1. Unit 9: Liquids and Solids
General Chemistry

2. States of Matter Kinetic-Molecular Theory
particles are always moving
Collisions happen – they are elastic – no energy lost
Gas: indefinite shape and volume
Movement is random, particles are independent of each other
Travel in straight lines until collision occurs
Take shape and volume of container
No intermolecular forces present (IM forces)
Liquid: definite volume, indefinite shape
Particles vibrate around a moving point, motion is limited, particles are closer together
Can slip past each other (in clumps)
Takes shape of container
Some intermolecular forces present (dispersion, dipole-dipole, H-bonds – strongest polar force)

3. States of Matter Solid: definite shape and volume
Relatively fixed position – IM forces between all particles, particles vibrate around fixed point
Arrangement of particles has a pattern – lattice (3-d structure)
State at room temperature depends on type of bonds and intermolecular forces
Plasma: matter at temperatures above 5000 oC
Made of electrons and positive ions (cations)
Creates electric and magnetic fields

4. Liquids Properties Higher density – close together Not compressible
Diffusion – particles spread out from area of high concentration to area of low concentration
Surface tension and capillary action (rise)
Surface tension is caused by unequal forces on the
Interior molecule: feels equal force (pull) in all directions from surrounding molecules
Surface molecule: feels pull across and down into liquid
Pull in toward center of liquid
Apparent elasticity of surface
Liquids form spheres when dropped

5. Liquids
Capillary action: attraction of liquid surface to solid surface
Water rises in small diameter tube (crawls up)
Meniscus
Paper towels
Roots

6. Changing State
Melting/freezing happen at the same temperature – difference is energy being gained or lost
Melting: ordered 3-d arrangement (lattice)of solid breaks down, particles start to slide past each other, some IM forces break
Freezing: particles slow down and settle into 3-d ordered arrangement (forms lattice, IM forces form
Vaporization: change into gaseous state, all IM forces break (Condensation is reverse)
Vapor is gaseous state of substance that normally is liquid or solid at room temp.
Evaporation: energetic particles on surface escape as gas
Boiling: bubbles of gas form throughout the liquid, rise and escape
Sublimation: solid changes directly into gas (Deposition is reverse)

7. Equilibrium: two opposing changes happen at the same rate in a closed system
In a closed container of liquid, equilibrium will be reached where as many particles are returning to liquid as are escaping as vapor
Vapor: gaseous state of substance that normally is liquid or solid at room temp.
Equation written with double arrows (can proceed in either direction)
H2O(l) + heat ↔ H2O(g)

8. Le Chatelier’s Principle: (applies to any equilibrium)
When a system is stressed, it readjusts to reduce the stress, coming to a new equilibrium (it tries to fix what has been done)
Possible stressors: change in temperature, concentration, pressure
Ice skating: skate on a thin layer of water produced because of pressure on the ice (causes it to melt to get molecules closer together)

9. What would happen if we increase temperature on a water/vapor equilibrium?
Adding heat would cause more vapor to form initially until a new equilibrium is reached and the rates of vapor formation are the same.
There would be more H2O molecules present as vapor is in the air than before and a little less liquid water than before.

10. Melting/Boiling Points
M.P.: temperature at which the vapor pressure of the solid equals the vapor pressure of the liquid
B.P.: temperature at which the vapor pressure above the liquid equals atmospheric pressure – changes with altitude
Volatile substance: has a low boiling point so particles evaporate quickly (only has small attractive forces such as dispersion)

11. Solids
Crystals: all true solids are composed of crystals – rigid structure with particles arranged in a repeating pattern
Pattern is determined by bond type and particle size
Classified by external shape: angles and length of axes
Cubic: salt, alum, iron pyrite
Hexagonal: quartz, ice (& snowflakes)
Made of unit cells: simplest repeating unit in crystal
3-d structure is called a lattice

12. Lattice Types: comparison chart
Ionic
Metallic
Molecular
Covalent Network
Lattice positions
ions
atoms
molecules
Force holding lattice together
Electrostatic attraction of positive to negative (ionic bond)
Delocalized electrons (metallic bond)
Intermole-cular forces (dispersion, dipole, H bonds)
Covalent bonds between all atoms
Melting points
high
low
Very high
Examples:
NaCl, CaCO3
Cu, Ag, Au
CO2, H2O
Diamond, silicon carbide

13. Some substances have more than one shape
Can be created by different temperatures or pressures during formation Example: Carbon
Graphite: arranged in layers (covalent) but between layers only has dispersion forces –easily broken so layers slide over each other
Diamond: one huge molecule as all carbons are covalently bonded

14. Amorphous solids Without crystalline form (no pattern)
No definite melting point: gradually softens
Often classed as a supercooled liquid (no crystals form as cooled)
Examples: glass, butter

15. Liquid Crystals
Lose order in one or more dimensions at its melting point (stays in layers or parallel arrangement like soldiers)
Some we use show the ability to change colors or go from transparent to opaque under certain conditions
Current: LCD displays (calculators, cell phones, TVs, ipod)
Temperature: mood rings, thermometers, stress cards

16. Phase Diagrams
Show physical states of matter for a substance at specific temperature and pressure
Lines in between are equilibrium points between the two states (both states exist in equilibrium – state change happens here)
See water diagram – label the
triple point, critical point,
normal melting point, normal
boiling point, & also the phase
changes occuring along each line

17. Water Unique properties are a result of hydrogen bonding
Liquid at room temp. (so small should be a gas)
Solid is less dense than liquid – hydrogen bonds force an open hexagonal pattern so molecules are farther apart(ice floats on liquid water)
High surface tension and capillary rise
High boiling and critical points (higher than expected for its size and bonds)
Water reaches maximum density at 4 oC.