States of Matter Newport High School Academic Chemistry Modified from a PowerPoint found at

States of MatterStates of MatterStates of MatterStates of Matter

Newport High SchoolAcademic ChemistryNewport High SchoolAcademic Chemistry

Modified from a PowerPoint found at http://www.nisd.net/communicationsarts/pages/chem/

Lesson 1 – Kinetic Molecular TheoryLesson 1 – Kinetic Molecular Theory

Essential Questions:• How does the kinetic-molecular theory

describe the properties of gases that include: expansion, fluidity, low density, compressibility, diffusion, and effusion?

Vocabulary: kinetic-molecular theory, ideal gas, elastic collision, diffusion, effusion, real gas

Lesson 1 – The Kinetic Molecular Theory

Lesson 1 – The Kinetic Molecular Theory

Kinetic-molecular theory is based on the idea that particles of matter are always in motion.

Can be applied to solids, liquids, and gases.

Assumptions of GasesAssumptions of Gases

Ideal gas – one that behaves in a way to fit all 5 assumptions of the kinetic-molecular theory of gases.

5 Assumptions• Gases consist of large numbers of tiny particles that

are far apart.• Collisions between particles are elastic.• Gas particles are in continuous, rapid, random motion.• There are no forces of attraction between gas particles.• The temperature of a gas depends on the average

kinetic energy of the particles of the gas.

ExpansionExpansion

Gases do not have a definite shape or volume.

Completely fill any container they occupy.

FluidityFluidity

Gas particles glide past each other because attractive forces are insignificant.

They are considered fluid, because they flow.

Low densityLow density

Density is about 1/1000 the same substance in the liquid or solid state.

Reason is that gas particles are very far apart.

CompressibilityCompressibility

Gas particles become more crowded when compressed.

Volume can be decreased due to large amount of space between particles.

Examples of compression of gases?

Diffusion and EffusionDiffusion and Effusion

Diffusion – spontaneous mixing of the particles of two substances caused by random motion

http://www.indiana.edu/~phys215/lecture/lecnotes/diff.html

Effusion – process where gas particles pass through a tiny opening

Lesson 2 - LiquidsLesson 2 - Liquids

Essential Question:• How does the kinetic-molecular theory of

liquids describe the motion of particles and the properties of liquids?

Vocabulary: Fluid, surface tension, capillary action, vaporization, evaporation, freezing

Kinetic Molecular Theory & Liquids

Kinetic Molecular Theory & Liquids

Liquid = definite volume and takes the shape of its container

Particles are closer togetherAttractive forces do existMore ordered than gasesParticles have lower mobility, but are

fluids

Properties of LiquidsProperties of Liquids

DensityCompressibilityDiffusionSurface TensionCapillary actionEvaporation and BoilingFormation of Solids

DensityDensity

Substances are hundreds of times denser in the liquid state than in the gaseous state

Water is one of the few substances that because less dense when it solidifies. • Most become more dense.

Densities differ so much between liquids that they can form layers.

CompressibilityCompressibility

They are relatively incompressibile.

• Particles are more closely packed.

• Volume only decreases by 4% at very high pressures.

DiffusionDiffusion

Liquids diffuse in the same way as gases.• Due to random motion of particles.• Is slower than gases since particles are closer together

and attractive forces exist.• Increase temperature, increase diffusion

Surface TensionSurface Tension

Surface tension – force that pulls adjacent parts of a liquid’s surface together, consequently decreasing surface area to the smallest possible size• This is why they form a sphere shape

Higher force of attraction = higher surface tension

Capillary actionCapillary action

Capillary action – attraction of the surface of a liquid to the surface of a solid

Attraction tends to pull the liquid molecules upward along the surface – reason for the meniscus

Responsible for the transportation of water in plants.

Evaporation & BoilingEvaporation & Boiling

Vaporization – process by which a liquid changes to a gas• Evaporation – vaporization without boiling• Boiling – energy is added in the form of heat

Formation of SolidsFormation of Solids

When a liquid is cooled, the average energy of the particles decreases.• This is called freezing or solidification.• Each substance has its own freezing

temperature.

Lesson 3 - SolidsLesson 3 - Solids

Essential Question:• How does the kinetic-molecular theory of

solids describe the motion of particles in solids and the properties of solids?

Vocabulary: Crystalline solids, crystal, amorphous solids, melting, melting point, supercooled liquids, crystal structure, unit cell

Kinetic Theory & SolidsKinetic Theory & Solids

Crystalline and amorphous solids Definite shape and Volume

• Shape stays due to arrangement of particles.• Volume only changes slightly with change in temperature.

Definite Melting Point• When particles overcome the forces holding them together.

