CHAPTER 10

States of Matter

Kinetic Molecular Theory

◦Kinetikos - “Moving”◦Based on the idea that particles of matter are always in motion ◦ The motion has consequences◦Explains the behavior of Gases, Liquids, and Solids

Ideal Gas – an imaginary gas that conforms perfectly to all assumptions of the KMT

Five Assumptions of the KMT1. Gases consist of large number of tiny particles

that are far apart2. The collisions between particles & w/ the

container wall are elastic (no net loss of KE)3. The Particles are in constant, rapid, random

motion, moving in straight lines.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. Remember: KE = ½ mv2

Atmospheric Pressure◦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.

Barometer Manometer

Physical Properties of Gases

◦Have indefinite shape and indefinite volume.◦Has mass◦Expands to occupy any space available.◦Easily compressed◦Gases exert pressure◦ Fluidity – ability to flow (just like liquids!)◦ Low density

Mixing of Gases

◦Different gases move easily through each other.◦Diffusion – spontaneous mixing of 2 gases.◦Low mass = High rate◦Effusion – gas passes through tiny opening.

“Real” Gas

◦Gas that does not behave completely to the assumption of the KMT.◦Reasons gases deviate from ideal behavior:◦High Pressure◦Low Temperature◦Polar molecules like water and ammonia

Liquids◦Definite volume but no definite shape.◦Liquids are fluids (just like gases!)◦Molecules are held together as 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.◦Result is that the surface acts like an elastic film.

Capillary ActionWater forms a meniscus (curved surface) by using cohesive and adhesive forces. This process is called capillary actionLiquids use cohesive and adhesive forces

to rise within a cylinder

Examples: Oil moving up a lamp wick, water creeping up a towel

Capillary Action

The thinner the tube, the taller the column of water will be!

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.

Solids

◦Definite shape and definite volume◦Definite melting point◦High density◦ Incompressible◦ Low rate of diffusion◦ Still happens, but millions of times slower than liquids

Crystals vs. Amorphous Solids◦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.◦ AKA supercooled liquids◦ Examples: Glass, plastic

Types of Crystals

Structures of SolidsUnit Cells

Structures of SolidsUnit Cells

Structures of SolidsClose 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

Crystal Bonding◦ 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

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

Solid Liquid Gas

Phase Changes

Melting(Add KE)

Vaporization(Add KE)

Freezing (Remove KE)

Condensation(Remove KE)

Sublimation

Deposition

• 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 Changes

• 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.

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 at which a substance can’t exist in the liquid state.

Phase Diagrams

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 (Enthalpy)

Molar Heat Problems◦ Determine the amount of heat needed to melt

100g of ice at 0oC.

100g 1mol 6.01kJx x =33.4 kJ1 18g 1mol

o Determine the amount of heat needed to change 100g of liquid water to steam.

100g 1mol 40.7kJx x =226 kJ1 18g 1mol

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.

◦Because water exhibits hydrogen bonding, it behaves differently than most covalent compounds!

Water’s the Exception

◦ Water expands when it freezes ◦Hydogen bonds force a hexagonal crystal pattern◦ Ice floats (less dense)◦At 3.98oC water begins to expand due to crystal formation in water.