Chemistry notes Chapter 12

Section 1 “Liquids”  Properties  Definite volume  Takes the shape of its container  Particles are in constant motion  Adjacent particles are held together with dipole-dipole forces; London dispersion forces and Hydrogen bonding (these forces work on gases too, but not as strongly as they work in liquids)

Intermolecular forces  Dipole-Dipole forces  The force of attraction between two polar molecules

London dispersion Forces  Intermolecular attraction from the constant motion of electrons and the creation of instantaneous dipoles

Hydrogen bonding  Hydrogen atom is bonded to a highly electronegative atom and attracted to another nearby highly electronegative atom in an adjacent molecule

Fluid  Any substance that can flow and therefore take the shape of its container.  Liquids and gases are considered fluids

Density  Most liquids are about 10% less dense than they are as a solid  Water is one of the few substances that becomes less dense as a solid

Pressure  Liquids under pressure only condense to about 4% smaller than original volume. This is very similar in solids. Gases under pressure condense many times more.

Diffusion  Much slower in liquids than in gases  Increased temperature increases diffusion rate

Surface tension  A force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size  This is why liquid droplets form a spherical shape

Capillary action  The attraction of the surface of a liquid to the surface of a solid  This causes the meniscus in a graduated cylinder and how roots transport liquids UP the plant

Phase Changes  Vaporization: the process of a liquid or a solid changing to a gas  Evaporation: a form of vaporization where the particles escape from the surface of a non-boiling liquid  Freezing: the physical change of a liquid to a solid by removing heat (also called solidification)

SECTION 2 “Solids”  Properties of a solid  Intermolecular forces are stronger in solids than in liquids  Molecules are more closely packed together  Molecular motion is restricted to vibrational movement around a specific point

Crystalline solid  Most solids are crystalline  Consists of crystals: orderly, repeating geometric pattern  Crystal structure: total 3 dimensional arrangement of particles  Lattice: the coordinate system that represents the arrangement of particles  Unit cell: the smallest portion of a crystal lattice that shows the 3 dimensional pattern

4 types of crystals  Ionic crystals:  usually a mix of group 1 or 2 with group 16 or 17: usually hard and brittle, high melting point and good insulators  Covalent network crystals:  Usually a single element covalently bonded with itself in a large network of atoms: examples include quartz, diamonds and oxides of transition metals: they are also hard and brittle with high melting points and are usually nonconductors or semiconductors

4 types of crystals (continued)  Metallic crystals  Metal atoms surrounded by a sea of pooled electrons: they are highly conductive: melting points vary greatly  Covalent molecular crystals ( low melting point, relatively soft, good insulators  Polar: water and ammonia held together by all types of intermolecular forces  Nonpolar: hydrogen, methane, benzene are all held together by weak london dispersion

Amorphous solid  “Without Shape”  Glass and plastics are amorphous  Particles are arranged in a random order  Some are said to “flow”: you can see old panes of glass that are thicker at the bottom

Melting Point  Physical change of solid to liquid at a certain temperature  Kinetic energy of particles overcome the forces that are holding them together  Crystalline solids have definite melting points whereas amorphous solids have no definite melting point and can become supercooled liquids

Density  Higher density than liquids and gases  Less compressible than liquids and are usually considered “incompressible”  If you compress a cork, it’s the air pockets inside the cork that is being compressed, not the wood

Diffusion  Diffusion occurs in solids, but only millions of times slower than in liquids  Zinc and copper plates that are compressed together for a long period of time  Rock / sediment that is compressed for hundreds or thousands of years

Section 3 “Changes of State”  Equilibrium: a dynamic condition in which two opposing changes occur at equal rate in a closed system  Condensation: the process by which a gas changes to a liquid  Liquid + heat energy ↔ vapor

Le Chatelier’s principle  When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position that minimizes the stress  Stress is usually a change in concentration, pressure or temperature

How Temperature affects Equilibrium  Increase in temperature favors the forward reaction (endothermic)  Decrease in temperature favors the backwards equation (exothermic)  Liquid + heat energy ↔ vapor

How concentration affects equilibrium  Decrease in vapor concentration (increase in volume) less condensation occurs – equilibrium shifts to the right  Liquid + heat energy ↔ vapor See table 12-3 on page 375

How Vapor pressure affects equilibrium  Equilibrium vapor pressure : the pressure exerted by a vapor in equilibrium with its corresponding liquid at a give temperature  EVP increases with increasing temperature

Volatile liquids  Liquids that evaporate easily  Ethanol is a good example of a volatile liquid. The intermolecular forces are very weak.

Boiling  Occurs when the EVP of the liquid equals the atmospheric pressure  Liquid turns to vapor within the liquid as well as on the surface  Boiling point: temperature at which EVP = atm. Pressure and boiling occurs for given substance  Energy used to separate intermolecular forces of liquids is stored in the vapor as potential energy

Molar Heat of Vaporization  The amount of heat energy needed to vaporize one mole of liquid at its boiling point  The greater the M.H. of V. the stronger the intermolecular forces of the liquid  Water has a very high MHV  Makes it affective as a cooling agent (it absorbs heat away from the surface as it evaporates)

Freezing and Melting  Occur at the same temperature but in reverse directions  Freezing point = The temperature at which the solid and liquid are in equilibrium at standard pressure  Solid + heat energy ↔ liquid

Molar Heat of Fusion  The amount of heat energy needed to melt one mole of solid at its melting point

Sublimation and Deposition  Solid + heat energy ↔ vapor  Low temp and pressure does not allow liquids to exist  Sublimation = change from solid to gas  Deposition = change from gas to solid  Both dry ice and iodine sublime at normal temperatures

Phase Diagram  Triple point : indicate the temperature and pressure condition at which the solid, liquid and vapor of the substance can coexist at equilibrium  Critical point : critical temp and press  Critical temperature : above this temperature, liquid cannot exist  Critical pressure : lowest pressure liquid can exist at the critical temperature

Phase diagrams

Liquid water is more dense than ice Water Ice Water molecules are most dense at 3.98°C

The only math in this section  Molar heat of fusion is kJ/mole  At standard pressure, molar heat of vaporization is kJ/mol 1. a) How much heat energy is absorbed when 47.0g of ice melts at STP? b) How much is absorbed when this same mass of liquid water boils? a) 15.7 kJ b) 106 kJ

The only math in this section  Molar heat of fusion is kJ/mole  At standard pressure, molar heat of vaporization is kJ/mol 2. What quantity of heat energy is released when 506 g of liquid water freezes? 169 kJ

The only math in this section  Molar heat of fusion is kJ/mole  At standard pressure, molar heat of vaporization is kJ/mol 3. What mass of steam is required to release 4.97 x 10 5 kJ of heat energy on condensation? 2.19 x 10 5 g