Chapter 12 Liquids and Solids
Liquids
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.
Meniscus
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
Solids
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
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
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
Water
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
transparent
odorless
tasteless
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
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.