GASES Chapters 13 and 14

Nature of Gases  Kinetic Molecular Theory (KMT)  Kinetic energy- the energy an object has because of its motion  According to kinetic theory, all matter consists of tiny particles in constant motion

A Model For Gases  The particles in a gas are considered small, hard spheres with insignificant volume  No attractive or repulsive forces  The motion of one particle is independent of that of all other particles

The motion of the particles in a gas is rapid, constant and random  Gases uniformly fill containers regardless of shape and volume  Travel in straight-line paths until a collision occurs  O 2 molecules at 20º C km/hr

All collisions between particles in a gas are perfectly elastic  KE transferred without loss from one particle to another  total KE remains constant (conservation of energy)

Gas Pressure  Results from the force exerted by a gas per unit surface area of an object  Gas pressure is the result of simultaneous collisions of billions of rapidly moving particles in a gas with sides of container

 Any empty space with no particles and no pressure is called a vacuum

Atmospheric Pressure  Results from the collisions of atoms and molecules in air with objects  Air exerts pressure on Earth due to gravity  A. P. decreases as you climb a mountain

 A barometer is a device used to measure atmospheric pressure  Mercury  In fair weather at sea level- 760 mm Hg

Pascal = S.I. Unit of Pressure  Atmospheric pressure ~ 100 kPa  1 atm = 760 mm Hg = kPa  STP= 0º C(273 K) and kPa (1 atm)

Vapor Pressure  Evaporation of a liquid- some particles at surface will evaporate and produce VAPOR PRESSURE

Vapor Pressure  A measure of the force exerted by gas particles above a liquid  Over time (in a closed container), the number of particles entering the vapor increases and some of the particles condense  Reaches an equilibrium

Liquid vapor (gas) evaporation condensation

 In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor and the liquid.  The rate of evaporation of liquid = The rate of condensation of vapor.

Different substances have different vapor pressures  As temperature , vapor pressure   REFERENCE TABLE H!!!!!!!

Boiling Point  When a liquid is heated to a temperature at which the particles throughout the liquid have enough KE to vaporize, the liquid begins to boil

What actually happens?  Bubbles of vapor form throughout the liquid, rise to the surface and escape to the air  The temperature at which the v.p. is just equal to the external pressure of the liquid is called the Boiling Point

B.P. and Pressure Changes  Since a liquid boils when v.p.=external pressure, boiling points vary  Because atmospheric pressure is lower at higher altitudes, boiling points decrease at higher altitudes

Case in Point  DENVER  Atmospheric pressure = 85.3 kPa  Water 95° C

Normal Boiling Point  Boiling Point of a liquid at a pressure of kPa

Properties of Gases  A gas can expand to to fill it’s container  The reverse is also true

Compressibility  A measure of how much the volume of matter decreases under pressure

Air Bags  The compression of the gas absorbs the energy of the impact

Kinetic Theory can explain compressibility  Gases are easily compressed because of the space between the particles in a gas  The volume of the particles in a gas is small compared to the overall volume of the gas

Factors Affecting Gas Pressure  Four variables are generally used to describe a gas:  (P) Pressure- kPa  (V) Volume- L  (T) Temperature- K  (n) number of moles- mol

Amount of Gas  Increasing the number of particles increases the number of collisions, which explains why the gas pressure increases

Material of the Container Matters  Forgiving vs. rigid  Once the pressure exceeds the strength of the container, the container bursts

Pressure Differences  If the pressure of the gas in a sealed container is lower than the outside air pressure, air will rush into the container when it is opened  The reverse is also true  High  Low

Aerosol Cans  Gas stored at high pressure  Pushing the spray button creates an opening  Gas flows through the opening to lower pressure outside

V, T, and P relationships  You can increase the volume and pressure by increasing its temperature  Heat-  temp and KE, impact on container is greater…..  T =  P =  V

Gas Laws  Boyle’s  Charles’  Gay-Lussac’s  Combined  Dalton’s  Graham’s

Boyle’s Law  PRESSURE AND VOLUME  As the pressure of the gas , the volume  (if T constant)  INVERSE Relationship  P 1 V 1 = P 2 V 2

Boyle’s Example  A balloon contains 30.0 L of helium gas at 103 kPa. What is the volume of the helium when the balloon rises to an altitude where the pressure is only 25.0 kPa?

Charles’s Law  TEMPERATURE AND VOLUME  As the temp. of an enclosed gas , the volume  (if P constant)  Direct relationship  Temp must be expressed in kelvins!!!  V 1 /T 1 = V 2 /T 2

Charles

Charles Example  A balloon inflated in a room at 24°C has a volume of 4.00 L. The balloon is then heated to a temperature of 58°C. What is the new volume if the pressure remains constant?

Gay-Lussac’s Law  PRESSURE AND TEMPERATURE  As the temp of an enclosed gas , the pressure  (if V is constant)  P 1 /T 1 = P 2 /T 2

Gay-Lussac Example  The gas in a used aerosol can is at 103 kPa and 25°C. What will pressure be at 928°C?

The Combined Gas Law  Describes the relationship among the pressure, temperature, and volume of an enclosed gas  Allows you to do calculations when only the amount of gas is constant

Combined Law Example  The volume of a gas-filled balloon is 30.0 L at 313 K and 153 kPa. What would the volume be at STP?

Ideal and Real Gases  Ideal Gas- one that follows the gas laws at all conditions of temperature and pressure  Would have to conform to kinetic theory  Its particles could have no volume and there could be no attraction between particles

Real Gases  No gas exists for which those assumptions are true  BUT- at many conditions, real gases behave very much like an ideal gas

Real Gases  Gases can condense and solidify because of the attractions between particles (intermolecular forces)  Real gases differ most from an ideal gas at low temperatures and high pressures

In order to be MOST like an ideal gas:  High Temperature  Low Pressure  Small, nonpolar gases are most ideal under ANY conditions  H 2 and He

Dalton’s Law  The contribution each gas in a mixture makes to the total pressure is called the partial pressure

Law of Partial Pressures  In a mixture of gases, the total pressure is the sum of all the partial pressures of the gases  P total = P 1 + P 2 + P 3 + ……

Dalton’s Example  What is the partial pressure of O 2 at kPa if:  P N = 79.1 kPa  P CO2 =.040 kPa  P other gases =.94 kPa

Graham’s Law  Diffusion: the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout  Effusion: a gas escapes through a tiny hole in its container  GASES OF LOWER MOLAR MASS DIFFUSE AND EFFUSE FASTER THAN GASES OF HIGHER MOLAR MASS

Graham’s Law  If two objects with different masses have the same KE, the lighter object will move faster