High Density• Results from the fact that the particles of a solid are more closely packed

than liquids or gases

Incompressibility Low Rate of Diffusion

• Millions of times slower than liquids

Crystalline SolidsCrystalline Solids

Crystal structure – 3D arrangement of particles• Unit cell – smallest portion of a crystal lattice

that shows 3D structure

Four Types of CrystalsFour Types of Crystals

Ionic• Positive and negative ions arranged in a regular pattern

• High melting points, hard and brittle, good conductors

Covalent network• Each atom is covalent bonded to its nearest neighbor.

• Very hard and brittle, high melting point, nonconductors or semiconductors

Metallic• Metal cations surrounded by a sea of valence electrons

• High melting point, good conductors

Covalent molecular crystals• Held together by intermolecular forces

• Low melting points, easily vaporized, soft, insulators

Amorphous SolidsAmorphous Solids

Particles not arranged in an orderly pattern

Most are cooled in ways that do not let them crystallize.

Glass, plastic, rubber, and asphalt are examples.

Have no definite melting point

Lesson 4 – Phase ChangesLesson 4 – Phase Changes

Essential Questions: • How do phase diagrams show the relationship

between the physical states of a substance and its temperature and pressure?

Vocabulary: phase, condensation, equilibrium, vapor pressure, volatile liquid, boiling, boiling point, molar enthalpy of vaporization, freezing point, molar enthalpy of fusion, sublimation, deposition, phase diagram, triple point, critical point, critical temperature, critical pressure

Phase ChangesPhase Changes

Phase ChangesPhase Changes

EvaporationEvaporation• molecules at the surface gain enough

energy to overcome IMF

VolatilityVolatility• measure of evaporation rate• depends on temp & IMF

Phase ChangesPhase Changes

Kinetic Energy

# o

f P

art

icle

s

p. 477

Boltzmann Distribution

temp

volatility

IMF

volatility

Phase ChangesPhase Changes

EquilibriumEquilibrium• trapped molecules reach a balance

between evaporation & condensation

Phase ChangesPhase Changes

Vapor PressureVapor Pressure• pressure of vapor above

a liquid at equilibrium

IMF v.p.temp v.p.

• depends on temp & IMF• directly related to volatility

temp

v.p

.

Phase ChangesPhase Changes

Boiling Point• temp at which v.p. of liquid

equals external pressure

IMF b.p.Patm b.p.

• depends on Patm & IMF

• Normal B.P. - b.p. at 1 atm

Which has a higher m.p.?• polar or nonpolar?• covalent or ionic?

Phase ChangesPhase Changes

Melting Point• equal to freezing point

polar

ionic

IMF m.p.

Phase ChangesPhase Changes

Sublimation

• solid gas

• v.p. of solid equals external pressure

EX: dry ice, mothballs, solid air fresheners

Heating CurvesHeating Curves

Melting - PE

Solid - KE

Liquid - KE

Boiling - PE

Gas - KE

Heating CurvesHeating Curves

Temperature Change• change in KE (molecular motion) • depends on heat capacity

Heat Capacity• energy required to raise the temp of 1

gram of a substance by 1°C• “Volcano” clip - water has a very high

heat capacity

Heating CurvesHeating Curves

Phase Change• change in PE (molecular arrangement)• temp remains constant

Heat of Fusion (Hfus)

• energy required to melt 1 gram of a substance at its m.p.

Heating CurvesHeating Curves

Heat of Vaporization (Hvap)

• energy required to boil 1 gram of a substance at its b.p.

• usually larger than Hfus…why?

EX: sweating, steam burns, the drinking bird

Phase DiagramsPhase Diagrams

• Shows the relationship among the solid, liquid, and vapor states.

• Each region represents a pure phase

• Line between regions is where the two phases exist in equilibrium

• Triple point is where all 3 curves meet, the conditions where all 3 phases exist in equilibrium!

Phase Diagrams cont.Phase Diagrams cont.

Critical temperature – temperature above which a substance cannot be in a liquid state

Critical pressure – lowest pressure at which the substance can exist as a liquid at the critical temperature

Critical point – critical temperature and pressure

Normal freezing point – point at which liquid freezes at sea level

Normal boiling point – point at which liquid boils at sea level

Lesson 5 - WaterLesson 5 - Water

Essential Questions:

• How are the properties of water determined by its structure?

• What happens to the energy of water when it changes state?

Vocabulary: Heat, energy

Structure of WaterStructure of Water

Two hydrogen atoms and one oxygen atom

Molecules in water are linked by hydrogen bonding

Why is water not a gas at room temperature?

Empty space between molecules is why ice has lower density.

Physical Properties of WaterPhysical Properties of Water

Works CitedWorks Cited

Modern Chemistry Textbookwww.nclark.nethttp://mrsj.exofire.net/chem/http://www.unit5.org/chemistry/