Gases

In the Van der Waals equation, the constant "a" corrects for intermolecular forces, and the constant "b" corrects for molecular size. In comparing the deviations of Xe and H2O from idea behavior, which has greater deviation due to: 1) intermolecular attractions? 2) molecular size?

1) "a" is larger for H2O because it has the ability to hydrogen bond, attracting other H2O particles more strongly than Xe particles can attract other Xe particles 2) "b" is larger for Xe because Xe atoms are larger than H2O molecules.

What are the two deviations from ideal behavior that real gases have?

1) Gas particles display intermolecular forces 2) Gas particles occupy a significant volume of their container (particle size is NOT negligibly small)

What are the two assumptions of the ideal gas law that can be invalid at low temperatures and high pressures?

1) The size of gas particles of particles is small compared to the space between them // At high pressures, gas molecules are crowded together, reducing the amount of empty space between them and causing them to occupy a significant amount of their container 2) There are no intermolecular forces between gas particles // At low temperatures, the effect of intermolecular forces is more prominent because particles have less kinetic energy to overcome them

How many Liters will 1 mole of ideal gas occupy at STP?

22.4 L

Standard temperature and pressure (STP)

273K/0°C and 1 atm

Avogadro's Law (at constant temperature and pressure, VOLUME and MOLES OF GAS are directly related)

According to Kinetic Molecular Theory: increasing the number of gas particles increases collisions with the surface of the container. The only way for pressure to remain constant is for the volume to increase so that the number of particles per unit volume (thus the number of collisions) remains constant. If the container is flexible (balloon), the walls will expand until the pressure of the gas inside matches the pressure of the gas outside. Thus, volume is directly proportional to the number of gas particles.

How does the Kinetic Molecular Theory explain the relationship between pressure and temperature in a rigid container (at constant volume)?

According to the model, kinetic energy of a gas increases with increasing temperature. As the particles move faster, they will collide with the walls of the container harder and more frequently, causing increased pressure.

Equation for calculating the gas density of an ideal gas:

D = m/V = MP/RT D= density (g/L) m= mass (g) V= volume (L) M = molar mass (g/mol) P = presssure (atm OR mmHg) R = constant T = temperature (K)

Charles's Law (at constant pressure and moles, VOLUME and TEMPERATURE are directly related)

Kinetic Molecular Theory states that an increase in temperature will increase the average kinetic energy of the gas particles. As the particles move faster, they will collide with the edges of their container more often. If the reaction is kept at constant pressure, then the particles must stay farther apart, and an increase in volume will compensate for the increase in particle collision with the surface of the container. The greater volume will spread the collisions apart over greater surface area so that the pressure (force per unit area) remains unchanged.

Equation for calculating the molar mass of an ideal gas:

PV = nRT --> PV = (m/M)RT --> M = mRT/PV n= mass/molar mass (m/M) M = molar mass (g/mol) m= mass (g) R = constant T = temperature (K) P = presssure (atm OR mmHg) V= volume (L)

Which assumptions of the kinetic molecular theory explain the behavior of gases described by Charles's Law (direct relationship between temperature and volume when pressure is constant)?

The average kinetic energy of gas particles is proportional to the absolute temperature in Kelvins. As the temperature increases, the collision frequency increases, causing the pressure inside a container to increase. Since the outside pressure stays the same, the volume of the gas will increase until the inside pressure matches the outside pressure. The average kinetic energy of the particles in a gas is proportional to the temperature of the gas. Because the mass of these particles is constant, the particles must move faster as the gas becomes warmer. If they move faster, the particles will exert a greater force on the container each time they hit the walls, which leads to an increase in the pressure of the gas. If the walls of the container are flexible, it will expand until the pressure of the gas once more balances the pressure of the atmosphere. The volume of the gas therefore becomes larger as the temperature of the gas increases.

How does Kinetic Molecular Theory explain why 1 mole of He and 1 mole of Kr have the same volume at STP?

Kinetic Molecular Theory states that gas particles have negligible size and do not interact. Thus the only property that distinguishes one type of gas from another is mass. Although the particles have different masses, they will all have the same average kinetic energy at a given temperature, with heavier gas particles moving slower than lighter gas particles. Therefore, they will exert the same force upon collision with a surface. Since (pressure X volume) will be the same, at a fixed pressure of 1atm, 1 mole of any gas that obeys the kinetic molecular theory will have the same volume, 22.4 L.

Boyle's Law (at constant temperature and moles, PRESSURE and VOLUME are inversely related)

Kinetic Molecular Theory states that gases can be compressed because most of the volume of a gas is empty space. As long as temperature is constant, kinetic energy of the particles remains the same. Although there is no change in the speed of the particles, they will hit the walls of the container more often. This leads to an increase in pressure. Thus, pressure increases as volume of the container decreases.

Dalton's Law of Partial pressures (total pressure of a gas mixture = sum of the partial pressures of its components)

Kinetic Molecular Theory states: particles have negligible size and do not interact. Thus, the only property that distinguishes one type of particle from another is its mass. However, even particles of different masses have the same average kinetic energy at a given temperature, so they exert the same force upon collision with a surface. Consequently, mixing different gases has the same effect as simply adding more particles. The sum the partial pressures of all components sum to the overall pressure.

Guy-Lussac's Law (at constant volume and moles, PRESSURE and TEMPERATURE are directly related)

Kinetic molecular theory states that increase in temperature leads to increase in the average kinetic energy of a gas particle. As kinetic energy increases, gas particles move faster and harder when they hit the walls of the container. If volume is constant, then pressure must increase as temperature increases.

Relationship between Pressure and Number of Moles of a gas

Kinetic molecular theory states that the pressure of a gas results from collisions between the particles and the walls of their container. An increase in the number of particles leads to an increase in the frequency of collisions. Thus, pressure increases as the amount of gas increases.

Basic Assumptions of the Kinetic Molecular Theory of ideal gases

1) Gases consist of particles so small that their size is insignificant compared to the space between them. Most of the volume of a gas is therefore empty space. 2) the particles are in constant random motion, colliding with the walls of their container and exerting pressure 3) collisions between gas particles are elastic and total kinetic energy is conserved 4) gas particles exert no attractive forces on each other 5) kinetic energy of gas molecules is directly proportional to absolute temperature

Does N2 behave more ideally at 1 atm or at 500 atm? Why?

N2 behaves more ideally at 1 atm because at that pressure, the molecules are sufficiently separated. This means that 1) there will be no intermolecular forces, and 2) they will not occupy a significant percentage of the space in which they are contained.