Liquids and Solids

Intermolecular Forces

Intermolecular Forces are relatively weak interactions that occur between molecules.  There are three main types of forces:  

Dipole-Dipole Forces

Molecules that have an asymmetric distribution of charge within the molecule are referred to as being polar. Since atoms differ in their ability to attract electrons to themselves (electronegativity), any situation in which two different atoms are bound together will result in an asymmetric distribution of charge. Almost every simple diatomic molecule consisting of two different kinds of atoms will be polar. The end of the molecule with the atom having the greater ability to attract electrons will be slightly negative and the other end will be slightly positive. The Dipole-dipole force then arises from the attraction between the slightly positive end of one molecule and the slightly negative end of another. 

As a rule, dipole-dipole attractions are stronger than London forces although it's hard to make simple comparisons because polar molecules will exhibit both dipole-dipole and London forces. 

Dipole-Dipole Forces animated.

Hydrogen Bonding

Illustrated above is an example of hydrogen bonding.

To fully understand hydrogen bonding, it's necessary to go back to the original idea of Coulombic forces. These are the attractive (or repulsive) forces between charged particles. The force of attraction between two opposite charges is proportional to the magnitude of their charges divided by the square of the distance between them. If we hook hydrogen up to an atom that is very good at attracting electrons (like N, O, or F) the hydrogen end of the bond becomes very positively charged and the other atom becomes negatively charged (i.e., polar). Now if we remember that hydrogen is the smallest atom on the periodic table, we realize that it is possible for two molecules to get very close together. The combination of high polarity and close approach result in the interaction being particularly strong. In fact, the interaction is so strong that it dwarfs all other dipole-dipole attractions and is given the special name of "hydrogen bonding".

London Dispersion Forces

London Dispersion Forces animated.

The simplest of the intermolecular forces, termed London dispersion forces after their discoverer Fritz London, arise (simplistically) from the weak attractive force of the electrons on one molecule for the nuclei of another molecule. There is no bond, i.e., no sharing of electrons between the two molecules.

Because London forces arise from interactions present between all molecules, these forces are always present. For non-polar molecules these are the only forces present. As a general rule for simple molecules, if a covalent bond is set at a relative value of 100, London forces have values between 0.001 and 0.1. In otherwords, very weak forces.

The Liquid State

When considering the three states of matter, the middle state is known as the Liquid State. It has some of the characteristics of the solid state and some of the characteristics of the gas state.  Like the solid state, the liquid state tends to have very little ability to be compressed.  Like the solid state, the liquid state is associated with somewhat low kinetic energy.  Like the gas state, the liquid state tends to flow, or have fluidity.  Like the gas state, the liquid state tends to take on the shape of the container that it is in.  As the middle state, liquids are associated with the common phase changes freezing and boiling.  In addition, other concepts such as viscosity, evaporation, and surface tension are also associated with the liquid state.

Structures and Types of Solids

There are 5 main types of solids.

1. Ionic Solids

2. Covalent Solids

3. Polar Molecular Solids

4. Nonpolar Molecular Solids

5. Metallic Solids

To a large extent, the primary difference between the types of solids is the mechanisms that hold them together as solids. These mechanisms will be responsible for many physical characteristics, such as melting points, boiling points, hardness and water solubility.

Solid state crystal structures.

Carbon and silicon:  Network Solids

Carbon

Carbon exists as a pure element at room temperature in three different forms: graphite (the most stable form), diamond, and fullerene.

The most stable form of carbon is graphite. Graphite consists of sheets of carbon atoms covalently bonded together. These sheets are then stacked to form graphite.

Silicon Dioxide

Silicon dioxide (SiO₂), also called silica, occurs naturally in many forms. Quartz is essentially pure silicon dioxide. Sand is composed of small quartz fragments. Many precious gems are quartz containing colored impurities. Amethyst is quartz colored red by the presence of iron(III) ions. Agate and onyx are also quartz containing impurities. Flint is silica colored black by carbon.  Quartz has a very complicated crystal structure, which involves interwoven helical chains.  All silicates involve silicon in tetrahedral environments surrounded by oxygen atoms.

Ionic Solids

• High melting point and boiling point. Ionic bonds are strong.
• Soluble in polar solvents such as water. The ions have favorable interactions with the polar solvent and thus it takes less energy to remove them from the crystal
• Poor electrical conductivity. The electrons are localized around the ions, not spread out like in a metal.

Vapor Pressure and Changes of State

A picture summarization of vapor pressure.

During vaporization, molecules in a liquid gain enough kinetic energy to break the bonds holding them to the rest of the molecules in the liquid and enter the gas. This can also occur in reverse: gas molecules can collide with the surface of the liquid and re-enter the liquid.

If you place a liquid in a closed container, at first there are very few molecules of the liquid in vapor form, and so the reverse process doesn't occur very often. As more of the liquid converts to vapor, the rate of the reverse process increases until finally there are the same number of molecules entering and leaving the liquid. Once this equilibrium is reached, the number of molecules in vapor form doesn't change, and we can compute the pressure of the vapor. The equilibrium is temperature dependent: as the temperature rises so does the vapor pressure. For example, at 25^oC, water has a vapor pressure of 24 mmHg, at 100^oC it has a vapor pressure of 760 mmHg.

In an open container, the vapor molecules can spread out to a great distance, and the liquid never reaches the equilibrium that it does in a closed container. This is why puddles left by a rainstorm evaporate over time.

Once you know the vapor pressure of a liquid at a certain temperature, you can use the ideal gas law to relate V and n to the pressure.

Phase Diagrams

Equilibrium can exist not only between the liquid and vapor phase of a substance but also between the solid and liquid phases, and the solid and gas phases of a substance.

A phase diagram is a graphical way to depict the effects of pressure and temperature on the phase of a substance:

• The vapor pressure curve is the border between the liquid and gaseous states of the substance
  • For a given temperature, it tells us the vapor pressure of the substance
  • The vapor pressure curve ends at the critical point.

• The line between the gas and solid phase indicates the vapor pressure of the solid as it sublimes at different temperatures
• The line between the solid and liquid phases indicates the melting temperature of the solid as a function of pressure
  • For most substances the solid is denser than the liquid
  • An increase in pressure usually favors the more dense solid phase
  • Usually higher temperatures are required to melt the solid phase at higher pressures
• The "triple point" is the particular condition of temperature and pressure where all three physical states are in equilibrium
• Regions not on a line represent conditions of temperature and pressure where only one particular phase is present
  • Gases are most likely under conditions of high temperature
  • Solids are most likely under conditions of high pressure

Phase Diagram for Water

• The frozen state of water (ice) is actually less dense than the liquid state, thus, the liquid state is more compact than the solid state
  • Increasing pressure, which will favor compactness of the molecules, will thus favor the liquid state

• The melting curve slopes to the left, unlike most compounds
• At 100 °C the vapor pressure of water is 760 torr or 1 atm, thus at this temperature water will boil if it is at 1 atm of pressure
• At pressures below 4.58 torr, water will be present as either a gas or solid, there can be no liquid phase