Chapter 12 Liquids and Solids


Least common state of matter in the universe -- narrow range of temperatures and pressures

Properties of Liquids and the Kinetic-Molecular Theory

a) definite volume
b) takes the shape of its container
c) particles in constant motion

in liquid the particles are closer than in gases and have more effective attractive forces

in liquid the particles have lower mobility because of the closeness of adjacent molecules

fluid - definition

flowing downhill
flowing uphill e.g. helium

Relatively High Density

most liquids are thousands of times denser than their gases at normal temperature and pressure

because of the closeness of the particles

water is an exception of most substances being more dense as a solid than as a liquid

at same temperature, different liquids can vary greatly in density e.g. Figure 12-1 page 364

Relative Incompressibility

at 20 degrees Celsius and at a pressure of 1000 atm pressure, the volume of water decreases by only 4% -- typical of liquids as well as solids -- due to closeness of particles in both liquids and solids

because of the closeness of liquid molecules when pressure is exerted on them, that pressure is transmitted in all directions

Ability to Diffuse

Liquids, like gases, diffuse and mix with other liquids.

Figure 12-2 page 364

Diffusion is due to constant random motion of particles, but much slower in liquids than in gases because of the closeness of particles.

Also slower because of the attractive forces present in liquids.

At higher temperature, the diffusion of liquids increases because of the increased speed of the particles.

Surface Tension

Surface tension - definition

Due to attractive forces of particles and is directly proportional to the attractive forces.

Water has a higher surface tension than many liquids because of its hydrogen bonding.

Molecules at the surface act differently than molecules in the body of the liquid.

Surface tension is why liquid drops have a spherical shape -- a sphere has the smallest surface area for a given volume.

Figure 12-3 page 365

Capillary action - definition

Closely related to surface tension.

Depends on attraction between liquid and the surface of the solid, usually a tube of some kind.

Continues until capillary action is balanced by the pull of gravity.

Can happen between water molecules and paper fibers such as in Figure 12-4 page 365

Movement of water in plants.


Evaporation and Boiling

Vaporization - definition

Evaporation is a form of vaporization.

Evaporation - definition

Liquid bromine added to bottle, and soon have vapor above the liquid.

Figure 12-5 page 366

perfume to wrist area

Each particle in a liquid has a different kinetic energy; a molecule with higher than average kinetic energy that is also at the surface of the liquid can overcome the attractive forces of the particles beside and beneath it and escape from the liquid and goes into the gaseous state.

Important in nature - evaporation of sea water leaves behind the salt which increases the concentration of the salt in the sea.

Evaporation of water from the surface of the earth is what comes back down as rain and snow.

Perspiration -- keeps us cool by removing high energy water molecules from our skin which takes heat away from our body.

Boiling is when a liquid turns into a vapor in the body of the liquid and forms a vapor bubble in the body of the liquid. More on boiling later.

Formation of Solids

Cooling a liquid sufficiently allows attractive forces to pull the particles into an even more orderly arrangement.

Freezing - definition

All liquids freeze e.g. water at 0 degrees Celsius; paraffin (candle wax) at room temperature; ethanol at -115 degrees Celsius.

Homework: 12.1


definite shape
definite volume

Properties of Solids and the Kinetic-Molecular Theory

particles more closely packed than liquid -- thus stronger intermolecular forces (dipole-dipole, London dispersion forces, hydrogen bonding)

particles in relatively fixed positions

only vibrational movement

particles more orderly than other physical states

Figure 12-6 page 367 -- physical appearance of solids, liquids, gases

Two types of solids: a) crystalline solids which consist of crystals - definition; b) amorphous solids - definition

Definite Shape and Volume

maintain definite shape without a container

crystalline solids have regular geometric arrangement of particles reflecting of the internal structure

amorphous solids have a definite shape but not the regular geometric arrangement of particles

volume of solids change little with changes in temperature or pressure

solids have a definite volume because the particles are packed closely an have very little empty space, thus little chance of compression

Definite Melting Point

Melting - definition

Melting Point - definition

When a solid melts the kinetic energy of the particles overcomes the attractive forces holding the particles together

Crystalline solids have a definite melting point. Amorphous solids have a temperature range which is their melting point.

