Intermolecular Forces, Liquids and Solids

AP Chapter 11

Molecular Comparison

Intermolecular Forces

• In liquids, the intermolecular forces are strong enough to keep the molecules in close proximity, but the molecules have enough kinetic energy to move past each other.

• In solids, the forces are much stronger; molecular motion is restrained and particles occupy specific locations in a 3 dimensional geometric pattern.

Types of Intermolecular Forces

• Dipole-dipole forces – the forces that act between polar molecules

• London dispersion forces – (Van der Waals) the forces that act between non-polar molecules, as well as all molecules

• Hydrogen bonding – occurs between hydrogen and another element of high electronegativity. (O-H, N-H, F-H)

Dipole-dipole attraction

Ion-Dipole Forces

• Ion-dipole forces occur between solutions in which ionic compounds are dissolved in polar solvents.

• The strength of the bond increases with the difference in electronegativity within the molecule.

• The greater the electronegativity inside the molecule, the stronger the dipole forces between molecules.

Dipole-dipole Forces

Hydrogen Bonding in Ice

London Dispersion Forces

• London Dispersion Forces increase in strength with an increase in molecular mass.

• This is the reason why the halogens go from gas to solid as the atomic mass increases.

Comparing Intermolecular Forces

Viscosity

• The stronger the intermolecular forces, the greater the viscosity, or resistance to flow, of a liquid.

• Surface tension of liquids also increase as intermolecular forces increase in strength.

• Surface tension is defined as a measure of the tendency of a liquid to maintain a high surface area.

Surface Tension

Capillary Action

• The adhesion of a liquid to the walls of a narrow tube and the cohesion of a liquid account for the capillary action and the formation of a meniscus at the surface of a liquid.

Phase Changes

• Substances exist in more than one state of matter.

• Phase changes are transformations from one phase to another.

• Solid to liquid is melting (fusion), solid to gas is sublimation, liquid to gas is vaporization.

• Heats of fusion, vaporization and sublimation are endothermic processes.

Phase changes, continued.

• The reverse changes (freezing and sublimation) are exothermic processes (negative values.)

Heating and Cooling Curves

Formulas to Use in Calculations

• q = mCΔT

• q = mHfus

• q = mHvap

Critical Temperature and Pressure

• Critical temperature is the highest temperature at which a liquid can exist.

• Above the critical temperature, the motional energies of the molecule are greater than the attractive forces that lead to the liquid state regardless of how much pressure is applied.

• The critical pressure is the pressure required to create a liquid at this critical temperature.

Vapor Pressure

• Vapor pressure of a liquid indicates the tendency of a liquid to evaporate.

• Molecules can escape the surface of a liquid and into the gas phase by evaporation.

• The pressure exerted back onto the surface of the liquid by the vapor is called the vapor pressure.

Dynamic Equilibrium

• The condition in which two opposing processes are occurring simultaneously at equal rates is called dynamic equilibrium.

• When some molecules are escaping the surface of a liquid and some molecules are condensing at an equal rate, the system is in dynamic equilibrium.

• This process happens only in a closed system!

Volatility

• Substances with very high vapor pressure, such as gasoline, evaporate more quickly than substances with low vapor pressure, such as motor oil.

• Liquids that evaporate readily are said to be volatile.

Boiling Point• A liquid boils when its vapor pressure equals

the external pressure acting on the surface of the liquid.

• This is the point where bubbles of vapor form within the liquid.

• The normal boiling point of a liquid at 1 atm pressure is the normal boiling point.

• The temp. at which a liquid boils increases with increasing external pressure.

Phase Diagrams

• A phase diagram is a graphic way to summarize the conditions under which equilibria exist between the different states of matter.

• These diagrams allow for the prediction of the phase of a substance that is stable at any given temperature and pressure.

Phase Diagram definitions• The line T to C is the vapor pressure curve of

the liquid. It represents the equilibrium between the liquid and the gas phases.

• The vapor pressure curve ends at C, the critical point, which is the critical temperature and critical pressure for that substance.

• Beyond that point, the liquid and gas phases become indistinguishable and the state is a supercritical fluid.

• The line that separates the solid phase from the gas phase represents the change in the vapor pressure of the solid as it sublimes at different temperatures.

• An increase in pressure usually favors the more compact solid phase, so higher temperatures are required to melt the solid at higher pressures.

• Point T, where the three curves intersect, is known as the triple point. All three phases are in equilibrium at this temperature and pressure.

Structures of Solids

• In a crystalline solid, the atoms, ions or molecules are ordered in well-defined, 3-dimensional arrangements.

• The solids have flat surfaces or faces that make definite angles with one another.

• An amorphous solid is a solid in which particles have no orderly structure.

Unit Cells

• The repeating unit of a solid is called a unit cell.

• While there are several ways to choose the unit cell, it is generally the smallest unit cell that shows the symmetry of the entire pattern.

A 2-dimensional analog of a lattice and its unit cell. This shows the repeating pattern of the lattice.

Crystalline Solids

• A crystalline solid can be represented by a 3-D array of points called a crystal lattice.

• Each point in the lattice is called a lattice point, and represents an identical environment within the solid.

• It acts like a scaffolding for the whole crystal.

Lattice points are at the corners only.

There is a lattice point at the center of the unit cell.

This unit cell has lattice points at the center of each face, as well as each corner.

This is a space-filling view of cubic unit cells. Only the portion of each atom that belongs to the unit cell is shown.

In the crystal structure of NaCl, we can center either the Na+ (purple) or the Cl- ions (green) on the lattice points. In both cases, they are face-centered cubic.

Close-Packing of Spheres

• Many solids have close-packed structures in which spherical particles are arranged in order to leave the minimum amount of space.

• There are two forms of close-packing; cubic close-packing and hexagonal close-packing.

Close-Packing of Spheres

• In both cases, each sphere has a coordination number of 12, which means the each sphere has 12 equidistant neighbors.

• There are 6 neighbors in one plane, 3 below and 3 above.

Hexagonal close-

packing

Cubic close-packing

Bonding in Solids

• The physical properties of solids, such as melting point and hardness depend on both the arrangement of the particles and on the attractive forces between them.

Types of Crystalline Solids

Covalent Network Solids• Covalent network solids consist of atoms

held together by covalent bonds in large networks or chains.

• Because covalent bonds are much stronger than intermolecular forces, these solids are much harder and have higher melting points than regular molecular solids.

• Diamond and graphite are 2 allotropes that are network solids.

Ionic Solids

• Ionic solids consist of ions held together by ionic bonds.

• The strength of the ionic bonds depends on the charges of the ions.

• NaCl, with charges of +1 and -1, has a melting point of 801°C. MgO, with charges of +2 and -2, has a melting point of 2852°C.

Unit cells of some common ionic solids.

Metallic Solids

• Metallic solids consist of metal atoms.

• They usually have hexagonal close-packed, cubic close-packed (face-centered) or body-centered cubic structures.

• Each atom is typically surrounded by 8 or 12 adjacent atoms.

• Held together with delocalized valence electrons – (sea of mobile electrons)

Sea of mobile

electrons in

metallic bonds