Chapter 10

Kinetic Theory (Kinetikos - “Moving”)

Based on the idea that particles of matter are

always in motion

The motion has consequences

Behavior of Gases

Physical Properties of Gases

Ideal Gas – an imaginary gas that conforms

perfectly to all assumptions

Five Assumptions of the KMT

1. Gases consist of large number of tiny particles

2. The Particles are in Constant Motion, moving in

straight lines.

3. The collisions between particles & w/ the container

wall are elastic.

4. There are no forces of attraction or repulsion between

the particles of a gas.

5. The average K.E. of the particles is directly

proportional to the Kelvin Temperature.

KE = ½ mv2

Pressure exerted by the column of air in the atmosphere.

Result of the earth’s gravity attracting the air downward.

Barometer – device used to measure the atmospheric pressure on earth.

Manometer – device used to measure the pressure of a gas in an enclosed container.

Physical Properties of Gases

Defined as having indefinite shape and volume.

Gases have mass

Expands to occupy any space available.

Easily compressed

Different gases move easily through each other.

Diffusion – spontaneous mixing of 2 gases.

Low mass = High rate

Effusion – gas passes through tiny opening.

Gases exert pressure

Fluidity – ability to flow.

Low density

Real Gas

Gas that does not behave completely to the

assumption of the KMT.

Deviation from ideal behavior:

High Pressure

Low Temperature

Polar molecules

10-2 Liquids

Definite volume but no definite shape.

Liquids are fluids.

Liquids attractions are a result of the IMF’s.

Liquid Properties

Surface Tension

Force that tends to pull adjacent parts of a liquid’s

surface together, decreasing the surface area to the

smallest possible size.

Imbalance of forces at the surface of a liquid.

Capillary Action

Attraction of the surface of a liquid to the surface of a

solid.

Viscosity

“Friction” or Resistance to motion, that exist

between molecules in a liquid.

High Viscosity = Low Flow

Stronger IMF = Higher Viscosity

Increase KE = Low Viscosity

Evaporation/Boiling

Vaporization

Process in which a liquid changes to a gas.

Evaporation

Process in which particles escape the surface of

a non-boiling liquid and enter the gas phase.

This is caused by a greater KE at the surface of

the liquid.

Boiling

Conversion of a liquid to a gas within the liquid

as well as at its surface.

10-3 Solids

Definite shape and definite volume.

Crystals

Crystals have an ordered, repeated structure.

The smallest repeating unit in a crystal is a

unit cell.

Three-dimensional stacking of unit cells is the

crystal lattice.

Amorphous Solids

Lack internal order but yet exhibit a solid like

substance.

Jello is similar but it’s considered a colloidal

suspension.

Structures of Solids Unit Cells

Structures of Solids Unit Cells

Structures of Solids Close Packing of Spheres

Close Packing of Spheres

A crystal is built up by placing close packed layers of

spheres on top of each other.

There is only one place for the second layer of

spheres.

There are two choices for the third layer of spheres:

Third layer eclipses the first (ABAB arrangement). This

is called hexagonal close packing (hcp).

Third layer is in a different position relative to the first

(ABCABC arrangement).

Close Packing of Spheres

Each sphere is surrounded by 12 other

spheres (6 in one plane, 3 above and 3

below).

Coordination number: the number of spheres

directly surrounding a central sphere.

Crystal Bonding

Metallic Solids (mobile valence electrons)

Low to High melting points

Metallic bonds hold the particles together

Molecular Solids (lowest melting pts)

Low melting points

Intermolecular forces hold the particles together

Ionic Solids (hard, brittle and non-conducting)

High melting points

Strong electrostatic force of attraction

Covalent – Network Solids (strong covalent bonds between neighboring atoms)

High melting points

Atoms covalently bonded to the same type of atoms

10-4 Changes of State

Phase change

Conversion of a substance from one of the 3

physical states of matter to the other.

Always involves a change in energy.

Equilibrium

Equilibrium ()

Dynamic condition in which 2 opposing

changes occur at equal rates in a

closed system.

Components under equilibrium

Phase – any part of the system that has

uniform composition and properties.

System – sample of matter being studied.

Concentration - #particles per unit of

volume

Evaporation/Condensation

Evaporation – rate in which a liquid changes to a gas under its boiling point.

Condensation – rate in which a gas changes to a liquid.

Phase change : Evaporation Condensation

Liquid + Heat Vapor

Vapor Liquid + Heat

Freezing/Melting

Freezing – rate in which a liquid changes to a solid.

Melting – rate in which a solid changes to a liquid.

Phase change : Freezing Melting

Solid + Heat Liquid

Liquid Solid + Heat

Phase Change

Phase change that occurs when

a solid changes to a gas without

passing through the liquid phase.

Sublimation/Deposition

Possible Changes of State

Changes of State Name Example

Gas Liquid Condensation H2O(g) H2O(l)

Liquid Gas Vaporization Br(l) Br(g)

Liquid Solid Freezing H2O(l) H2O(s)

Solid Liquid Melting H2O(s) H2O(l)

Solid Gas Sublimation CO2(s) CO2(g)

Conversion of a liquid to a vapor, when the vapor

pressure of the liquid is equal to the atmospheric

pressure.

Vapor Pressure – Amount of pressure caused by

the vapor of a liquid in a closed container.

Boiling Point – Temperature at which a liquid’s

vapor pressure equals the atmospheric pressure.

Normal Boiling Point – Temperature at which a

liquid boils at Standard Pressure.

Boiling

2 Factors that cause boiling:

Lowering the atmospheric pressure, by

placing the liquid in a vacuum.

Increasing the vapor pressure, by

increasing the temperature of the liquid.

Factors Affecting Boiling

Graph of Temperature vs. Pressure that

indicates points in which a substance will be

a gas, liquid or a solid.

Triple Point – Temperature and Pressure at

which a substance has all three phases at

equilibrium.

Critical Point – Point in which a substance

can’t exist in the liquid state.

Phase Diagrams

Phase Graph of Water

Phase Diagram

Molar Heat of Fusion/Solidification

Amount of heat needed to change 1 mole of a

substance from a liquid to a solid or solid to a

liquid.

Solid Liquid (Molar Heat of Fusion)

Liquid Solid (Molar Heat of Solid)

Water: Molar Heat of Fusion

(Hfus) = 6.01 kJ/mol

Molar Heats (Enthalpy)

Molar Heat of Condensation/Vaporization

Amount of heat needed to change 1 mole of a

substance from a liquid to a gas or a gas to a

liquid.

Gas Liquid (Molar Heat of Condensation)

Liquid Gas (Molar Heat of Vaporization)

Water: Molar Heat of Vaporization

(Hvap) = 40.7 kJ/mol

Molar Heats Cont.

Molar Heat Problem

Determine the amount of heat needed to melt

100g of ice at 0oC.

Determine the amount of heat needed to

change 100g of liquid water to steam.

10-5 Water

Water is present in a large abundance

throughout our life.

70%-75% earth’s surface is water

60%-90% of the mass of most living things is

water.

Water’s the Exception

Water expands

when it freezes

Less dense than

water

Reason for ice

floating

3.98oC water

begins to expand

due to crystal

formation in

water.

Water’s Hydrogen Bonding

Two Types of Strong Interactions

Cohesive forces – between molecules of the

same type

Adhesive forces – between different types

Adhesive & Cohesive

Water in a tube exhibits a curved surface called a meniscus.

Adhesive forces occur between the glass and water

Drawing the water up along the glass

Cohesive forces in the water help hold the curve in the water level