Unit 9: Liquids and Solids General Chemistry

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)

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

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

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

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)

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)

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)

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.

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)

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

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

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

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

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

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

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.