CHAPTER 12 LIQUIDS and SOLIDS

Liquids – the least common state of matter in the universe Liquids exist at relatively narrow ranges of temperature and pressure

Liquid Particles have less KE than gas particles Properties of Liquids Particles are in constant motion Particles are much closer together than gases Liquid Particles have less KE than gas particles Attractive forces between particles are strong enough they can not be ignored (dipole-dipole, London dispersion and hydrogen bonds) Movement is more limited than in gases due to these attractive forces

Properties of Liquids Definite Volume- volume can only be changed slightly by temp & pressure. Fluidity- particles have ability to flow past one another. Relatively High Density – most are thousands of times denser as liquids than as gases (solids are only slightly more dense than liquids). Relatively Incompressible – even at very high pressure and low temperature liquids compress only slightly because the particles are already so close together.

Ability to diffuse- particles will mix uniformly with one another but at a much slower rate than gases. Surface Tension– A force that tends to pull adjacent parts of a liquid’s surface together due to attractive forces between particles. Capillary Action- The attraction of the surface of a liquid to the surface of a solid Vaporization - Liquid or solid changing to a gas, evaporation is the change of state from a non-boiling substance. Boiling- Change of a liquid to bubbles of vapor (different from vaporization though)

Liquid – Molecules have Enough Kinetic Energy to Slide Past One Another.

Formation of Solids: When KE is low enough that attractive forces hold particles in fixed positions, the particles become solid. Freezing/solidification- the physical change from a liquid to solid by removing heat.

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SOLID AS A ROCK SOLIDS

Solids have less KE than either liquids or gases. Molecules are tightly packed and intermolecular forces are strongly exerted. Particles are held in a relatively fixed position. Two types of Solids: Crystalline and Amorphous

Crystalline solids: consist of crystals (substance in which the particles are arranged in an orderly, geometric, repeating pattern Amorphous solids: one in which the particles are arranged randomly

PROPERTIES of SOLIDS Definite shape and volume Definite melting point Melting: solid to liquid by addition of heat Melting point: temp at which a solid becomes a liquid Supercooled liquids: substances that retain certain liquid properties even at temperature at which they appear to be solid. (Amorphous solids.. Might they be called liquids that do not flow?)

PROPERTIES OF SOLIDS High density and incompressibility- particles are more densely packed Very Very Low rate of diffusion- Rate of diffusion depends upon: ____________ Temperature: As temperature increases, the kinetic energy of the particles increases so the speed of particles also increases which thus increases the rate of diffusion. Mass of the particle: As the mass of particle increases, rate of diffusion decreases.

Crystalline Solids Crystal structure: total three-dimensional arrangement of particles of a crystal. Unit cell: smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice. The unit cells are then repeated to make the crystal. There are seven basic crystalline systems. See Figure 12-8 page 369

Binding forces in Crystals ionic crystals – Positive and negative ions arranged in a regular pattern. Ex: NaCl, KOH,MgF2 Properties include: hard, brittle, high melting points, good insulators. covalent network crystals – Each atom is covalently bonded to its neighboring atom. Essentially giant molecules. Ex: diamond, SiO2, SiC Properties include: high melting points, nonconductors or semiconductors, hard and brittle. metallic crystals – Sea of electrons model. Electrons belong to the entire crystal. Properties: high conductivity, melting points vary greatly covalent molecular crystals -- Covalently bonded molecules that are held together by intermolecular forces (Van der Waals). Properties: low melting point, soft, good insulators, easily vaporized.

Amorphous Solids Amorphous – Greek.. without shape Do not have a regular natural shape, but can hold their shape for a long time. Solids that flow Ex: Glass, rubber and plastics

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CHANGES of STATE Equilibrium - dynamic condition in which 2 opposing changes occur at equal rates in a closed system. A dynamic equilibrium occurs when you have a reversible reaction in a closed system. Nothing can be added to the system or taken away from it apart from energy. In other words, the system must be closed. Phase: any part of a system that has uniform composition and properties. Ex: liquid phase, gas phase, solid phase.

Trapped molecules reach a balance between evaporation & condensation Equilibrium Trapped molecules reach a balance between evaporation & condensation An increase in condensation = An increase in evaporation

Examples of Changes of State Change of State Process Example solid  liquid melting ice  water solid  gas sublimation dry ice  CO2 gas liquid  solid freezing water  ice liquid  gas vaporization liquid Br  Br vapor gas  liquid condensation water vapor  water gas  solid deposition water vapor  ice

A system will remain at equilibrium until something occurs to change this condition. Le Chatelier’s principle- when a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium. Stressors include changes in: 1. concentration 2. pressure 3. temperature A change in equilibrium is called a shift in equilibrium.

Equilibrium Vapor Pressure Equilibrium Vapor Pressure (EVP) -The pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature. Equilibrium vapor pressure of a liquid increases as temperature increases but not in direct proportion Increasing temp of liquid increases average kinetic energy; thus, increasing the number of molecules that have enough energy to escape from the liquid phase Increased evaporation rate increases the concentration of molecules in the vapor phase increasing the equilibrium vapor pressure. Every liquid has a specific EVP at a given temp (like having specific b.p. & f.p.)

Volatile & Nonvolatile liquids Volatile Liquids- evaporate readily, have weak intermolecular forces. Ex: ether, alcohol, ammonia. A substance with a high equilibrium vapor pressure at normal temperatures is often referred to as volatile. Nonvolatile Liquids – evaporate slowly due to stronger bonding.

Molar Heat of Vaporization Boiling - conversion of a liquid to a vapor within the liquid, as well as at its surface. boiling point (b.p.) of a liquid is the temp at which EVP equals the atmospheric pressure. Boiling occurs when EVP = atm. pressure Boiling point is dependent on atmospheric pressure. The lower the atmospheric pressure is, the lower the boiling point is. At the boiling point, all of the heat absorbed goes to evaporate the liquid Molar Heat of Vaporization The amount of heat that is required to vaporize one mole of liquid at its boiling point. Water has an unusually high heat of vaporization

Freezing and Melting Molar Heat of Fusion Freezing Point-Melting point The temperature at which the solid and liquid are at equilibrium at 1 atm pressure. At this point the solid and liquid have the same average KE. Molar Heat of Fusion The amount of heat that is required to melt one mole of solid at its melting point. The structural order of the particles decrease in as energy is added.

Sublimation and Deposition sublimation - change of a solid directly to a gas (dry ice) deposition - change of a gas directly to a solid (formation of frost on a cold surface) Dry ice iodine

During any change of state, the temp will remain constant until the change of state is complete. water melts at 0oC and freezes at 0oC a glass of ice water at 0oC will always have the same amt. of ice and water! even if heat is added to the system, the water/ice solution will remain at 0oC until all the ice is melted, then its temp will increase. Same goes for boiling

Phase Diagrams Pressure vs. Temperature Triple point - indicates the temp and pressure conditions at which a solid, liquid and gas can coexist at equilibrium. Critical temperature (tc) - the temp above which the substance cannot exist in liquid state. Critical pressure (Pc) - the lowest pressure required for a substance to exist as a liquid at critical temp Pc of H2O = 217.75 atm Pressure vs. Temperature Indicates the conditions under which gaseous, liquid, and solid phases of a particular substance exist. (pg 381) Along AB line: rate at which solid sublimes to form a gas = rate at which gas condenses to form a solid Along BC line: rate at which liquid boils to form a gas = rate at which gas condenses to form a liquid Along BD line: rate at which solid melts to form a liquid = rate at which liquid freezes to form a solid Phase Diagrams

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