Unit 9, page 1
Unit 9 Gases Notes
Matter can exist as a solid, liquid or gas.
We can not observe individual particles of a solid, liquid or gas.
Scientist study groups of particles as they occur as a solid, liquid or gas.
Scientists developed the kinetic-molecular theory of matter to describe the behavior of the atoms and molecules that make up matter.
This theory is based on the idea that particle of matter are always in motion.
The theory can be used to explain the properties of solids, liquids and gases with regards to energy of the particles and
the forces between the particles.
Kinetic Molecular Theory of Gases
This theory provides a model of an ideal gas
Ideal gas: An imaginary gas that fits all 5 assumptions of the kinetic molecular theory
There are no ideal gases in real life, but some gases can come pretty close under certain circumstances.
Monatomic and diatomic gases under low pressure and high temperature are the closest to acting ideal.
Real gas: does NOT meet all 5 assumptions
The theory is based on 5 Assumptions
1) Gases consist of many tiny particles that are far apart relative to their size.
2) Collisions between gas particles and container walls are elastic collisions. No kinetic energy is lost in the collision. (Example: bouncy ball vs clay ball)
This does not ever happen
3) Gas particles are always moving rapidly, randomly, and continuously. 4) Gas particles are not attracted to each other. No intermolecular forces.
This does not ever happen
5) The kinetic energy of gas particles depends on the temperature high temperature = high energy
low temperature = low energy
Properties of Gases
ALL gases have these properties
Expansion: gases take the shape and volume of the container. Example: filling a balloon
Fluidity: gas particles slide past each other
Low Density: Gases have low density. Gases have densities 1/1000 of the same substance as a liquid or solid
Gases are easily compressed. Gases can be compressed so that the particles are forced close together and the volume is significantly decreased.
Gases spread out and mix with one another without being stirred.
Diffusion: The random mixing of different gas particles.
Example: smell of sprayed perfume moving across the room
Unit 9, page 2
Gases and Pressure
Pressure: The pushing force of an object against another object.
Example: The force acting on a person is gravity.
A flat shoe exerts less pressure on the floor than a stiletto high heel shoe. The force of gravity is constant.
The flat shoe has a larger area (denominator) than the high heel. A large denominator will make the pressure value smaller.
Gas molecules exert pressure on the surface they collide with. This is an essential skill! The atmosphere has pressure!
The atmospheric pressure at the beach is greater than the atmospheric pressure in the mountains. At the beach there is more air on top of you than when you are in the mountains.
Atmospheric Pressure can be measured using a barometer. Units of Pressure
1) mm Hg (millimeters of mercury) 2) torr (named after Torricelli) 3) pascal (named after Pascal) 4) atm (atmosphere)
Memorize these conversion factors!
1 atm = 760 torr = 760 mm Hg = 101 kPa
Examples:
Convert 2.5 atm to torr
Convert 789 mm Hg to atm
Convert .347 kPa to atm
Gases and Temperature
Temperature can be measured in Celsius (°C) or Kelvin (K) Know how to convert between these two units of temperature
Degrees Kelvin = °C + 273 Example:
Convert 25 °C (room temperature) to K 25C + 273 = 298K
Unit 9, page 3 Conditions where
temperature = 0 °C and pressure = 1 atm.
Gas Laws
The gas laws are mathematical relationships between volume, temperature, pressure, and the amount of gas.
Boyle’s Law: Pressure and Volume As volume decreases, pressure increases As volume increases, pressure decreases Think about why this is true:
-Pressure is caused by gas molecules hitting the container.
-If the volume of the container is decreased, the gas molecules will hit the container more. -Hitting the container more = increase in pressure.
Charles’ Law: Volume and Temperature When temperature increases, volume increases. When temperature decreases, volume decreases. Think about why this is true:
-Increase in temperature causes gas molecules to move faster. -Faster molecules = more collisions.
-More collisions = more pressure on the inside of the container. -More pressure from the inside of the container = bigger volume.
Gay-Lussac’s Law: Pressure and Temperature When temperature increases, pressure increases. When temperature decreases, pressure decreases.
Combined Gas Law
This gas law combines Boyle’s, Charles’, and Gay-Lussac’s Laws into one mathematical relationship:
OR T1P2V2 = T2P1V1
Example: A sample of air has a volume of 140.0 mL at 67°C. At what temperature will its volume be 50.0 mL?
Use combined gas law and ignore the pressure. V1= 140.0mL T1= 67°C V2=50.0mL T2 =?
T1 = 340K
T1P2V2 = T2P1V1 T1V2 = T2V1 T2 = 121K
T2 = -151°C
Partial Pressure
Unit 9, page 4
Dalton’s Law of Partial Pressure:
Total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases.
Total pressure = the sum of individual partial pressures Ptotal = P1 + P2 + P3 etc…
Gases Collected by Water Displacement
Liquid water will always have a small amount of water vapor molecules existing at the surface of the water.
Gases are usually produced and collected over water.
Because there is a small amount of water vapor molecules near the surface of the water, the gas collected by water displacement will be a mixture of water vapor and the collected gas. Those water vapor molecules exert pressure because they are gases. The total pressure inside the collection device is due to the gas, and the water vapor molecules.
Ptotal = Pgas + Pwater vapor
If a gas is collected by bubbling through water, then need to know the partial pressure of the water vapor at the temperature of the water.
Look it up on a chart: Book Page 899. Table A-8 Or it will give it to you in the problem.
Example of Partial Pressure Problems
Unit 9, page 5 Ptotal = PO2 + PH2O
PO2 = Ptotal – PH2O
PO2 = 731 torr – 17.5 torr
PO2 = 713.5 torr
Standard Molar Volume of a Gas
1 mole of gas (any gas) at STP occupies a volume of 22.4 L Equality: At STP 1 mol = 22.4 L
If have an equality can use it as a conversion factor. Use this to convert between moles and liters! Same Amount of Gas = Same Volume
1 mole of O2 has the same volume as 1 mole of H2
0.5 mole of O2 has the same volume of 0.5 H2
And so on…
However, 1 mol of O2 weighs more than 1 mol of H2, even though both occupy a volume of 22.4 L
Example: What volume will 0.068 mol of O2(g) occupy at STP?
Example: How many moles of H2(g) are in 40.6 L, at STP?
Example: How much does 0.098 L of SO2(g) weigh at STP?
CAN ALSO USE THE IDEAL GAS LAW
Look what you can convert using mole!
Mole to molecules/atoms using Avagadro’s number = 6.02 x 10
23Mole to grams using molar mass
Mole A to Mole B using mole ratio
Moles to Liters using standard molar volume = 22.4 L
Ideal Gas Law
Unit 9, page 6 -temperature (T) MUST BE IN K
-number of moles of a gas (n)
PV = nRT R is the ideal gas constant
The value of R is calculated by rearranging the equation to solve for R and plugging in the values for 1 mole of a gas at standard conditions (1 atm of pressure and 0°C, 273 K)
Sample Ideal Gas Law Problems:
Example: What is the pressure in atm exerted by a 0.5 mol sample of nitrogen gas in a 10. L container at 298 K?
P = ? mol = .5mol N2 Vol = 10.L T= 298K
PV = nRT P = 1.2 atm
Example: What is the volume in liters, of 0.250 mol of oxygen gas at 20°C and 0.974 atm?
V = ? mol = .250 mol T= 20°C + 273 = 293K P = .947 atm
PV = nRT V = 6.17 L
STEPS to Solving Gas Problems 1. list variables with units 2. Write equation
3. Manipulate equation for variable that you are trying to solve for 4. Plug in values for variable including units
