Chapter 12

States of Matter

12.1 Kinetic Molecular Theory

• Describes the behavior of matter in terms

of particles in motion

– Makes several assumptions about the size,

motion, and energy of gas particles

Assumptions of the Kinetic

Molecular Theory

1. Gases consist of small particles

2. Gases take up little volume relative to the

volume of empty space around them so

size of individual particles is zero

3. Gas particles move in constant, random

straight lines until they collide with other

particles or with the walls of the container

– Collisions are elastic –

– Collisions cause gas pressure -

4. Gas particles do not attract or repel each

other

5. The average kinetic energy of gas

particles is directly proportional to the

KELVIN temperature of the gas

12.2 Gases

• Gases expand, diffuse, exert pressure,

and can be compressed because they are

in a low density state consisting of tiny,

constantly moving particles

Explaining the Behavior of Gases

• Kinetic molecular theory helps explain the

behavior of gases

– Blowing up a balloon

• Gases have low densities

– Why?

• Gases can be compressed and can

expand

– Explain

• Gases diffuse and effuse

– Diffusion = gas particles move from an area of

high concentration to low concentration

• Ex

– Effusion = gas particles escape through tiny

openings

• Ex

12.3 Forces of attraction

• Intermolecular forces (dispersion forces,

dipole-dipole forces, and hydrogen bonds)

determine a substance’s state at a given

temperature

Intermolecular forces

• Inter- means between or among

• Intermolecular forces can hold together

identical particles or two different types of

particles

• Weaker than intramolecular forces (bonds)

Dispersion Forces

• Weak forces that result from temporary

shifts in the density of electrons in electron

clouds

• Exist between all particles

– Weak for small particles

– Get stronger as the number of electrons

involved increases

– F2

– Cl2

– Br2

– I2

Dipole-dipole forces

• Attraction between oppositely charged

regions of polar molecules

– Polar molecule =

• Neighboring polar molecules orient

themselves so that oppositely charged

regions align

Hydrogen Bonds

• Dipole-dipole attraction that occurs

between molecules containing a hydrogen

atom bonded to a flourine, oxygen, or

nitrogen atom

• Explain why water is liquid at room

temperature while compounds of similar

masses are gases

12.4 Liquids and Solids

• The particles in solids and liquids have a

limited range of motion and are not easily

compressed.

Liquids

• Kinetic molecular theory also applies to

liquids and solids

– Must take intermolecular forces into account

to apply it

• Density and compression

– Much denser than gasses

• Due to intermolecular forces holding particles

together

– Incompressible

• Why can you compress a gas but not a liquid?

• Fluidity – both gases and liquids are

classified as fluids because they can flow

and diffuse

– Liquids diffuse more slowly because

intermolecular attractions interfere with the

flow

• Viscosity - measure of the resistance of a

liquid to flow

– Attractive forces – stronger intermolecular

forces = higher viscosity

– Particle size – larger molecules = higher

viscosity

– Temperature – lower temperature = higher

viscosity

• Surface tension – the energy required to

increase the surface area of a liquid by a

given amount

– Caused by intermolecular forces pulling down

on the particles on the surface of a liquid

which stretches it tight like a drum

• Stronger the attraction between particles

in a liquid = greater surface tension

• Surfactant – lowers the surface tension of

water by disrupting hydrogen bonds

between water molecules

• Cohesion – force of attraction between

identical molecules

• Adhesion – force of attraction between

molecules that are different

Solids

• Solid particles have as much kinetic

energy as liquids or gasses but much

stronger attractive forces between

particles

– Limit the motion of particles to vibrations

• Density of solids –

almost always greater

than density of liquids

– Exception = water

• Crystalline solids – solid whose atoms,

ions, or molecules are arranged in an

orderly, geometric structure

– Unit cell = smallest arrangement of atoms in a

crystalline solid that has the same shape as

the whole crystal

• Categories of crystalline solids

– Classified based on the types of particles they

contain and how they are bonded together

• Molecular solids

– Molecules are held together by dispersion

forces, dipole-dipole forces or hydrogen

bonds

– Most are not solid at room temperature

– Poor conductors

• Covalent network solids

– C or Si, can form multiple covalent bonds

which allow it to take many forms

– Allotrope – element that can exist in different

forms at the same state

• Ionic solids

– Made of cation + anion

– Each ion is surrounded by ions of the

opposite charge

– High melting point

– Brittle

• Metallic solids

– Positive metal ions surrounded by a sea of

mobile electrons

Amorphous solids

• Particles are not arranged in a regular,

repeating pattern

• Does not contain crystals

• Forms when molten material cools too

quickly for crystals to form

– Glass

– Rubber

– Some plastics

12.5 Phase changes

• Matter changes phases when energy is

added or removed

Phase changes that require energy

• Melting

– Heat flows from an object at a higher

temperature to an object at a lower

temperature

– Ice absorbs heat which does not raise

temperature but is used to break hydrogen

bonds

– When hydrogen bonds are broken molecules

can move further apart into the liquid phase

• Melting point – temperature in which

forces holding a solid together are broken

and it becomes a liquid

• Vaporization – process by which liquid

changes to vapor

– Vapor – gaseous state of a substance that is

normally liquid at room temperature

– Evaporation – when vaporization occurs only

at the surface of a liquid

– Vapor pressure – the pressure exerted by a

vapor over a liquid

– Boiling – temperature at which the vapor

pressure of a liquid equals the atmospheric

pressure

– Energy being input causes molecules to move

around more and vaporize

• Sublimation – changing from solid to gas

without becoming a liquid

– Dry ice

– Moth balls

– Solid air fresheners

Heating Curve

Phase changes that release energy

• Freezing

– Heat flows out of warmer object into cooler

object

– Molecules slow down & become less likely to

flow past one another

– Intermolecular forces cause the molecules to

become fixed into set positions

– Freezing point – temperature in which a liquid

becomes a solid

• Condensation – process by which a gas or

vapor becomes a liquid

• Deposition – substance changes from gas

or vapor to solid without first becoming a

liquid

– frost

Phase Diagrams

• Temperature and pressure both effect the

phase of a substance

– Have opposite effects

• Phase diagram – graph of pressure vs

temperature that shows which phase a

substance will be in under different

conditions.

• Triple point = point at which all three

phases exist at the same time