Embed Size (px)
Citation preview
UNIT 6 GASES
OUTLINE
• States of Matter, Forces of Attraction
•Phase Changes
•Gases
•The Ideal Gas Law
• Gas Stoichiometry
STATES OF MATTERRemember that all matter exists in three physical states: Solid Liquid Gas
INTERMOLECULAR FORCESWhy are some substances gases at room temperature (eg.
CO2) while some substances are solid at room temperature
(eg. Iron)?
Intermolecular forces of attraction is the general term
to describe how particles are held together in a substance.
1. Dispersion
2. Dipole-Dipole
3. Hydrogen bonding
INTERMOLECULAR FORCES•Strong intermolecular forces = higher
melting point, boiling point.
Intermolecular means between molecules
Intramolecular means inside molecules (ie.
chemical bonding)!
Do not confuse these two things
1. DISPERSION FORCEDispersion forces are also called London dispersion
forces.
Dispersion forces are weak forces that result from
temporary changes in the density of electron
clouds.
•Dispersion forces exist between ALL particles in
ALL substances.
•Bigger electron cloud (ie. substance with more
electrons) = stronger dispersion force. Eg. Gr.17
1. DISPERSION FORCEDispersion forces are also called London dispersion forces.
Dispersion forces are weak forces that result from
temporary changes in the density of electron clouds.
2. DIPOLE-DIPOLE FORCESDipole-dipole forces are the attractions between
oppositely charged parts of polar molecules’ dipoles.
•Dipole-dipole forces ONLY exist between
particles with a permanent dipole. Polar
molecules only. Eg. HCl
•Dipole-dipole forces are slightly stronger than
dispersion forces, but if that molecule is very large,
dispersion forces are more significant.
2. DIPOLE-DIPOLE FORCES
3. HYDROGEN BONDNOT a chemical bond!!
A hydrogen bond is a type of dipole-dipole attraction.
A H that is chemically bonded to an O, N or F inside a molecule: positive.
An O, N or F that is chemically bonded to a H inside a molecule: negative.
•On two different molecules (which contain H-O/N/F bond) one H and another O/N/F are electrostatically attracted.
• This is the strongest type of intermolecular force.
INTERMOLECULAR FORCESFor each of the following compounds, determine the main
intermolecular force.
1. Nitrogen 8. SiH2O
2. Carbon tetrachloride
3. H2S
4. Sulfur monoxide
5. N2H2
6. BH3
7. CH4O
INTERMOLECULAR FORCES: NOTES!•Strong intermolecular forces = Stronger forces of attraction between particles = harder to separate = higher melting point, higher boiling points!
•Dispersion forces are weakest when comparing substances of similar size.
•Dispersion forces are stronger than other IMF is the substance is much larger than the other substances!
INTERMOLECULAR FORCES: NOTES!•Dispersion forces are weakest when
comparing substances of similar size.
•Dispersion forces are stronger than other
IMF is the substance is much larger than
the other substances!
* See Question 5 on IMF Worksheet
PHASE CHANGES
•Phase Diagrams (12.4)
Most substances exist in three states, depending on the temperature and pressure.
Phase diagram is a graph of pressure versus temperature that shows the phases a substance exists for different T and P.
PHASE CHANGESPhase changes require energy:
Melting:
Ice has water molecules that are close together, held by H-bonding. Heat is transferred to the water molecules, and the molecules absorb enough energy to break these IMFs so that the molecules move further apart, into the liquid phase.
PHASE CHANGESVaporization: liquid changing to gas
Vapor Pressure: Pressure exerted by a
vapor over the surface of a liquid.
Boiling point: The temperature where
the vapor pressure of a liquid equal the
external atmospheric pressure.
PHASE CHANGES
Phase changes that require energy:
Melting
Vaporization
Sublimation: Solid directly to gas.
PHASE CHANGESPhase changes that release energy:
Freezing
Condensation
Deposition
STOICHIOMETRY: SOLUTIONS CH.14
• What is the critical temperature of
compound X?
• If you were to have a bottle containing
compound X in your closet, what
phase would it most likely be in?
• At what temperature and pressure will
all three phases coexist?
