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Unit 12

States of Matter

2

Gases

• Kinetic – movement

• Kinetic theory – explains the

behavior of particles in terms of

their motions

3

Kinetic Theory of Gases

Kinetic Theory of Gases

1. Negligible volume

2. Move in rapid, constant straight-line motion

3. Collide elastically

• No net loss in kinetic energy

4. Far apart with no attractive or repulsive force

4

• Kinetic Energy – energy of motion

– KE = ½ (m)(v)2

• m = mass

• v = velocity

• Temperature – measure of average kinetic energy.

– Molecules within a substance don’t all move at the

same speed

• They have a constant average speed!

Kinetic Energy and Temperature

5

Which has the higher average kinetic energy?

Red

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Kinetic Energy and

Temperature

• The higher the temperature, the

faster moving the particles

• At 0 K all molecular motion

would stop

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Explaining the Behaviors of Gases

• Low Density

– – Gases are far apart

• Low mass

• High volume

volume

massDensity

• Compression

– If gases are far apart,

they can be pushed

together

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• Expansion

– If gases are moving in constant straight-line

motion, they will expand to fill any container.

Explaining the Behaviors of Gases

9

• Diffusion – tendency of molecules to move from

an area of high concentration to low concentration

until uniformly mixed.

– If particles are in rapid straight-line motion, the will

spread out.

– Perfume sprayed in a room

Explaining the Behaviors of Gases

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• Effusion – rate at which gases escape from

a small hole in a container.

Explaining the Behaviors of Gases

Both diffusion and effusion depend upon the

velocity of the particles!

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Graham’s Law

Graham’s Law

The rate of effusion of a gas is inversely

proportional to the square root of the gases

molar mass.

RateA Molar MassB

RateB Molar MassA

=

12

Graham’s Law Derived

If both gases are in the same

room, then they have the same

temperature.

Would they be moving at the same velocity?

KEKE BA

2

BB

2

AA Vm2

1 Vm

2

1

No, the lighter gas is faster!

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What happens if we play with this

equation?

A

B

2

B

2

A

m

m

V

V

2

BB

2

AA Vm2

1 Vm

2

1

A

B

B

A

m

m

V

V You get

Graham’s law!

Multiply both sides by 2 and get the velocities and masses on the same side.

Get rid of the squared values by taking the square root of both sides.

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Graham’s Law Problems

1. Finding Rate Ratios (How much faster…)

1. Decide what is “A” (listed first) and “B”

2. Setup equation appropriately and solve.

2. Finding Molar Mass

1. Establish rate ratio

2. Which gas is faster?

• This gas becomes the bigger number in the rate ratio and establishes what is “A” and what is “B”.

3. Setup equation appropriately and solve.

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Graham’s Law

How much faster does oxygen diffuse (or

effuse) than nitrogen.

RateA Molar MassB

RateB Molar MassA

=

Which is A? O2

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Graham’s Law

If O2 is A, then N2 would be B. The rate of

oxygen to nitrogen would be the x.

oxygen

nitrogen

MM

MM x

32.0

28.0 x

X = .935 Oxygen is .935 times faster

than nitrogen (less than 1

means it’s slower)!

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What is Wind?

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• Pressure – simultaneous collisions of billions of

molecules with an object

• Pressure =

– The force is created by the collisions

• More collisions higher pressure

• Harder collisions higher pressure

– Thumbtacks and snowshoes take advantage of

this to work.

Gas Pressure

force

area

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Gas Pressure

• Vacuum- empty space means no collisions

• Atmospheric pressure - air molecule collisions – We live at the bottom of an ocean of air

– If you go up a mountain, atmospheric pressure decreases because the depth of air above you is less.

