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

Lesson 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  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 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.

Expansion  Gases do not have a definite shape or volume.  Completely fill any container they occupy.

Fluidity  Gas particles glide past each other because attractive forces are insignificant.  They are considered fluid, because they flow.

Low density  Density is about 1/1000 the same substance in the liquid or solid state.  Reason is that gas particles are very far apart.

Compressibility  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 Effusion  Diffusion – spontaneous mixing of the particles of two substances caused by random motion   Effusion – process where gas particles pass through a tiny opening

Lesson 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  Liquid = definite volume and takes the shape of its container  Particles are closer together  Attractive forces do exist  More ordered than gases  Particles have lower mobility, but are fluids

Properties of Liquids  Density  Compressibility  Diffusion  Surface Tension  Capillary action  Evaporation and Boiling  Formation of Solids

Density  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.

Compressibility  They are relatively incompressibile. Particles are more closely packed. Volume only decreases by 4% at very high pressures.

Diffusion  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 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 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 & 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 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 - 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 & 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 Solids  Crystal structure – 3D arrangement of particles Unit cell – smallest portion of a crystal lattice that shows 3D structure

Four 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 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 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 Changes

 Evaporation molecules at the surface gain enough energy to overcome IMF  Volatility measure of evaporation rate depends on temp & IMF

Phase Changes Kinetic Energy # of Particles p. 477 Boltzmann Distribution tempvolatilityIMFvolatility

Phase Changes  Equilibrium trapped molecules reach a balance between evaporation & condensation

Phase Changes  Vapor Pressure pressure of vapor above a liquid at equilibrium IMFv.p.tempv.p. depends on temp & IMF directly related to volatility temp v.p.

Phase Changes  Boiling Point temp at which v.p. of liquid equals external pressure IMFb.p.P atm b.p. depends on P atm & IMF Normal B.P. - b.p. at 1 atm

 Which has a higher m.p.? polar or nonpolar? covalent or ionic? Phase Changes  Melting Point equal to freezing point polar ionic IMFm.p.

Phase Changes  Sublimation solid  gas v.p. of solid equals external pressure  EX: dry ice, mothballs, solid air fresheners

Heating Curves Melting - PE  Solid - KE  Liquid - KE  Boiling - PE  Gas - KE 

Heating 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 Curves  Phase Change change in PE (molecular arrangement) temp remains constant  Heat of Fusion (  H fus ) energy required to melt 1 gram of a substance at its m.p.

Heating Curves  Heat of Vaporization (  H vap ) energy required to boil 1 gram of a substance at its b.p. usually larger than  H fus …why?  EX: sweating, steam burns, the drinking bird

Phase 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.  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 - 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 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 Water

Works Cited  Modern Chemistry Textbook   