Chapter 12 – Liquids and Solids

Which one represents a liquid? Why?

Liquids have a definite volume, but not a definite shape. The particles are closer together than gases so the intermolecular forces are now a factor.

Hydrogen Bonding Dipole-Dipole London Forces

Due to the greater attraction of the intermolecular forces, the particles are more orderly than gases, and have a lower mobility.

Although there liquids have a lower mobility, they are still able to flow, thus classified as a fluid. They are, therefore, able to diffuse.

Due to the greater attraction of the intermolecular forces, the particles are more orderly than gases, and have a lower mobility. Although there liquids have a lower mobility, they are still able to flow, thus classified as a fluid. They are, therefore, able to diffuse. The more orderly arrangements causes liquids to be about 1000 times more dense than the gas form of the substance. The typical density of liquids makes them incompressible.

In addition, liquids will exhibit surface tension. The stronger the intermolecular forces, the more they can pull the surface molecules together.

The greater the strength of the intermolecular forces, the smaller the resulting surface area will be. (A sphere has the smallest surface area for it’s volume of any other geometric shape.)

Often liquids not only experience intermolecular forces within themselves, but can experience intermolecular forces with the walls of a container. The attraction of a liquid to the walls of a container results in capillary action – which forms the meniscus, and in a capillary tube can even pull the liquid up the tube.

Which one represents a Solid? Why?

Solids particles are more closely packed than even liquids (usually the most dense phase). Due to very large intermolecular forces, solids have definite volume and definite shape.

The only motion the particles are able to have is vibrational and rotational, there is no more ability to have translational motion (as with liquids).

Solids particles are more closely packed than even liquids (usually the most dense phase). Due to very large intermolecular forces, solids have definite volume and definite shape. The only motion the particles are able to have is vibrational and rotational, there is no more ability to have translational motion (as with liquids). Solids are also incompressible and have a low rate of diffusion.

There are two types of solids: Amorphous solids have particles that are randomly arranged.

There are two types of solids: Amorphous solids have particles that are randomly arranged. Amorphous solids do not have a definite melting point, but rather become more and more fluidic when more and more heat is added. Thus they are sometimes called supercooled liquids. Glass is an amorphous solid, and old windows are thicker at the bottom and thinner at the top, adding to the draftiness of old buildings.

Crystalline solids have particles arranged in an orderly, geometric, repeating pattern (a crystal). Crystalline solids will have a definite melting point, which can often be useful in identifying a particular solid.

Crystal systemLattices: triclinic monoclinic simplebase-centered orthorhombic simplebase-centeredbody-centeredface-centered

hexagonal rhombohedral (trigonal) tetragonal simplebody-centered cubic (isometric) simplebody-centeredface-centered

The total 3-dimensional arrangement of the particles in a crystal is referred to as the crystal structure. The entire crystal structure in a sample is the crystal lattice. The smallest portion of the crystal lattice that shows the 3-dimensional pattern (as in the previous graphics) is called the unit cell.

Solids can be categorized by shape based on the unit cell (as in the previous graphics) or on the binding force. When based on the binding force there are four types of crystals: Molecular, Ionic, Covalent Network, and Metallic.

Form of unit particles Atoms Nonpolar Molecules Polar Molecules Forces between particles London H-bonding Dipole-dipole Properties Soft, very low melting point, poor conductors Soft, low to medium melting points, poor conductors ExamplesAr, Kr CH 4, Sugar, Dry Ice Ice, NI 3 Molecular Solids

Form of unit particles Positive and Negative Ions Atoms connected in a covalently bonded network Atoms Forces between particles Electrostatic Attraction Giant molecules of covalent bonds Metallic Bond (sea of electrons) Properties Hard and Brittle, high melting point, poor conductors Very Hard, very high melting point, poor conductors Soft to very hard, low to very high melting points, Very good conductors, malleable, ductile Examples NaCl, Ca(NO 3 ) 2 Diamond (C), Quartz (SiO 2 ) Iron, Lead Ionic Covalent NetworkMetallic

