Liquids and Solids Chapter 12

Liquid Has a definite volume and indefinite shape Particles are in constant motion Closer together than gases Less kinetic energy than gases Greater attractive forces than gases Chemistry chapter 12

Fluid Substance that can flow and therefore take the shape of its container Chemistry chapter 12

Fluid density At normal pressure, most liquids are thousands of times denser than their gases. Particles are closer together Different liquids can vary greatly in density Chemistry chapter 12

Density Chemistry chapter 12

Incompressibility Liquids are much less compressible than gases Particles are closer together Liquids can transmit pressure throughout themselves Chemistry chapter 12

Diffusion Liquids diffuse in other liquids in which they can dissolve Much slower than gases Particles closer together Attraction between particles slows them down Faster at higher temperatures More kinetic energy Chemistry chapter 12

Surface Tension Force that pulls parts of a liquid’s surface together, causing it to have the smallest possible size From attractive forces between molecules Liquid droplets take a spherical shape Chemistry chapter 12

Capillary action The attraction of the surface of a liquid to the surface of a solid Causes meniscus Water can travel up paper Water traveling up a plant Drawing blood in capillary tube Chemistry chapter 12

Vaporization Changing from a liquid to a gas Evaporation Boiling Higher energy particles escape from the surface of a nonboiling liquid perfume Boiling Bubbles of vapor that appear throughout the liquid and travel to the surface Chemistry chapter 12

Freezing Changing a liquid to a solid by removing heat Energy of particles decreases until they are pulled into a more orderly arrangement Chemistry chapter 12

Discussion Describe the liquid state using kinetic molecular theory. Explain why liquids in a test tube form a meniscus. Compare and contrast vaporization and evaporation. Chemistry chapter 12

Solid Has definite volume and definite shape Particles are in constant motion Much closer together than liquid or gas Much stronger intermolecular forces Held in relatively fixed position – only vibrate Most ordered state of matter Chemistry chapter 12

High density and incompressibility Substances are generally the most dense in the solid state Slightly denser than liquids, much denser than gases Virtually incompressible Sometimes we compress air pockets in the solids Wood, cork, etc. Chemistry chapter 12

Diffusion Very, very slow A few atoms may diffuse if clamped together for a long time Chemistry chapter 12

Melting Change of a solid to a liquid by addition of heat Melting point – temperature at which something melts Chemistry chapter 12

Crystalline solids Consist of crystals Fragments have geometric shapes Particles are arranged in an orderly, geometric, repeating pattern Fragments have geometric shapes Have definite melting points When the crystal structure breaks apart Chemistry chapter 12

Crystal structure Total three-dimensional arrangement of particles in a crystal Chemistry chapter 12

Lattice Coordinate system that represents the arrangement of particles in a crystal. Chemistry chapter 12

Unit cell Smallest portion of a crystal lattice that shows the 3D pattern of the entire lattice Each crystal lattice contains many unit cells packed together Has one of seven types of symmetry – see page 369 Chemistry chapter 12

Ionic crystals Positive and negative ions in a regular pattern Hard and brittle High melting points Good insulators Chemistry chapter 12

Covalent network crystals Individual atoms connected by covalent bonds Giant molecules Diamond Quartz Very hard and brittle Rather high melting points Nonconductors or semiconductors Chemistry chapter 12

Metallic crystals Metal atoms surrounded by a sea of electrons High electrical conductivity Varying melting points Chemistry chapter 12

Covalent molecular crystals Covalently bonded molecules held together by intermolecular forces Low melting points Easily vaporized Relatively soft Good insulators Chemistry chapter 12

Amorphous solids Noncrystalline solids The particles are arranged randomly Glass Plastics Can be molded Fragments have irregular shapes Chemistry chapter 12

Amorphous solids Made by cooling molten substances in a way that prevents crystallization Also called supercooled fluids Retain certain fluid characteristics even at temperatures at which they appear to be solid Can flow over a wide range of temperatures Chemistry chapter 12

Discuss Account for each of the following properties of solids: Definite volume Relatively high density Extremely low rate of diffusion What is the difference between an amorphous solid and a crystalline solid? Chemistry chapter 12

Possible changes of state Melting: solid to liquid Sublimation: solid to gas Freezing: liquid to solid Vaporization: liquid to gas Condensation: gas to liquid Deposition: gas to solid Chemistry chapter 12

Closed system Matter cannot enter or leave, but energy can Chemistry chapter 12

Equilibrium A dynamic condition in which two opposing changes occur at equal rates in a closed system. The same number of particles are entering and leaving. The total number stays the same. Chemistry chapter 12

Equilibrium Chemistry chapter 12

Phase Any part of a system that has uniform composition and properties. Liquid or gas Chemistry chapter 12

An Equilibrium equation When a substance changes state, it either absorbs or gives off energy, usually as heat. Chemistry chapter 12