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• Units of pressure

–Pascal (Pa)-SI unit

–Atmosphere (atm)-Typical sea level

pressure

–Millimeters mercury (mm Hg)-

barometric reading

Gas Pressure

760 mm Hg = 1 atm = 101.3 kPa = 14.7 psi

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Measuring Pressure – Barometer

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Measuring Pressure – Manometer

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Dalton’s Law of Partial

Pressures

Dalton’s Law

At constant volume and temperature, the total

pressure exerted by a mixture of gases is

equal to the sum of the individual (partial)

pressures.

Ptotal = P1 + P2 + P3 • • • •

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Dalton’s Law

?

200 kPa + 500 kPa + x = 1100 kPa

x = 1100 kPa – (500 kPa + 200 kPa)

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Dalton’s Law

• Dalton’s Law Correction

– Because most gases are COLLECTED OVER

WATER, you will often use this law to correct

the pressure.

Ptotal = Pgas + Pwater Solving for the gas yields:

Pgas = Ptotal – Pwater This is usually

the pressure used

in a gas law

problem (next

chapter)!

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13.2 Forces of Attraction

What separates gases from solids

and liquids?

(What do solids and liquids have that

gases don’t)

Attractive Forces

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Forces of Attraction

There are two types of attractive forces

1. Intramolecular forces – forces within a

compound that hold it together.

2. Intermolecular forces – forces between

compounds that hold these compounds

together.

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Intramolecular Forces

Force Model Basis of

Attraction

Energy

(kJ/mol) Example

Ionic Opposite

charges 4000-400 NaCl

Covalent

Shared

electron

pair

1100-150 H2

Metallic

Metal

cations and

delocalized

electrons

1000-75 Fe

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Intermolecular Forces

Forces Basis of

Attraction

Energy

(kJ/mol) Example

London

dispersion

forces

Momentary

dipole – shifting

of electron

cloud

40-.05 H2 to H2

Dipole-

Dipole

Partial charges

of polar

molecules 25-5 HCl to HCl

Hydrogen

Bonding H bonded to an

N, O, or F 40-10

H2O to

NH3

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• Molecules are close together, but not so

tightly packed they can’t move around.

• Unlike gases, liquids are held together

by attractive forces (intermolecular

forces – forces between molecules)

– This gives liquids a fixed volume

Nature of Liquids

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• Molecules are not moving fast enough to break attractive force and become a gas

• Not very compressible – too close together

Nature of Liquids

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Nature of Liquids

• Density

– Intermolecular force holds liquids together causing a

small volume – this means the density is higher than as

a gas

• Fluidity

– Molecules aren't so close together they can’t move, but

intermolecular forces make them less fluid than gases

– Viscosity – resistance of a liquid to flow

• Decreases with increase in temperature because the

motion allows the molecules to overcome

intermolecular forces

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Nature of Liquids

• Surface Tension

– Energy needed to

increase the surface area

of a liquid

– Caused by an imbalance

of intermolecular force

on the surface of a liquid

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Surface Tension

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Nature of Liquids

• Capillary Action – Cohesion – force of attraction between identical

molecules

– Adhesion – force of attraction between different molecules

• The adhesive force between water and glass is stronger than the cohesive force between water molecules

– Meniscus

– Capillary action

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• Tightly packed - held together by

stronger intermolecular forces than

liquids or gases

• Do not flow - molecules vibrate

around fixed points

• Incompressible

The Nature of Solids

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The Nature of Solids

• Density

– Generally the solid form of a

substance is the most dense

• ~10% more dense than liquids

• Packing tightly together decreases

volume

– H2O is the exception

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• Crystalline

– Most solids are crystalline

– Crystal - the atoms, ions or molecules are

arranged in a very orderly, repeating 3D

pattern known as the crystal lattice

The Nature of Solids –

Types of Solids

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The Nature of Solids –

Types of Solids

• Unit Cell

– Smallest group of particles that retains the

shape of the crystal.

• Allotropes

– 2 or more different molecular forms of the

same element in the same state.

• Diamond and graphite

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Types of Crystalline Solids

Type Unit Particle

Characteristics

of

Solid

Example

Atomic atoms •Soft

•Very low melting point

•Poor conductivity Noble gases

Molecular molecules •Fairly soft

•Low to medium-high mp.

