Chemistry Chapter 10 notes

Physical Characteristics of Gases

Kinetic molecular theory of matter

• All matter is composed of tiny particles which are in constant motion

• This explains observed properties of matter

Kinetic molecular theory of matter (KM)

• Ideal gas- an imaginary gas which perfectly fits all assumptions of the kinetic molecular theory of matter

• IE: Ideal gas behaves exactly as a gas should, no deviations

• Kinetic molecular theory of gases based on 5 assumptions

5 Assumptions

1. Gases consist of large numbers of particles very far apart from one another relative to their size

– Most of space occupied by gases is empty space

– This explains compressibility of gases and their low density (compared to liquids and solids)

2. Collision between gas particles/particles and gas particles/container are elastic

– Elastic collision- no net loss of kinetic energy.

– KE is transferred, but total KE of 2 particles remains the same as long as the temperature is constant

3. Gas particles are in continuous, rapid, random motion and therefore have kinetic energy

• Their KE is high enough to overcome any attractive forces between particles (except near the temperature of condensation)

4. There are no forces of attraction or repulsion between gas particles

• When gas particles collide, they immediately bounce apart

5. Average kinetic energy of gas particles depends on the temperature of the gas – For any object KE= ½ m v2

– Where m= mass and v = velocity– All gases at the same temperature have same

KE, so lighter particles (H) have higher average speed than heavier particles (O)

KM theory and the nature of gases

• Expansion– Gases fill any container and take it’s shape– Gases have no definite shape or volume

• Fluidity– Gas particles glide past one another– Behave much like liquids– Gases and liquids are both considered fluids

KM theory and the nature of gases

• Low density– Gases typically have about 1/1000 the density

of the same substance in a liquid or solid state

• Compressibility– Due to their low density gases can be

compressed dramatically

KM theory and the nature of gases

• Diffusion– Gases randomly mix with other particles to

even distribution– Rates of diffusion depend on the speed of

particles, diameter of particles and attractive forces between particles

– Lighter gases diffuse more rapidly than heavier ones

KM theory and the nature of gases

• Effusion– Movement of gas particles through a tiny

opening– Rates of effusion are directly proportional to

the velocity of the particles

Real gases

• Do not behave completely according to kinetic molecular theory

• 1873 Van der Waals noted that forces between particles of gases caused deviation from ideal gas behavior

• Deviation is most significant at high pressure and low temperature

• KM theory holds truest in gases with little attraction between particles (ex. Noble gases)

Pressure

• When describing a gas you must specify characteristics: Volume, temperature, number of molecules and pressure

• You’ve got the first 3!

• Pressure is force per unit area on a surface or Pressure = force/area

Atmospheric pressure

• Pressure exerted by gases of the atmosphere

• At sea level approximately 10.1 N/cm2

• Barometers are used to measure atmospheric pressure

• Oldest barometer- mercury column measurement expressed in mm of Hg– Normal atmospheric pressure at sea level and

0°C = 760 mm Hg = 1 atmosphere

Pressure!

• SI units for pressure are derived

• 1 Pascale (Pa) = 1 Newton / meter2

• Pressure often expressed in kilopascals (kPa)

• 1 atmosphere = 1.01325 x 105 Pa

(or 101.325 kPa)

• See table 10-1 on p. 311

STP

• Standard Temperature and Pressure are needed to compare gas volumes

• STP = 1 atmosphere and 0°C

Gas Laws

Boyles law

• Relates pressure and volume of a gas at constant temperature

• Pressure and volume are inversely proportional

• PV = k or V=k1/P

• K is a constant for a given sample of gas

Boyles and changing pressure

• Because k is a constant for a given sample of gas and we know that the product of pressure and volume will always equal k

• P1V1 = k and P2V2 = k we can set P1V1

equal to P2V2

• P1V1= P2V2 and solve for any one of the 4 values

Charles Law

• Relates temperature and volume of gases at constant pressure

• 1787 Charles found that volume of a gas changes 1/273 of original volume for each 1°C change in temperature (with a starting point of 0°C and at constant pressure)

Charles and absolute zero

• Kelvin 0 = -273.15°C

• K= °C + 273.15

• This is useful because it is directly proportional to gas volume

• Charles law: Volume of a fixed sample of gas at constant pressure varies directly with Kelvin temperature

Charles…

• V/T= k or V = kT

• K is a constant based on quantity of gas and pressure

• Same thing can be done with Charles for changing volume or temperature as was done with boyles for changing pressures

• V1/T1 = V2/T2

Gay-Lussacs Law

• Relates pressure and temperature of a gas at constant volume

• P/T = k or P= kT

• K is a constant depending on quantity and volume of gas

• P1/T1 = P2/T2 useful when faced with changing pressures and temperatures

Combined gas law

• Merges three laws just mentioned

• PV/T = k

• k is a constant related to the amount of gas

• P1V1/T1 = P2V2/T2

• if any one quantity is unchanging one of the other gas laws can be derived

Daltons combined pressures

• The total pressure of a mixture of gases is the sum of the individual pressure of each gas alone

• PT= P1 + P2 + P3…

• This can be used no matter how many gases are in combination

Law of Combined Pressures

• Is useful when dealing with gases collected over water

• Gases collected this way are mixed with water vapor, this exerts water vapor pressure

• To measure pressure of gas and water vapor in collection bottle, raise bottle until water level in and out are same.

• At that point pressure inside bottle = atmospheric pressure

• Patm = P gas + P H2O

• Obtain atmospheric pressure from barometer in lab and subtract water vapor pressure at given temp (from table A8 in book)