Le Châtelier’s Principle A system remains in equilibrium until a stress occurs on the system. Stress: change in concentration, pressure, or temperature When a system is disturbed by a stress, it attains a new equilibrium position that minimizes the stress. Chemistry chapter 12

Shifting equilibrium Shifts to the right or left, depending on which part of the equation gains concentration. See table 12-3 on page 375 Chemistry chapter 12

Equilibrium vapor pressure The pressure exerted by a vapor in equilibrium with its liquid at a given temperature Increases as temperature increases But not directly Chemistry chapter 12

Kinetic-molecular theory Increasing the temperature increases the energy and speed of the liquid particles This means more particles evaporate, leading to higher vapor pressure Chemistry chapter 12

Caution Equilibrium vapor pressure depends only on temperature. If the system is not in equilibrium, gas laws must be used. Chemistry chapter 12

Volatile liquids Evaporate easily Weak forces of attraction between particles Ether, acetone Chemistry chapter 12

Boiling Conversion of a liquid to a vapor within the liquid as well as at the surface. Occurs when the equilibrium vapor pressure equals atmospheric pressure. All the heat absorbed goes to evaporate the liquid, so the temperature remains constant. Chemistry chapter 12

Chemistry chapter 12

Cooking If atmospheric pressure is lower (high altitudes), liquids boil at lower temperatures and food takes longer to cook. If the pressure is increased (pressure cooker), liquids boil at higher temperatures and food cooks faster. If the pressure is decreased, it boils at low enough temperatures to avoid scorching milk and sugar. (evaporated and sweetened condensed milk) Chemistry chapter 12

Molar heat of vaporization The amount of energy needed to vaporize one mole of liquid at its boiling point. (or the amount of energy released when one mole of vapor condenses) A measure of the attraction between particles. Chemistry chapter 12

Normal freezing point Temperature at which the solid and liquid are in equilibrium at 1 atm pressure. When a liquid freezes, energy is lost and order is gained. Chemistry chapter 12

Clarification Boiling point is the same as condensation point. Freezing point is the same as melting point. Chemistry chapter 12

Molar heat of fusion The amount of heat energy required to melt one mole of solid at its melting point (or the amount of energy released when one mole of a liquid freezes) Depends on the attraction between particles. Chemistry chapter 12

Phase diagram A graph of pressure versus temperature that shows the conditions under which the phases of a substance exist Reveals how the states of a system change with changes in temperature or pressure Chemistry chapter 12

Water’s phase diagram Chemistry chapter 12

Curves on diagram AB shows where solid and vapor can exist at equilibrium AC shows liquid and vapor at equilibrium AD shows solid and liquid at equilibrium Usually a positive slope, but water is different Chemistry chapter 12

Triple point Point A Shows the temperature and pressure at which solid, liquid, and vapor can exist in equilibrium Chemistry chapter 12

Critical temperature Temperature above which the substance cannot exist in the liquid state. Chemistry chapter 12

Critical pressure The lowest pressure at which the substance can exist as a liquid at the critical temperature Chemistry chapter 12

Critical point Point C Where critical pressure meets critical temperature Chemistry chapter 12

Water’s phase diagram Chemistry chapter 12

Carbon Dioxide’s phase diagram Chemistry chapter 12

Discuss Section review on page 382 Chemistry chapter 12

Water Bent molecule Hydrogen bonding in liquids and solids Bond angle = 105° Hydrogen bonding in liquids and solids Usually 4 – 8 molecules per group in liquid water Without them, water would be a gas at room temperature Ice has hexagonal arrangement Empty spaces lead to low density Chemistry chapter 12

Water Has highest density at 3.98 °C When it melts, molecules can crowd together When it gets hotter, increased kinetic energy makes them spread apart Needs a lot of energy to completely break hydrogen bonds and vaporize Chemistry chapter 12

Water Pure liquid water is transparent, odorless, tasteless, and almost colorless Odors or tastes are caused by impurities Molar heat of fusion: 6.009 kJ/mol Relatively large Density of ice: 0.917 g/cm3 Molar heat of vaporization: 40.79 kJ/mol Quite high Chemistry chapter 12

Calculating heat energy Chemistry chapter 12

Example Find the mass of liquid water required to absorb 5.23 x 104 kJ of heat energy on boiling. 2.31 x 104 g Chemistry chapter 12

You try How much heat energy is absorbed when 16.3 g of ice melts? 5.44 kJ Chemistry chapter 12

You try Calculate the quantity of heat energy released when 783 g of steam condenses. 1.77 x 103 kJ Chemistry chapter 12

Specific Heat The amount of energy required to raise the temperature of one gram of substance by one one kelvin. J/(g∙K) Used to calculate the energy absorbed or released by a substance during a temperature change No change of state Example: warm water  cool water Chemistry chapter 12