•Poor conductivity I2, H2O, CO2

Covalent

Network

Atoms connected by

covalent bonds

•Very hard

•Very high mp.

•Poor conductivity

Diamond (C)

and Quartz

(SiO2)

Ionic Ions •Hard and brittle

•High mp.

•Poor conductivity NaCl

metallic Metal cations

surrounded by mobile

valence electrons

•Soft to hard and malleable

•Low to high mp.

•Excellent conductivity Any metal

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• Amorphous

– No ordered internal structure

– No definite melting point, melt over a

range or gradually “soften”

– Example Glass, rubber, and many

plastics

The Nature of Solids –

Types of Solids

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1. Melting point - temperature that a solid turns

into a liquid

– As the solid is heated, the molecules vibrate faster

until these vibrations break the intermolecular forces

that hold the solid structure together

– The stronger the intermolecular force the higher the

melting point

– Ionic compounds have stronger intermolecular

forces than molecular compounds

Phase Changes Requiring Energy

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Phase Changes Requiring Energy

2. Vaporization – process of a liquid changing to a gas or vapor

– Evaporation - Liquid to gas below boiling point

• Happens only on surface

• Only particles moving fast enough to overcome intermolecular forces escape to become a gas

• Cooling process - as high speed particles escape this lowers the average kinetic energy and therefore the temperature

• Heating liquid makes more high speed particles and rate of evaporation increases

44 What if this were the energy needed for

vaporization

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– Vapor pressure - vaporized particles will

exert an outward pressure • increasing temperature will increase vapor pressure

Phase Changes Requiring Energy

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Why do different substances have different vapor pressures?

Stronger

Intermolecular

Forces!

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– Boiling Point - when vapor pressure equals

atmospheric pressure

• Unlike evaporation, boiling happens at a set

temperature and is not limited to the surface

• Temperature of liquid never gets above the

boiling point

Phase Changes Requiring Energy

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3. Sublimation – change of a solid

directly to a gas

– Solid CO2 is called “dry ice” because it

sublimates from a solid to a gas without

becoming a liquid

– Ice sublimes and this is how “frost free”

refrigerators work

Phase Changes Requiring Energy

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Boiling Point Curve

0

20

40

60

80

100

120

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17

Tem

pe

ratu

re

(C)

Time (s)

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1. Condensation – gas or vapor becomes a

liquid

– slow moving gas particles that attract each

other and become a liquid

– high speed particles that enter the liquid

become trapped

2. Deposition – gas to solid (snowflakes)

3. Freezing – liquid to a solid

Phase Changes Releasing Energy

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What happens to a liquid in a closed container?

– At first high evaporation and low condensation

Phase Change Equilibrium

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• Closed Container Continued

– Over time evaporation rate will stay constant (stays at

constant temp in room) and condensation rate

increases as the number of gaseous particles increases

– Eventually the rates will become equal and dynamic

equilibrium is reached (particles are still traded but in

equal amounts so no net change)

Phase Change Equilibrium

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Phase Diagram

• Phase Diagram – diagram that shows

relationship between solid, liquid and

gas states for a substance in a closed

container at different temperatures and

pressures

Phase Diagram of Water

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57

• Triple point - temp and pressure where all three states of matter co-exist in equilibrium

• Critical point – above this temperature, the substance can only exist as a gas

Phase Diagram

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Plasma: The Fourth State of Matter

• Gaseous mixture of electrons and

positive ions

• Generally occurs only at high

temperatures where the kinetic

energy is enough to strip electrons

off of gaseous atoms

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• Happens to a small degree (very

unstable and only lasts for a brief

moment in time) at lower

temperatures in fluorescent lights,

neon signs, lightning and the

“northern lights.”

Plasma: The Fourth State of Matter

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Graphite Diamond

Two

allotropes of

carbon

http://www.creative-chemistry.org.uk/molecules/carbon.htm