Amorphous solids are sometimes called supercooled liquids.

Supercooled liquids - definition

Particles in supercooled liquids are arranged randomly as in a liquid but they are not true liquids because the particles are not constantly changing their positions.

High Density and Incompressibility

Solids are slightly denser than liquids.

solid hydrogen is the least dense solid and osmium is the most dense

less compressible than liquids

wood and cork are compressible because of air trapped in pores and it is the pores that are compressed

Low Rate of Diffusion

diffusion can occur in solids but only over long periods of time and is a million times slower than diffusion in liquids

Crystalline Solids

can be a single crystals or groups of crystals fused together

crystal structure - definition

lattice - definition

unit cell - definition

Each crystal lattice contains many unit cells packed together - figure 12-7 page 368

A crystal and its unit cells can have any one of seven type of symmetry and allows us to classify crystal by their shape -- figure 12-8 page 369

No need to memorize these crystal symmetries.

Binding Forces in Crystals

can also describe crystals in terms of the types of particles and the types of chemical bonding between particles -- table 12-1 page 370

Four type of crystals based on this:

a) ionic crystals: particles are positive and negative ions; ions can be monatomic or polyatomic; usually group 1 or 2 elements with group 17 or 17 elements; binding forces are attractive forces between positive and negative ions; properties are hard, brittle, high melting points, good insulators.

b) covalent network crystals: particles are single atoms; binding forces are covalent bonds; e.g. diamond, Cx, quartz, (SiO2)x, Figure 12-9, page 371, silicon carbide, (SiC)x, and many oxides of transition metals.

hard and brittle, high melting points, usually nonconductors or semiconductors

c) metallic crystals: particles are metal atom kernel and sea of valence electrons;

high electrical conductivity, melting points vary -- table 12-1 page 370

d) covalent molecular crystals: particles are covalently bonded molecules held by intermolecular attractive forces; if non polar, then only London dispersion forces hold the molecules together, e.g. hydrogen molecule, methane, benzene; in polar covalent molecular crystals the molecules are held together by London dispersion forces, dipole-dipole forces and sometimes hydrogen bonds e.g. Water, ammonia;

low melting points; easily vaporized, relatively soft, good insulators.

Amorphous Solids

without shape

do not have a regular, natural shape as in crystals

hold their shape for long time but some can flow e.g. very old window glass and other types of glass.

Homework: 12.2

Changes of State

table 12-2, page 372 different changes in states and their names


Equilibrium - definition

a closed system means matter can not enter or leave but energy can

Equilibrium and Changes of State

figure 12-10 page 373

initially just a single phase - liquid

phase - definition

surface molecules evaporate and enter gaseous state above the liquid

some of the vapor molecules then will enter the liquid state

condensation - definition

initially there are no vapor molecules, only liquid molecules

as time goes on the number of vapor molecules increases and eventually some of the vapor molecules enter the liquid state figure 12-10 b

at some point in time the number of liquid molecules entering the vapor state is the same as the number of vapor molecules entering the liquid state -- then we have equilibrium. figure 12-10 c

the process of vapor = liquid continues once equilibrium is reached

Equilibrium is based on the RATE of vapor molecules becoming liquid equal to the rate of liquid molecules becoming vapor -- not on the amounts of vapor and liquid at equilibrium

once equilibrium is reached and maintained, the amount of vapor and liquid remains the same because the rate of conversion from one to the other is the same

An Equilibrium Equation

liquid + heat ---> vapor

vapor ---> liquid + heat

liquid + heat = vapor

two opposing arrows or equal sign

forward reaction -- left to right

reverse reaction -- right to left

LeChatelier's Principle

LeChatelier's Principle - statement

a stress can be a change in concentration, pressure, temperature

Equilibrium and Temperature

given: liquid + heat = vapor

and the temperature is raised from 25 to 35 degrees Celsius

fwd reaction is endothermic, reverse is exothermic; an increase in temperature favors the endothermic reaction so the equilibrium shifts to the RIGHT to relieve the stress which means that the fwd reaction occurs faster than the reverse because of the increase in temperature; eventually the reverse reaction catches up and a new equilibrium is reached with both the fwd and rev rxns occurring at a higher rate than at 25 degrees Celsius

at the higher temperature, the concentration of vapor (the product of the favored reaction) is higher than it was at the lower temperature -- concentration here refers to the amount of vapor at the new equilibrium compared to the volume of the container

take original equilibrium and lower the temperature from 25 to 15 degrees Celsius -- describe what happens

