Unit 3 AP chem gases only

Assumptions of Kinetic Molecular Theory

1) gases made of particles with negative volumes 2) no intermolecular forces 3) continous, random motion, collisions 4) collisions are elastic 5) average KE is proportional to absolute temperature

standard pressure/atmospheric pressure

1.00 atm = 760 mm Hg = 760 torr = 101.3 kPa

An equimolar mixture of N2(g) and Ar(g) is kept inside a rigid container at a constant temperature of 300 K. The initial partial pressure of Ar in the mixture is 0.75atm. An additional amount of Ar was added to the container, enough to double the number of moles of Ar gas in the mixture. Assuming ideal behavior, what is the final pressure of the gas mixture after the addition of the Ar gas?

2.25atm, because doubling the number of moles of Ar doubles its partial pressure.

At 10 degrees C, 20g of oxygen exerts a pressure of 2.1 atm in a rigid,7.0 L cylinder. Assuming ideal behavior, of the temperature of the gas was raised to 40 degrees C, which statement indicates the new pressure?

2.3 atm, because the pressure P increased by the proportion 313/283

STP (standard temperature and pressure)

273 K and 1 atm

Effusion

A process by which gas particles pass through a tiny opening. They are going from a place with higher pressure to a place with lower pressure

A gas mixture at 0°C and 1.0atm contains 0.010mol of H2, 0.015mol of O2, and 0.025molof N2. Assuming ideal behavior, what is the partial pressure of hydrogen gas (H2) in the mixture?

About 0.20atm, because H2 comprises 20% of the total number of moles of gas.

A gaseous air‑fuel mixture in a sealed car engine cylinder has an initial volume of 600.mL at 1.0atm. To prepare for ignition of the fuel, a piston moves within the cylinder, reducing the volume of the air‑fuel mixture to 50.mL at constant temperature. Assuming ideal behavior, what is the new pressure of the air‑fuel mixture?

About 12atm, because the volume of the gas mixture decreased by a factor of 12.

van der walls equation

An equation of the state for gases that modifies the ideal gas equation to account for intrinsic molecular volume and intermolecular forces (P+n2a/v2)(V-nb)=nRT - need to know that a is adding pressure since it is lower in real gases - b is used to subtract volume since it is higher in real gases

Why do not all real gases behave ideally at high pressure and low temperatures

As pressure is increased the particles are pushed closer together which results in the volume of the gas molecules becoming significant. As the temperature decreases the particles are moving slower and the intermolecular attractions become significant.

How does temperature affect diffusion

As temperature increases, the rate of diffusion increases since the particles are moving faster

A student is doing experiments with CO2(g). Originally, a sample of gas is in a rigid container at 299K and 0.70 atm. The student increases the temperature of the CO2(g) in the container to 425K. Describe the effect of raising the temperature on the motion of the CO2(g) molecules. Calculate the pressure of the CO2(g) in the container at 425K. In terms of the , briefly explain why the pressure of the CO2(g) in the container changes as it is heated to 425K.

As you increase the temperature, or the average , of the CO2(g) molecules, the speed of the molecules increases as well. 0.70 atm/299 K = P2/425 K, which gets us 0.99 atm. Faster-moving gas particles collide more frequently and forcefully with the walls of the container, increasing the overall pressure.

The gases CO2(g) and NH3(g) can be liquefied at 20°C by compressing them to sufficiently high pressures. A student claims that NH3(g) can be liquefied at a lower pressure than CO2(g) can be liquefied. Which of the following is the best justification for this claim?

CO2 is a nonpolar molecule that has London dispersion intermolecular forces that are weaker than the dipole-dipole and London dispersion forces between the polar NH3 molecules.

Gay-Lussac's Law explained

Describes constant relationship between temperature and pressure under constant volume

Charles' Law explained

Describes the direct relationship between volume and temperature under constant pressure (when temperature increase, collisions increase, so volume also has to increase to keep pressure constant)

Avogadro's Law explanation

Direct relationship between volume and moles of gas in a sample

If the volume of a gas is less than you would expect based on the ideal gas law, the best explanation for this is because the...

IMFs are stronger

Formula for kinetic energy

KE=1/2mv^2

Conditions causing gas to deviate from ideal behavior

Low temperature high pressure

Density of gases

MM = dRT/P

NO2 and CO2 have a similar molecular weight. Which gas would you predict to deviate from as ideal gas? Justify your selection.

