Chapter 11: Intermolecular Forces and Liquids and Solids • Phases include solids (molecules held rigidly in position), liquids (molecules packed close together with definite volume) and gases (molecules readily compressed due to lack of strong forces). • Intermolecular forces are the attractive forces between molecules (as opposed to intramolecular forces, which hold the atoms in a molecule together). Intermolecular forces are weaker than intramolecular forces, meaning they require much less energy to break. As the strength of intermolecular forces increases, the boiling point of the substance also increases. o Dipole-dipole forces are the attractive forces between polar molecules, or those that have dipole moments. o Ion-dipole forces are the attractive forces between an ion (cation or anion) and polar molecule. How strong they are depends on the charge and size of the ion, as well as the magnitude of the dipole moment and size of the molecule. Cation-dipole interactions are stronger than anion-dipole interactions because cations are usually smaller (and therefore their charge is more concentrated). o Dispersion forces occur between polarizable electron clouds, and are defined as attractive forces that result from temporarily induced dipoles in atoms or molecules. In nonpolar molecules, the only intermolecular forces present are dispersion forces. § Induced dipoles are caused by the proximity of an ion or polar molecule, which separates the positive and negative charges in an atom or molecule. o Hydrogen bonding is a dipole-dipole interaction between hydrogen and an electronegative F, O or N atom. The strength is determined by electronegativity (N is weakest, O is stronger, F is strongest). o EXAMPLES: CCl4 is nonpolar because the identical Cl charges around C “balance out,” therefore it only experiences dispersion forces. HBr is polar because Br more strongly attracts the H atom; it experiences dispersion and dipole-dipole forces. • Intermolecular forces gives rise to surface tension, or the amount of energy necessary to stretch the surface of a liquid by a unit area (i.e. 1 cm2). Strong intermolecular forces indicate high surface tension. Therefore, water has considerably high surface tension due to hydrogen bonding. o Capillary action, when water rises spontaneously in a capillary tube, is an example of surface tension. Two forces cause this: cohesion (intermolecular attractions between like molecules) and adhesion (attractions between unlike molecules). • Viscosity measures a fluid’s resistance to flow; the greater the viscosity, the slower it flows. Liquids with strong intermolecular forces have stronger viscosities (therefore, water also has a particularly high viscosity because of strong hydrogen bonding).• The difference between water and other polar molecules is that water has the capability of forming two hydrogen ponds, giving water an extensive three-dimensional network. Solid ice is less dense than liquid water, because as it is heated, water molecules are released from intermolecular (three-dimensional) hydrogen bonding. • Crystalline solids have rigid order, meaning its atoms, molecules, or ions occupy specific positions. They are arranged so the next attractive intermolecular forces are at their maximum. The basic, repeating structural units of crystalline solids are unit cells. o Each sphere representing an atom, ion or molecule is a lattice point. o Packing spheres show the arrangement of crystalline solids: o There is 1 atom in a simple cubic cell, 2 atoms in body-centered cubic cell, and 4 atoms in a face-centered cubic cell. • Bragg Equation for calculating distance between planes of atoms in a lattice: • Atoms of covalent crystals are held together in three-dimensional networks entirely by covalent bonds; examples are diamond, graphite and quartz. These are hard, have high melting points, and are generally poor conductors. • Lattice points of molecular crystals are occupied by molecules, and held together by van der Waals forces and/or hydrogen bonding. Examples are ice, I2, P4, and S8. These are more easily broken apart than ionic or covalent crystals, and therefore have lower melting points. • Each lattice point in a metallic crystal is occupied by an atom of the same metal. These are generally very dense, have low to high melting points, and are good conductors (include all metallic elements).• Amorphous solids lack a well-defined 3D arrangement of atoms (such as glass). Glass is an optically transparent fusion product of inorganic substances, which have cooled to a rigid state without crystallizing. • Phase changes are transitions of matter from one phase to another, and require an energy input. o Evaporation (or vaporization) is the process by which a liquid transforms into a gas; the higher the temperature, the more molecules that evaporate. § When liquid evaporates, the molecules in gaseous form exert a vapor pressure above the liquid. § As the concentration of molecules in gas phase increase, some gas molecules return to liquid state. This is called condensation. § The rate of evaporation remains constant; the rate of condensation increases as the number of vapor molecules increases. • Molar heat of vaporization (∆Hvap) measures the strength of intermolecular forces in a liquid, and is defined as the energy needed to vaporize 1 mole of liquid. It is directly proportional to the strength of intermolecular forces. o Clausius-Clapeyron Equation: o To calculate boiling point: R is the gas constant 8.314 J/K•mol and C is a constant • Boiling point is the temperature at which vapor pressure equals external pressure. o The higher the heat of vaporization, the higher the boiling point • Critical Temperature is the temperature at which a substance’s gas phase cannot be made to liquify; Critical Pressure is the minimum pressure required to liquefy a substance at critical temperature • Freezing is the transformation of a liquid to solid. The melting or freezing points of a liquid are the temperatures at which the solid and liquid phases coexist at equilibrium. • Molar heat of fusion (∆Hfus) is the energy (in kJ) required to melt 1 mole of a solid. This quantity is smaller than heat of vaporization, because the molecules of a liquid are still closely packed together. • Supercooling occurs when a substance is cooled below its freezing point.•
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