Equilibrium and Concentration

given: liquid + heat = vapor

increase the volume available to the system: more volume means the same number of molecules of vapor in more volume -- decreased concentration of vapor molecules -- the change from vapor to liquid slows while the change from liquid to vapor remains the same i.e. the rate of liquid to vapor is now higher than the rate of vapor to liquid -- we say that the fwd reaction is favored.

table 12-3 page 375

Equilibrium Vapor Press of a Liquid

equilibrium vapor pressure - definition

figure 12-11 page 376 - vapor pressure setting up

figure 12-12 page 377 - vapor pressure curves

equilibrium vapor pressure depends on
a) what the liquid is
b) temperature

notice from the graph that an increase in temperature does not increase the vapor pressure proportionally -- not a straight line.

Equilibrium Vapor Pressure and the Kinetic Molecular Theory

increasing the temperature of a liquid increases its average kinetic energy

an increase in average kinetic energy increases the number of molecules that can escape from the liquid to the vapor state

an increase in the number of vapor molecules increases the vapor pressure

Volatile and Nonvolatile Liquids

the stronger the attractive forces between the molecules in a liquid the lower will be its vapor pressure at any given temperature; the weaker the forces of attraction, the high will be its vapor pressure at any given temperature

volatile liquids - definition


boiling - definition

boiling point - definition

the lower the atmospheric pressure the lower the boiling point - affects cooking

during boiling the temperature is constant -- i.e. as long as liquid and vapor coexist the temperature is constant

heating curve figure

heating curve - temperature vs heat

heat absorbed during boiling goes toward converting the liquid molecules into the vapor state i.e. to overcome attractive forces holding the molecules in the liquid state not toward increasing the average kinetic energy

increasing the pressure above a boiling liquid -- pressure cooker

vacuum evaporator -- used to remove water from milk and sugar solutions without affecting the milk or sugar because of the lower temperatures

normal boiling point - definition

figure 12-12 page 377

Energy and Boiling

To keep a liquid boiling you must continually supply heat.

We continue to heat, the liquid continues to boil but its temperature remains constant.

heating curve figure

Heat added during boiling goes to overcome attractive forces between molecules so they can enter the gaseous state.

Molar Heat of Vaporization

molar heat of vaporization - definition

related to the attractive forces in the liquid - greater attractive forces, greater the size of the molar heat of vaporization

characteristic of the liquid

water's is high because of the hydrogen bonding

makes water a good cooling agent

figure 12-13 page 379

Freezing and Melting

freezing involves the loss of heat energy by the liquid and can be represented by

liquid ---> solid + heat energy

normal freezing point - definition

liquid and solid are in equilibrium at the freezing point - heating curve - constant temperature

at the freezing point, the particles of the liquid and the solid have the same average kinetic energy so energy loss during freezing is a loss of potential energy

as long as solid and liquid are present the temperature remains the same

Molar Heat of Fusion

molar heat of fusion - definition

heat absorbed increases the potential energy of the solid as its particles are pulled apart, overcoming the attractive forces holding them together

size depends on the attraction between the particles of the solid

Sublimation and Deposition

at the proper low temperature and pressure a liquid cannot exist and a solid exists in equilibrium with its vapor

solid + heat energy = vapor

sublimation - definition

deposition - definition

e.g. dry ice - solid carbon dioxide; iodine; ordinary ice sublimes slowly at temperatures lower than its melting point - explains how a thin layer of snow can eventually disappear, even if the temperature remains below zero degrees Celsius.

formation of frost on a cold surface is an example of deposition

Phase Diagrams

phase diagram - definition

also shows how the states of a system change with changing temperature or pressure

Figure 12-14 page 381

three curves show equilibrium positions for various physical states

AB: solid (ice) - vapor

AC: liquid - vapor

AD: liquid - solid

notice negative slope of AD because ice is less dense than liquid water - increase in pressure lowers the melting point; most substances have their liquid solid curve with a positive slope

Point A is the triple point

triple point - definition

for water is it 0.01 degrees Celsius

Point C is the critical point - for water it is 373.99 degrees Celsius

critical point - definition

critical temperature - definition

critical pressure - definition

AD has a slope that indicates that ice melts at a higher temperature with decreasing pressure.

