Gas Liquid Solid
Entropy
High Medium
Low
Intermolecular Forces
Negligible Moderate High
Spacing
Widely spaced Close
packed Close
packed
Motion
Active Active Fixed
location
Entropy is often described as "disorder", but is also described as "possible states". High entropy has a high number of possible states, this makes predictions more difficult, hence the term "disorder".
Physical state is a result of a balance between the kinetic energy of the particles and the intermolecular forces.
Intermolecular Forces (forces that occur between molecules) occur in the condensed states (liquids and solids)
Intermolecular forces are also called cohesive forces
Nonpolar molecules also have attractive forces between them. These are called London Dispersion Forces.
Motion of electrons around an atom or molecule can result in a momentary charge separation. This instantaneous dipole can induce a similar dipole in a neighboring atom resulting in an electrostatic attraction. This attraction is weak; however, it can be significant and is more significant for larger molecules and atoms.
Larger atoms and molecules have more electrons and a larger electron "cloud". This results in an increased ability to form instantaneous dipoles (increased polarizability).
The strength of intermolecular forces can be seen in physical properties such as boiling points and melting points. Stronger intermolecular forces result in higher melting points and higher boiling points.
Molecules with a dipole moment have a charge separation, an area of positive charge and an area of negative charge. These charges interact with the charges on adjacent molecules. The molecules tend to orientate themselves to minimize electrostatic repulsion and maximize electrostatic attraction. This results in a net electrostatic attraction between molecules with dipole moments. This is called Dipole-Dipole Attraction.
A special case of Dipole-Dipole Attraction is called hydrogen bonding. This occurs when hydrogen is bound to a highly electronegative atom (N, O, F). This results in a particularly strong attractive force because of the large polarity combined with the small size of the hydrogen, which allows the molecules to get closer.
Ion-dipole forces are the attractive electrostatic forces that form between an ion and polar molecules. The polar molecules line up the partial charge of opposite polarity with the ion. This creates the attractive force. This force is important for the dissolution of ionic compounds in polar solvents such as water.
In order of increasing strength.
Dispersion forces < dipole-dipole forces < hydrogen bonding < ion-dipole forces
Viscosity is a measure of the liquid's resistance to flow.
Strong intermolecular forces increase a liquids viscosity.
Molecular complexity increases viscosity since complex molecules can become entangled with each other.
Surface Tension is the resistance of a liquid to increase its surface area. A molecule on the surface of a liquid only has a fraction of the intermolecular attractions that a molecule in the interior of the liquid has. Increasing the surface area requires breaking intermolecular forces.
Stronger intermolecular forces results in stronger surface tension.
Surface tension allows insects to float in the surface of water, and makes liquids form spherical drops.
Capillary Action is the spontaneous rising of a liquid in a narrow tube.
Capillary Action is the result if the interplay of cohesive forces (within a material) and adhesive forces (the forces between materials).
The curved surface of a liquid in a tube is called the meniscus.
A concave meniscus (high on the sides and low in the center) indicates that the adhesive forces are larger than the cohesive forces
A convex meniscus (high in the center and low on the sides) indicates that the cohesive forces are larger than the adhesive forces.
With strong cohesive forces and stronger adhesive forces, the liquid is pulled up the tube.
Heat of Fusion (Enthalpy of Fusion)(DHfus) is the heat energy required to melt a solid (or the heat released when a liquid freezes).
Q = n DHfus
When a solid and liquid remain in equilibrium, the temperature remains steady at the melting point.
More precisely the normal melting point is the temperature at which the solid and liquid states have the same vapor pressure under conditions where the total pressure is 1 atmosphere.
The process of liquid becoming gas. This process is endothermic.
Heat of Vaporization (Enthalpy of Vaporization)(DHvap) is the energy required to vaporize a mole of a liquid (or the energy released when the liquid vapor condenses).
at a pressure of 1 atm.