NO2, it is polar due to the lone pairs on N, therefore the dipole dipole IMF is stranger that CO2

Combined Gas Law

P1V1/T1=P2V2/T2

Boyle's Law Equation

P1V1=P2V2

Dalton's law mole fractions

Pa = mols a(Ptotal)

If the volume of a gas is greater than you would expect based on the ideal gas law, the best explanation for this is because the...

Particle volume is great

Graham's Law of Effusion Equation

Rate A/Rate B = square root of molar mass B/molar mass A

The graph above shows how a particular real gas deviates from ideal behavior at very high pressures. Based on this information, which of the following is most likely the gas and gives the reason based on kinetic molecular theory?

SO2, because it has the largest molecular volume.

Rate of effusion equation

Square root of mm gas A / mm gas B

The student measures the actual pressure of CO2(g) in the container at 425K and observes that it is less than the pressure predicted by the ideal gas law. Explain this observation.

The attractive forces between CO2(g) molecules result in a pressure that is lower than that predicted by the ideal gas law. Since the particles are attracted to each other, they aren't colliding with the walls of the container as often as ideal gases with no attractive forces would.

A 1 L sample of helium gas at 25C and 1 atm is combined with a 1 L sample of neon gas at 25C and 1 atm. The temperature is kept constant. Which of the following statements about combining the gases is correct?

The average kinetic energy of the helium atoms and neon atoms do not change when the gases are combined.

How does molecule size affect diffusion

The bigger the molecules, the slower the diffusion. This is because these molecules contain more mass and make slower movements.

The diagrams above represent two samples of gas in containers of equal volume at . Which of the following correctly compares the two samples in terms of their deviation from ideal gas behavior and explains why?(sample 1 has 4 atoms and sample 2 has 9 atoms)

The gas in sample 2 would deviate more from ideal behavior because the atoms are closer together, leading to an increase in intermolecular attractions.

Diffusion

The mixing of gases

The graph above shows the distribution of molecular speeds for four different gases at the same temperature. What property of the different gases can be correctly ranked using information from the graph, and why?

The molecular masses of the gases, because the gas molecules have the same average kinetic energy and mass can be calculated using the equation KEavg=12mv2.

Ar(g) deviates more from ideal behavior at extremely high pressures than Ne(g) does. Which of the following is one reason for this difference?

The particle volume of Ar is greater than that of Ne

Two sealed, rigid 5.0L containers each contain a gas at the same temperature but at a different pressure, as shown above. Also shown are the results of transferring the entire contents of container 1 to container 2. No gases escape during the transfer. Assuming ideal behavior, which statement is correct regarding the total pressure of the gases after they are combined?

The total pressure of the gases in the mixture is the sum of the initial pressures of oxygen gas and nitrogen gas because pressure only depends on the total amount of gas when volume and temperature are held constant.

Charles law equation

V1/T1=V2/T2

Avogadro's Law Equation

V1/n1 = V2/n2

Pressure vs ideal gases

When the gas particles are close together due to a large number of particles, this can cause more attractive forces. The pressure of real gases is usually lower than the pressure of ideal gases due to attractive forces. When particles are attracted to each other, IMFs become significant and the particles aren't hitting the walls of the container as often.

ideal gas behavior

high temperature and low pressure

Dalton's Law of Partial Pressures

states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture

Rules for effusion

temperature increases the rate of effusion while a higher mass decreases the rate of effusion

Maxwells equation

vrms = sqrt (v^2) = sqrt (3RT/M) R - gas constant = 8.314 J/ mol K = 8.314 kg m^2/s^2 mol K T - absolute temp in Kelvins M - Molar mass (kg/mol)

Temperature

A measure of the average energy of motion of the particles of a substance.

Boyle's Law explained

A principle that describes the inverse relationship between the pressure and volume of a gas at constant temperature

Equimolar samples of CH4(g) and C2H6(g) are in identical containers at the same temperature. The C2H6(g) deviates much more from ideal behavior than the CH4(g) does. Which of the following best helps explain this deviation?

C2H6 molecules have a larger, more polarizable electron cloud than CH4 molecules do.

Diagram 1 above shows equimolar samples of two gases inside a container fitted with a removable barrier placed so that each gas occupies the same volume. The barrier is carefully removed as the temperature is held constant. Diagram 2 above shows the gases soon after the barrier is removed. Which statement describes the changes to the initial pressure of each gas and the final partial pressure of each gas in the mixture and also indicates the final total pressure?

The partial pressure of each gas in the mixture is half its initial pressure; the final total pressure is half the sum of the initial pressures of the two gases.

pressure

the number of times particles hit the walls of the container