UNIT 6 GASES...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

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)


The theory can explain properties of

gases.

1. Gases expand because:

-constant, random motion

-no attractive forces between particles


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


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)


Diffusion: movement of one

substance through another

Effusion: gas escaping through

an opening.

Gases can diffuse and effuse because:

constant, random motion


Graham’s law of

effusion: Heavier gases

effuse more slowly than

lighter gases.


The temperature is a measure of the

average kinetic energy of the gas

particles in a sample.

KE = ½ m v2


The temperature is a measure of the

average kinetic energy of the gas

particles in a sample.

KE = ½ m v2


Gas Pressure: Gas particles

exert pressure when they collide

with the walls of their container.Units: atmosphere (atm), pascal (Pa)


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.



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.



Knowing the pressure, volume of

the container, and temperature of

a gas allows you to calculate the

number of moles of gas!


Real gases:

1. Particles of real gases do have a

physical volume

2. Particles of real gases can exert

attractive forces on each other


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

UNIT 6 GASES...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

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