Below the triple point, the temperature of sublimation decreases with decreasing pressure.

Figure 12-15 summarizes all changes of states.

Homework: 12.3


covers about 75% of earth's surface
70-90% of the mass of living things is water

Structure of Water

polar covalent bonds between the oxygen and each hydrogen

bent geometry

bond angle 105o

molecules held together by hydrogen bonding

the higher the temperature the fewer molecules that are grouped together

other molecules similar in size and mass to water, if they are non polar, they are gases at room temperature, viz methane, CH4; hydrogen bonding in water

ice has its molecules arranged in a hexagonal structure due to rigid hydrogen bonds; the empty spaces created is why the density of ice is less than liquid water and ice floats on water

from zero to 4 degrees Celsius, heating causes the molecules to collapse into the open spaces; above four degrees the kinetic energy of the molecules due to the heating causes the molecules to start moving apart; temperature of maximum density of water is 4 degrees Celsius

Physical Properties of Water

almost colorless
normal freezing point is 0 degrees Celsius
normal boiling point is 100 degrees Celsius
molar heat of fusion of ice is 6.009 kJ/mol - relatively large
molar heat of vaporization is 40.79 kJ/mol

both are high because of hydrogen bonding

steam heat

Sample Problem 12-1, page 386

sample problem 12-1

Homework: 12.4

end of notes

A fluid is a substance that can flow.

Surface tension is a force that tends to pull adjacent parts of a liquid's surface together, thereby decreasing surface area to the smallest possible size.

Capillary action is the attraction of the surface of a liquid to the surface of a solid.

Vaporization is the process by which a liquid or solid changes to a gas. Back

Evaporation is the process by which particles escape from the surface of a nonboiling liquid and enter the gaseous state.

Freezing is the physical change of a liquid to a solid by removal of heat.

An amorphous solid is one in which the particles are arranged randomly.

A crystal is a substance in which the particles are arranged in an orderly, geometric, repeating pattern.

Melting is the physical change of a solid to a liquid by the addition of heat.

The melting point is the temperature at which a solid becomes a liquid.

Supercooled liquids are substances that retain certain liquid properties even at temperatures at which they appear to be solid.

Crystal structure is the total three-dimensional arrangement of particles of a crystal.

Lattice is a coordinate system used to represent the arrangement of particles in a crystal.

Unit cell is the smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice.

Equilibrium is a dynamic condition in which two opposing changes occur at equal rates in a closed system.

A phase is any part of a system that has uniform composition and properties.

Condensation is the process by which a gas changes to a liquid. Back

When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position that minimizes the stress. Back

The equilibrium vapor pressure is the pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature. Back

A volatile liquid is a liquid that evaporates readily. Back

Boiling is the conversion of a liquid to a vapor within the liquid as well as at its surface and occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure. Back

The boiling point is the temperature at which the equilibrium vapor pressure of a liquid equals the atmospheric pressure. Back

The normal boiling point is the temperature at which a liquid will boil when the atmospheric pressure is exactly one atmosphere. Back

Molar heat of vaporization is the heat energy needed to vaporize one mole of liquid at its boiling point. Back

Normal freezing point is the temperature at which the solid and liquid are in equilibrium at 1 atm pressure. Back

Molar heat of fusion is the amount of heat energy required to melt one mole of solid at its melting point. Back

Sublimation is the change of state from a solid directly to a gas. Back

Deposition is the change of state from a gas directly to a solid. Back

A phase diagram is a graph of pressure versus temperature that shows the condition under which the phases of a substance exist. Back

The triple point of a substance indicates the temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium.

The critical point of a substance indicates the critical temperature and critical pressure

The critical temperature (tc) is the temperature above which the substance cannot exist in the liquid state.

The critical pressure (Pc) is the lowest pressure at which the substance can exist as a liquid at the critical temperature.