Gases- Part 2Gas StoichiometryDalton’s Partial PressureKinetic Molecular TheoryEffusion and DiffusionReal Gases

Gas Stoichiometry

• Standard Temperature and Pressure • 1.000 atm • 0°C (in equation 273.15K)

• Molar Volume = 22.4 L of an ideal gas at STP

Gas Stoichiometry Example

• What volume will 1.18 moles of O2 occupy at STP?

Gas Stoichiometry• Moles of Reactants to Moles of Products• Mass or volume (at STP) of reactants

Ideal Gas Law

• Moles of reactants

Stoichiometry

• Moles of products

Ideal Gas Law

• Mass or Volume (at STP) of products

**Balanced Chemical Equation First!**

Gas Stoichiometry Example

• 0.500 L of H2(g) are reacted with 0.600 L of O2(g) at STP. What volume will the H2O occupy at 1.00 atm and 350°C?

(Hint: Limiting Reactant Problem)

Gas Stoichiometry Practice

HCl(g) can be prepared by reaction of NaCl with H2SO4. What mass of NaCl is required to prepare enough HCl to fill a 340. mL cylinder to a pressure of 151 atm at 20.0°C?

Molar Mass of a Gas

Molar Mass of a Gas Example

• A gas at 34.0°C and 1.75 atm has a density of 3.40 g/L. Calculate the molar mass of the gas.

Molar Mass of a Gas Practice

• Air is a mixture of about 21.0% oxygen and 79.0% nitrogen (we’ll neglect the minor components and water vapor in this example). What is the density of air at 30.0°C and 1.00 atm?

Dalton’s Law of Partial Pressures

• For a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone.

PTOTAL = P1 + P2 + P3 + …

• The total pressure of the mixture and the Ideal Gas Law equation:

**How can this be so … if the gases are different?**

Dalton’s Partial Pressure Example

• A volume of 2.0L of He at 46°C and 1.2 atm pressure was added to a vessel that contained 4.5 L of N2 at STP. What is the total pressure and partial pressure of each gas at STP after the He is added?

Mole Fraction

• The ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture.

Dalton’s Law Example

• The mole fraction of argon in dry air is 0.00934. How many liters of air at STP will contain enough argon to fill a 35.4 L cylinder to a pressure of 150. atm at 20.°C?

Dalton’s Practice

• Assume the mole fraction of nitrogen in the air is 0.8902. Calculate the partial pressure of N2 in the air when the atmospheric pressure is 820 torr.

Kinetic Molecular Theory

• Recall as many components of the KMT.

• Again, KMT explains/provides the rationale behind our Gas Laws.

Kinetic Molecular Theory• 1. The volume of the individual particles of a gas can be

assumed to be negligible.• 2. The particles are in constant motion. The collisions of

the particles with the walls of the container are the cause of the pressure exerted by the gas.• 3. The particles are assumed to exert no forces on each

other.• 4. The average kinetic energy of a collection of gas

particles is assumed to be directly proportional to the Kelvin temperature of the gas.

Temperature is a measure of Kinetic Energy.

Application of KMT and Gas Laws – Write it out!!• Using the postulates of the Kinetic Molecular Theory

describe/explain the following laws:• Boyle’s Law• Gay-Lussac’s Law• Charles’s Law• Avogadro’s Law• Ideal Gas Law

Temperature and Kinetic Energy

• Temperature is an index of the random motions of the particles in a gas.

**The idea that temperature is a measure of the average kinetic energy of a gas is of critical importance!!**

Root Mean Square Velocity

• The average velocity of gas particles.

• R = 8.314 J/K mol• T = temperature in Kelvin• M = mass of a mole of the gas in kilograms (kg/mol)**Watch Units!**• Reality is that particles have a LARGE range of velocities.

Root Mean Square Velocity

• Calculate the root mean square velocity for the atoms in a sample of oxygen gas at • 0.00°C (3 sig figs)• 300. °C

**Increase in rms velocity agrees with prediction – KMT**

Effusion and Diffusion• Diffusion• The mixing of gases• The rate of diffusion is the rate of the mixing of gases. • Smell diffuses throughout a room (particle collision)

• Effusion• The passage of a gas through a tiny orifice into an evacuated

chamber.• The rate of effusion is the speed at which the gas is transferred

into the chamber.• Thomas Graham (1805-1869) Scottish chemist

• Rate of effusion of a gas is inversely proportional to the square root of the mass of its particles.

• The higher the molar mass, the slower the rate of effusion through a small orifice.

Graham’s Law of Effusion

• The higher the molar mass, the slower the rate of effusion through a small orifice.

• M = molar mass (not concentration/molarity here)

Graham’s Law of Effusion

• How many times faster than NO2 would He gas effuse?

• **Does the answer make sense?**

Graham’s Law of Effusion

• Calculate the rate of effusion of PH3 molecules through a small opening if NH3 molecules pass through the same opening at a rate of 8.02cm3/s. Assume the same temperature and equal partial pressures for the two gases.

• Which gas would effuse faster, Ne or CO2? How much faster?

Diffusion

• Even though gases travel very rapidly, mixing is a relatively slow process.

• The distance traveled, during mixing, is inversely related to the molar mass of the gas. (similar to Graham’s Law)

Real Gases

• Ideal gas behavior can be thought of as the behavior approached by real gases under certain conditions.

• High temperature• Low pressure

• The “ideal gas” model fails under high pressure and low temperature. Van Der Waal’s Equation used to account for the finite volume of a gas and the interaction between gas particles (negated by KMT).

Chemistry in the Atmosphere• Key ideas• The atmosphere is composed of 78% N2, 21% O2 0.9% Ar, and

0.03% CO2 along with trace gases.• The composition of the atmosphere varies as a function of distance

from the earth’s surface. • Heavier molecules tend to be near the surface due to gravity.

• Upper atmospheric chemistry is largely affected by ultraviolet, x-ray, and cosmic radiation emanating from space. The ozone layer is especially reactive to ultraviolet radiation.

• Manufacturing and other processes of our modern society affect the chemistry of our atmosphere. Air pollution is a direct result of such processes.

**Higher ozone levels that are characteristic of smog cause lung and eye irritation and can be very dangerous for people with asthma, emphysema, and other respiratory conditions. Ozone can also react with other pollutants.**

Individual Practice Problems

• Starting on page 220• 51, 53, 55, 57, 63, 65, 66, 69, 75, 85, 86, 88

For next time…

• Molar Volume of a Gas Pre-lab!!

• Continue practice problems and review of Ch. 4 and 5 • Check the student topic check list and begin

preparing for Exam 2