Q = n DHvap
Water has a large heat of vaporization at 40.7 kJ/mol
The evaporation of water is commonly used for cooling (swamp coolers for houses, sweating for humans, and cooling towers for power plants)
Some solids also have a measurable vapor pressure.
The process of going from a solid directly to gas is sublimation.
The heat of sublimation (DHsub) is equal to the sum of DHfus + DHvap at the same conditions (temperature and pressure).
When we heat a solid, the temperature of the solid increases. When the solid starts to melt, the temperature stabilizes at the melting point. After all the solid melts, the temperature of the liquid increases as we continue to add heat. When the liquid starts to boil the temperature stabilizes at the boiling point. After all the liquid has vaporized, then the gas temperature increases as we continue to add heat.
The Critical Temperature is the temperature above which the vapor cannot be liquefied (through a phase transition) regardless of the pressure applied
The Critical Temperature is the temperature above which there is no clear transition between vapor and liquid
The critical pressure is the pressure required to produce liquefaction at the critical temperature.
The critical point is the set of critical temperature and critical pressure.
Temperatures or pressures above the critical point are referred to as supercritical.
Vapor Pressure is the pressure of a gas in equilibrium with its liquid state (or solid state).
This equilibrium is when the rate of evaporation is equal to the rate of condensation.
Liquids with high vapor pressures are called volatile, and they evaporate rapidly
Strong intermolecular forces result in a low vapor pressure
Vapor pressure is temperature dependent and increases rapidly with increasing temperature
Boiling occurs when the vapor pressure of the liquid is at or above the pressure above the liquid
The normal boiling point is the temperature at which the vapor pressure of the liquid is exactly 1 atmosphere
A phase diagram represents phases as a function of temperature and pressure
Phase diagrams represent closed systems where no material can escape and no other material is present
As the expression goes, "A picture is worth a thousand words", a phase diagram reveals much information about a material.
The triple point is the one set of temperature and pressure where three states of matter coexist at equilibrium
The solid-liquid line in the phase diagram for water has a negative slope. Most other substances have a positive slope for this line. For water, this means that the liquid is more dense than the solid. It also means that increased pressures result in a lower melting point.
Ice sublimes at temperatures below 0C. The phase diagram indicates that ice should not sublime at normal atmospheric pressure. However, we shouldn't look at the atmospheric pressure but at the pressure of the water vapor. As long as the water vapor is below 4.58 torr, then ice can sublime.
The solid-liquid line for carbon dioxide is positive, this is typical. It indicates that the solid is more dense than the liquid
The triple point occurs at 5.1 atm and –56.6°C
This means that at 1 atm carbon dioxide will transition from solid directly to gas. Dry ice is solid carbon dioxide, which sublimates instead of melting.
The critical point is 72.8 atm and 31°C. Carbon dioxide above the critical point is used to extract various materials referred to as supercritical fluid extraction, and is used as an environmentally friendly dry cleaning agent.
A crystalline solid consists of a well-ordered structure. Often crystalline structures exhibit flat surfaces (faces) and definite angles between the faces.
An amorphous solid lacks an orderly structure. Amorphous solids can arise from mixtures, large complicated molecules, or from quick freezing.
For example; glass, primarily silicone dioxide, can be amorphous as is most glass products, or if cooled real slow, can be made into "crystal" (often with other additives to help the crystal formation).
Atomic Solids
Molecular solids are the solids whose composite units are atoms or molecules held together by intermolecular forces. The intermolecular forces in molecular solids are; hydrogen bonding, dipole-dipole forces, and dispersion forces.
These are relatively weak forces, which results in a soft solid with low to moderate melting points, and poor thermal and electrical conductivity.
Comparison of atomic Separation within Molecules and Between Molecules
A lattice network of atoms held together by covalent bonds. This is a very strong arrangement.
These tend to be hard materials with very high melting points, and poor thermal and electrical conductivity.
Example: diamond, quartz.
Ionic solids are the solids whose composite units are formula units of ions.
