Phases of Matter

Chapter 7

Early Theories

• Early Greeks discovered 2 basic laws:
  1. Matter is composed of particles
  2. Basic particles are very small
• Democritus called particles atoms from atomos meaning indivisible
• Aristotle taught 4 basic elements: air, fire, earth, water
• Dalton proposed first modern atomic theory - said atoms were solid spheres, each element having different atoms
• Modern atomic theory improved by Thomson, Rutherford, Bohr, and others
• Atoms composed of many smaller particles.

Sizes

• Molecule: smallest chemical unit of substance that has its chemical properties; size: from 3 to >200 angstroms (Å) (1 angstrom = 10^-10m)
• Atom: smallest unit of element that can exist alone or in combinations. .6 - 5 angstroms
• Atomic mass units (u). 1 u = 1/12 mass of carbon 12 atom = 1.66 x 10^-27kg
• Mass number is nearest integer to atomic mass in u

Kinetic Theory

• Molecules of matter are always in motion
• Collisions between particles are completely elastic
• Amount of motion depends on average kinetic energy of molecules (temperature)

Phases

• Phases of matter:
  • solid: atoms close together in regular pattern,
  • liquid: atoms close together but no pattern,
  • gas: atoms spread apart, no pattern
• Vapor: gaseous form of substance usually a liquid at STP
• Fluid: liquid or gas
• Fourth phase is plasma

Forces Between Molecules

• Attractive forces depend on distance
• Forces greatest in solids, less in liquids
• Attractive forces very small in gases
• If atoms forced too close together, repulsive forces occur
• Solids and liquids hard to compress
• Forces are electrical, gravity too weak to be a factor

Forces of Attraction

• Cohesion: attractive force between molecules of the same type; holds solids together
• Adhesion: attractive force between molecules of different types; makes different materials stick together; water on glass, glue on paper

Solid Phase

• Molecules vibrate about fixed positions
• Definite shape and volume
• Crystalline solids have regular pattern of atoms
• Amorphous solids have random arrangement

Properties of Solids

• Ductility: ability to be drawn into a wire; metals typically
• Malleability: ability to be hammered or rolled into flat sheets; metals
• Tensile strength: force per unit area needed to break a solid by pulling and stretching; a measure of cohesion

Elasticity

• The ability to return to original shape when deformed by force
• Elastic limit: where object won't regain shape
• Stress: deforming force divided by area of contact
• Strain: relative amount of deformation that occurs from stress
• Tension or compression produce linear or elongation strain changing length
• elongation strain = change in length/original length (Dl /l )
• When force is from side it's called shear strain
• When force is from all sides, it is volume strain
• Each type of stress/strain involves a constant called the modulus
• For elongation strain, Young's modulus (Y) is used Y = stress/strain
• To find length change, Dl = Fl / YA

Hooke's Law

• Robert Hooke (English, 1600s) found the force on a spring needed to stretch or compress it a distance x: F = kx where k is a spring constant
• Large spring constant means stiff spring
• k is in N/m
• Hooke's law also means strain is directly proportional to stress

Liquid Phase

• Definite volume, no definite shape
• Molecules close together, no definite arrangement, free to move
• Brownian Motion: Small particles suspended in liquid seen to move randomly due to molecular collisions

Properties of Liquids

• Diffusion: penetration of molecules of one type into a mass of molecules of another type; solids & gases diffuse also
• Viscosity: measure of how a liquid flows; low viscosity means rapid flow; dependent on temperature: motor oil

Properties of Liquids

• Meniscus: curved surface of liquid column; if adhesion with container is greater than cohesion, edges rise -- water; if adhesion < cohesion, center rises -- mercury
• Surface Tension: cohesive forces at surface cause appearance of thin "skin"; causes contraction, spherical drops
• Capillarity: elevation or depression of liquids in small diameter tubes; depends on whether liquid wets surface of tube and diameter of tube; due to cohesion and adhesion

Melting

• Change of phase from solid to liquid; reverse process is freezing
• Crystalline solids have definite melting points
• Melting requires breaking of bonds in crystal lattice; energy must be supplied
• As heat energy is added before melting, kinetic energy of molecules increases
• During melting, energy goes to raising potential energy of molecules to liquid phase
• Temperature (avg. kinetic energy) does not rise during melting
• Non-crystalline solids soften, then melt; no definite melting point
• Most solids have less distance between molecules, less potential energy than liquid

Special Properties of Water

• Water expands when it freezes
• Solid structure more open than liquid
• Liquid water contracts when cooled until 4 deg. C when ice crystals start to form; expands when cooled further
• Ice has 1.1 times volume of water
• Sb and Bi also expand when freezing

Effect of Pressure on Freezing Point

• For most substances, increased pressure will raise freezing point since solid takes up less space.
• For water effect of high pressure is opposite
• Regelation: melting under pressure, refreezing after pressure release.

Effect of Solutes on Freezing Point

• Freezing point lowered by dissolving substance in liquid
• Amount of change depends on liquid and solute
• Example: salt on icy roads

Properties of Gases

• No definite volume or shape
• olecules independent of each other
• exert pressure due to inertia of molecules
• diffuse readily into other gases
• also diffuse into liquids and solids to a lesser degree

Vaporization

• production of gas from another phase
• evaporation: liquid to gas
• sublimation: solid to gas (no liquid phase)
• due to above average kinetic energy from molecular collisions
• cools original liquid or solid because more energetic molecules leave
• rate depends on temperature

Vapor Pressure

• pressure due to molecules from vaporized substance
• increases with temperature
• vaporization and condensation occur simultaneously
• When rates of each are equal, equilibrium occurs and equilibrium vapor pressure exists
• Ratio of current vapor pressure to equilibrium vapor pressure (at current temp.) is relative humidity
• Temperature when current amount of vapor pressure will equal equilibrium vapor pressure is dew point

Boiling

• Rapid vaporization when vapor pressure of liquid equals pressure on its surface
• Temperature when this occurs at 1 atm pressure is normal boiling point
• Increased pressure over liquid raises boiling point; decreased pressure lowers boiling point
• Temperature doesn't change during boiling even if heat is added
• Dissolved solids raise boiling point
• Dissolved gases lower boiling point

Fluid Dynamics

• Fluids exert buoyant force on objects immersed in them
• Buoyant force equals weight of displaced fluid: Archimedes' principle
• Smooth flow of fluid is streamlined flow, shown by drawing streamlines
• Turbulence disrupts smooth flow

Pascal's Principle

• A change in pressure applied to enclosed fluid is transmitted undiminished to every portion of the fluid and to the walls of the container
• Principle behind hydraulic systems
• Can be used to multiply forces with hydraulic lever

Bernoulli's Principle

• For flow of fluid through tube, sum of pressure and kinetic energy per unit volume of fluid is constant
• Faster flow means lower pressure
• Airplane wings show Bernoulli's principle: lift created by low pressure area over wing where air travels faster

Fluid Friction

• Moving through fluid creates friction called drag
• Can be reduced by streamlining: cars, boats, airplanes, rockets, etc.
• Drag depends on density of fluid and cross-sectional area of object