Phases are any part of a system that has uniform composition and properties. Phase changes occur as a substance changes from one phase to another. Melting / Freezing – solid to liquid / liquid to solid Vaporizing (Boiling) / Condensing – liquid to gas / gas to liquid Sublimation / Deposition – solid to gas / gas to solid Evaporation – vaporization occurring below boiling point (temperature) – only molecules with high energy at the surface become a vapor (a cooling process)

Equilibrium is a dynamic condition in which two opposing changes occur at equal rates in a closed system. (Like ice in water at 0°C.) For chemical reactions, this is shown with a double sided arrow.

LeChatelier’s Principle – when a system that was at equilibrium is subjected to some outside stress, the system reacts in such a way to relieve the stress and return to another equilibrium.

If the concentration of a chemical is increased / decreased, shift to use up more / make more of that chemical.

If gases are involved, increasing / decreasing the pressure (or decreasing / increasing the volume) will shift the equilibrium to the side of the equation with the fewer / greater number of gaseous chemicals.

Le Chatelier’s Principle Review 1 Le Chatelier’s Principle Review 2

Equilibrium vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid at a given temperature. Volatile liquids evaporate easily due to weak intermolecular forces.

Boiling is the conversion of liquid to vapor, within the liquid as well as at its surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure. The temperature at which the liquid must be heated to reach an equilibrium with the atmospheric pressure is called the boiling point. Pressure cookers increase the vapor pressure to increase the temperature at which the water must be heated before it will boil – thus cooking food faster.

Molar heat of vaporization is the amount of heat energy needed to vaporize one mole of liquid at the boiling point. (The stronger the intermolecular forces, the higher the heat of vaporization.)

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

Molar heat of vaporization is the amount of heat energy needed to vaporize one mole of liquid at the boiling point. (The stronger the intermolecular forces, the higher the heat of vaporization.) Normal freezing point is the temperature at which the solid and liquid are in equilibrium at 1 atm. Molar heat of fusion is the amount of heat energy required to melt one mole of solid at its melting point.

Phase diagrams:

The triple point is the temperature at which all three states of matter are in equilibrium (solid, liquid and gas)

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

The critical pressure is the pressure above which the substance cannot exist in the liquid state if above the critical temperature.

You must be able to label on a blank chart: sublimation curve melting curve boiling curve normal melting and boiling points (temperatures) triple point critical temperature solid, liquid, and gas regions

Water is a very special compound! It is a polar covalent compound with angular molecular geometry. Hydrogen bonding between water molecules accounts for water’s high melting point (for it’s mass), high heats of fusion and vaporization (more energy to pull the molecules apart), and high melting and boiling points. In the solid phase, water has a hexagonal arrangement with more empty space between the molecules than as a liquid. This gives ice a lower density than the liquid.

Compare water’s phase diagram to carbon dioxide.

It should be obvious by now that heat has a large influence over the phases of matter. If it is desired to quantitize the amount of effect the heat can cause the following formula will be useful: Q = m × C p × ∆T where Q = heat (in calories or joules) m = mass (in grams) C p = specific heat ∆T = T final – T initial (in °C)

The specific heat is different with every substance, and is also different for each phase of a substance. For example: C p ice = 2.06 J/g °C C p water = 4.18 J/g °C C p steam = 2.08 J/g °C C p Al = J/g °C

During a phase change, the temperature is constant, therefore temperature cannot be a factor in the equation for heat during a phase change, thus: Q = m × H f or Q = m × H v where H f = molar heat of fusion H v = molar heat of vaporization

A great way to visualize what is occurring during a process that involves changing the heat of a substance is with a heating curve:

Each segment will require a different formula due to different values for C p, H f, or H v ! To solve the heat necessary to do the curve below, it would take 5 calculations and a sum!