained the same amount of particles at the same temperature and pressure 1 mole = 22.41 L @ STP (1 atm and 0°C)

Combined Gas Law Relationship between Volume, Pressure, Temperature, & # of moles Combination of Boyle’s & Charles’s Law P1V1 = P2V2 *Remember, temp must n1 T1 n2T2 be in Kelvin! Ex: If a child releases a 6.25-L balloon where the temperature is 28.50°C and the air pressure is 757.2 mmHg, what will be the volume of the balloon when it has risen to a height where the temperature is -34.35°C and the air pressure is 366.4 mmHg?

Group Quiz #2 At what temperature will a gas sample occupy 100.0 L if it originally occupies 76.1 L at 89.5°C? Assume constant P. If the pressure of a container is 1.50 atm at a volume off 37.3 mL, what will be the new volume if the pressure is changed to 1.18 atm?

Section 11.5 The Ideal Gas Equation

The Ideal Gas Law Includes all 4 gas variables PV = nRT P = atm or kPa V = liters n = moles T = Kelvin R = 0.08206 atm*L/mol*K or 8.314 kPa*L/mol*K

What is an Ideal Gas? One in which a sample of gas’s pressure-volume-temperature behavior is predicted by the ideal gas law equation Most of the time the differences between an ideal gas and a real gas are negligible

Example What is the volume of 5.12 moles of an ideal gas at 32.0°C and 749.5 torr?

Ideal Gas and Stoichiometry Consider the reaction of zinc with hydrochloric acid to produce zinc chloride and hydrogen gas. How many grams of zinc are necessary to produce 11.5 L of hydrogen gas at 825 torr and 25.0°C?

Using Ideal Gas Law for Density Ex: Calculate the density of CO2 at room temperature (25.0°C) and 1.00 atm.

Section 11.6 Real Gases

Real versus Ideal Gases Gas Laws and KMT assume ideal gas behavior Most gases exhibit ideal (or nearly ideal) behavior under ordinary conditions Conditions of high pressure and low temperature constitute a deviation of a real gas from an ideal gas

Section 11.7 Gas Mixtures

Dalton’s Law of Partial Pressures In a mixture of gases, each gas exerts a pressure as though it were by itself Each gas exerts its own pressure, regardless of the presence of any other gas Partial Pressure Ptot = P1 + P2 + P3 + … Pn

Mole Fractions Another way of showing the quantity of a gas in a mixture of gases Mole Fraction = moles of gas X total # of moles Remember: Mole fractions less than 1 Sum of mole fractions equals 1 Mole fractions are unitless

Mole Fractions Since n and P are proportional at a set T and V, mole fractions can also be found be dividing partial pressure by total pressure Mole Fraction = Partial Pressure Total Pressure

Dalton’s Law of Partial Pressures Ex: Three gases occupy a volume of 1.55 L. Gas 1 has a pressure of 1.30 atm; Gas 2 has a pressure of 3.10 atm; Gas 3 has a pressure of 2.70 atm. What is the total pressure? What is the mole fraction for each gas? Ex: A 1.00 L vessel contains 0.215 moles of N2 gas and 0.0118 moles of H2 gas at 25.5°C. Determine the partial pressure of each gas and the total pressure in the vessel.

Practice What is the mole fraction of CO2 in a mixture of 0.756 moles of N2, 0.189 moles of O2, and 0.0132 moles of CO2? If the total pressure in the container is 212.4 atm, what is the pressure of CO2?

Lab Example

Lab Example In the lab, you react aluminum with excess hydrochloric acid. As the reaction goes to completion, you observe that 19.12 mL of gas has been collected in the buret. The leveling bulb has been held in such a way where the liquid level is equal to that of the buret. If the pressure in the room is 0.9894 atm, the temperature is 22.5°C, and the vapor pressure of water at this temperature is 0.02687 atm, what is the number of moles produced of hydrogen gas alone?

Section 11.8 Reactions with Gaseous Reactants and Products

Gas Stoichiometry Stoichiometry where reactants and products are all gases Ex: How many liters of ammonia gas can be produced from 2.38 L of nitrogen gas reacted with excess hydrogen gas?

Gas Stoichiometry Recall: at a set temperature and pressure (typically STP), the volumes of 1 mole of any gas are equal. Use mole ratio Ex: N2(g) + 3H2(g)  2NH3(g) Must use ideal gas law if conditions are NOT at STP

Group Quiz #4 Determine the mass of NaN3 needed for an air bag to produce 100.0 L of N2 gas at 85.0°C and 1.00 atm according to the equation: 2NaN3(s)  2Na(s) + 3N2(g)