Ionic solids are held together by the attractive forces between positive ions and negative ions.
Ionic solids tend to be hard, brittle, have high melting points, and poor thermal and electrical conductivity
Determining the Number of Ions in a Unit Cell
The attractive force holding metallic atoms together in a lattice network can be described as positive metal ions in a sea of valence electrons.
The valence electrons are mobile throughout the solid.
The strength of the bonding increases as the number of valence electrons available for bonding increases.
Metallic solids tend to be soft to hard, low to high melting point, and excellent thermal and electrical conductivity
A regular 3-dimensional array of atoms (lattice points) of a crystalline solid.
The unit cell is the smallest unit of the array that shows the character of the crystalline solid and stacks together to form the solid. The unit cell usually includes fractions of atoms at the corners, edges, and faces.
There are three common types of unit cells; primitive cubic, body-centered cubic, and face centered cubic.
The primitive cubic has identical lattice points (atoms) in each of the eight corners.
The body-centered cubic has identical lattice points (atoms) in each of the eight corners plus another (different) lattice point in the center.
The face-centered cubic has identical lattice points (atoms) in each of the eight corners plus a set of a different set of six identical lattice points centered on each face of the cube.
The closest arrangement of spherical atoms in a three-dimensional structure (such as in metals) results in a 2-dimensional layer with each atom surrounded by 6-atoms.
There are two ways to stack these 2-dimensional layers to make the three dimensional structure. It every other layer lines up, this is the hexagonal close packing. If every third layer lines up, this is the cubic close packing.
In both cases, 12 equally close atoms surround each atom. Each atom has a coordination number of 12.
Hexagonal Closest Packed (hcp) Structure
Cubic Closest Packed (ccp) Structure
The wavelength of x-rays is approximately the same as the spacing of layers in crystals (about 2 – 20 °A), this causes the regular scattering or diffraction of the x-rays. Interference of the scattered x-rays results in a pattern of light and dark spots of the collector. This information can be analyzed to reveal the structure of the crystal.
X-Ray Diffraction
Using the Bragg Equation
Types of Crystalline Solids
Structure and Bonding in Metals
Calculating the Density of a Closest Packed Solid
Bonding Models for Metals
Band Model
Molecular Orbital (MO) Model
Metal Alloys
Alloy
Substitutional alloy
Interstitial Alloy
Carbon and Silicon: Network Atomic Solids
Silica
Silicates
Glass
Ceramics
Semiconductors
Semiconducting element (semiconductor)
n-type Semiconductor
p-type Semiconductor
p-n junction
Properties of Solids
Not Chem. 1A
Determining the Enthalpies of Vaporization
A relationship exists between vapor pressure, temperature, and the enthalpy of vaporization
ln(Pvap) = (-DHvap/R)(1/T) + C
Where R is the gas constant and C is a constant unique to each liquid
This equation has the form of a straight line (y = mx + b) when plotted with y = ln(Pvap), x = (1/T), m = slope = (-DHvap/R), b = intercept = C
This equation can be used to calculate the enthalpy of vaporization by measuring vapor pressure at several temperatures.
Calculating Vapor Pressure
If enthalpy of vaporization is known along with vapor pressure at one temperature, then vapor pressure at another temperature can be calculated.
The equation can be rearranged to give
ln(Pvap,T1/Pvap,T2) = (DHvap/R)(1/T2 – 1/T1)
Supercoooled
Supercooling can occur where a liquid is cooled below the freezing point and still remains liquid. This occurs when the alignment associated with the solid has not occurred. When this alignment is achieved, freezing can occur rapidly.
Superheated
Superheating can occur when a liquid is heated, usually rapidly, to a temperature above the boiling point without boiling occurring. When a bubble forms, it expands rapidly, ejecting a volume of liquid with it. This is referred to as "bumping" and has been reportedly to occasionally to occur in microwave ovens.
Q = n DHfus
Q = n DHvap