• If I have a bottle of compound X at a
pressure of 45 atm and temperature of
100°C, what will happen if I raise the
temperature to 400°C?
• Why can’t compound X be boiled at a
temperature of 200°C?
• If I wanted to, could I drink compound
X?
GASES•Gases look and behave very differently from
liquids and solids.
•To try and explain the properties of gases,
scientists came up with the kinetic-
molecular theory of gases.
•This theory describes the behaviour of matter
in terms of particles in motion.
KINETIC-MOLECULAR THEORY *IMPORTANT*
1. Gases are made up of small particles that
are separated from one another by empty
space.
2. Volume of the particles is small compared
with the volume of the empty space
between particles.
3. There are no forces of attraction
between particles.
KINETIC-MOLECULAR THEORY
4. Gas particles are in constant,
random motion.
5. Collisions between gas particles are
elastic (do not lose energy)
KINETIC-MOLECULAR THEORY
The theory can explain properties of
gases.
1. Gases expand because:
-constant, random motion
-no attractive forces between particles
KINETIC-MOLECULAR THEORY
2. Gases can contract because:
-gas particles are tiny and far apart
3. Gases have low density because:
-gas particles are tiny and far apart
-no attractive forces between particles
KINETIC-MOLECULAR THEORY
The theory can explain why gases have
low density (mass per volume) and can
be compressed or expanded (random
motion of particles fills the available
space)
KINETIC-MOLECULAR THEORY
Diffusion: movement of one
substance through another
Effusion: gas escaping through
an opening.
Gases can diffuse and effuse because:
constant, random motion
KINETIC-MOLECULAR THEORY
Graham’s law of
effusion: Heavier gases
effuse more slowly than
lighter gases.
KINETIC-MOLECULAR THEORY
The temperature is a measure of the
average kinetic energy of the gas
particles in a sample.
KE = ½ m v2
KINETIC-MOLECULAR THEORY
The temperature is a measure of the
average kinetic energy of the gas
particles in a sample.
KE = ½ m v2
KINETIC-MOLECULAR THEORY
Gas Pressure: Gas particles
exert pressure when they collide
with the walls of their container.Units: atmosphere (atm), pascal (Pa)
KINETIC-MOLECULAR THEORY
Dalton’s Law of Partial
Pressures: The total pressure of
a mixture of gases is equal to the
sum of the pressures of all the
gases in the mixture.
KINETIC-MOLECULAR THEORY
KINETIC-MOLECULAR THEORY
Partial pressures can be used to find the
amount of gas produced by a reaction.
Eg. Gas collected becomes a mixture. Total
pressure inside container is the sum of the
partial pressures of water vapor and new gas.
KINETIC-MOLECULAR THEORY
KINETIC-MOLECULAR THEORY
Knowing the pressure, volume of
the container, and temperature of
a gas allows you to calculate the
number of moles of gas!
KINETIC-MOLECULAR THEORY
Real gases:
1. Particles of real gases do have a
physical volume
2. Particles of real gases can exert
attractive forces on each other
KINETIC-MOLECULAR THEORY
Real gases behave like an “ideal gas”:
1. Low pressure (particles far apart)
2. High temperature (lots of
kinetic energy)
3. Weak attraction to each other
THE GAS LAWS•Boyle’s Law: At constant temperature,
increasing pressure decreases volume.
•Charles’s Law: At constant pressure,
increasing temperature increases volume.
•Guy-Lussac’s Law: At constant volume,
increasing temperature increases pressure.
THE IDEAL GAS LAW•Avogadro thought that because particles are
so small, 1000 large krypton gas particles would
occupy the same volume as 1000 small helium
gas particles.
•Equal volumes of gases at the same
temperature and pressure contain equal
numbers of particles.
THE IDEAL GAS LAW•So at 0oC and 1.00 atm of pressure, 1 mole of
gases (1 mol = 6.02 x 1023 particles) occupy
22.4 L of volume.
•This is true for ANY gas (no matter what size
the gas particle actually is)
•Eg. You have 3.50L of a gas. How many moles of
this gas do you have?
THE IDEAL GAS LAW•Combining all the laws!
•The ideal gas constant,
R = 0.0821 L atm/mol K
•PV = nRT
