Intermolecular Force: Liquids, Solids and Phase Changes

Chemistry 212 Outline:

Chap 12- Intermolecular Force: Liquids, Solids and Phase Changes

Phase changes: the changes from one phase to another

• (liquids and solids are called condensed phases (or condensed states) because their particles are extremely close)

12.1 Two types of electrostatic forces:

1)Intramolecular Forces (bonding forces): exist within each molecule and influence the chemical properties of the substance.

2) Intermolecular Forces ( nonbonding forces): exist between the molecules and influence the physical properties of the substance.

Types of Phase Change

• Condensing and freezing are exothermic (loosing heat) changes.

• Melting and vaporizing are endothermic ( gaining heat) changes

Section Summary:

• Because of the relative magnitudes of intermolecular forces and kinetic energy, the particles in a gas are far apart and moving randomly, this in a liquid are in contact but still moving relative to each other and those in a solid are in contact and fixed relative to one another in a rigid structure.

• These molecular-level differences in the states of matter account for macroscopic differences in shape, compressibility and ability to flow.

• When a solid becomes a liquid ( melting or fusion) or a liquid becomes a gas (vaporization), energy is absorbed to overcome intermolecular forces and increase the average speed between particles. When particles come closer together in the reverse changes ( freezing and condensation), energy is released.

• Sublimation is the changing of a solid directly into a gas. Each phase is associated with a given enthalpy change under specified conditions.

12.2 Quantitative Aspects of Phase Changes

Heating-cooling curve: shows the changes that occur when heat is added or removed from a particular sample of matter at a constant rate.

• Hess' s law: the total heat released is the sum of the heats released for the individual stages.

• Two key points stand out in this or any similar process (at constant pressure), whether exothermic or endothermic:

• Within a phase, a change in heat is accompanied by a change in temperature, which is associated with a change in average Ek as the most probable speed of the molecules changes. The heat lost or gained depends on the amount of substance, the molar heat capacity for that phase, and the change in temperature.

• During a phase change, a change in heat occurs at a constant temperature, which is associated with a change in Ep, as the average distance between molecules changes. Both physical states are present during a phase change. The heat lost or gained depends on the amount of the substance and the enthalpy of the phase change ( ΔHvap or ΔHfus ).

The Equilibrium Nature of Phase Changes

• In everyday experience, phase changes take place in open containers -the out doors

• In a closed container under controlled conditions, phase changes of many substances are reversible and reach equilibrium, just as chemical changes do.

• Vapor pressure: the pressure exerted by the vapor at equilibrium ( also called equilibrium vapor pressure)

• Behavior of a pure liquid in contact with its vapor is a general one for any system: when a system at equilibrium is disturbed, it counteracts the disturbance and eventually re-establishes a state of equilibrium.

The Effects of Temperature and Intermolecular Forces on Vapor Pressure

• The vapor pressure of a substance depends on temperature

• In general, the higher the temperature is, the higher the vapor pressure.

• The vapor pressure also depends on the intermolecular forces present.

• In general, the weaker the intermolecular forces are, the higher the vapor pressure.

The nonlinear relationship between vapor pressure and temperature shown in Fig. 12.6 can be expressed as a linear relationship between ln(P) and (1/T) using the Clausius-Clapeyron equation.

• Clausius-Clapeyron equation: gives a us a way of finding the heat of vaporization( the energy needed to vaporize 1 mol of molecules in the liquid state)

Vapor Pressure and Boiling Point

• Boiling point: temperature at which the vapor pressure equals the external pressure, which is usually that of the atmosphere

• Boiling point depends on applied pressure

• The normal boiling point is observed at standard atmospheric pressure (760 torr, or kPa)

Solid-Liquid equilibria:

• Because the phase remain in contact, a dynamic equilibrium is established when the melting rate equals the freezing rate. The temperature at which this occurs is called the melting point; it is the same temperature as the freezing point, differing only in the direction of energy flow.

• As we saw with the boiling point, the temperature remains fixed at the melting point as long as both phases are present.

Solid-Gas equilibria:

• Solids have much lower vapor pressures than liquids.

• Sublimation, the process of a solid changing directly into a gas, is much less familiar than vaporization because the necessary conditions of pressure and temperature are uncommon for most substances.

• Some solids do have enough high vapor pressure to sublime at ordinary conditions include dry ice (carbon dioxide), iodine and solid room disorders

• A substance sublimes rather than melts because the combination of intermolecular attractions and atmospheric pressure is NOT great enough to keep the particles near one another when they leave the solid state.

Phase Diagrams: Effect of Pressure and Temperature on Physical State

• Four features of Phase Diagrams:

1)Regions of the Diagram:

• Each region corresponds to one phase of the substance

• A particular phase is stable for any combination of pressure and temperature with its region

• If any of the other phases is placed under those conditions, it will change to the stable phase

• In general,the solid is stable at low temperature and high pressure,

• gas is stable at high temperature and low pressure.

• Liquid is stable at intermediate conditions

2)Lines between Regions:

• The lines separating the regions represent the phase-transition curves discussed earlier

• Any point along a line shows the pressure and temperature at which the two phases exist in equilibrium

• Note that the solid -line has a positive slope (slants to the right with increasing pressure) because, for most substances, the solid is more dense than the liquid.

• Because the liquid occupies slightly more space than the solid, an increase in pressure favors the solid phase.

• Water is the major exception

3)The critical point:

• The liquid-gas line ends at the critical point

• Example of a liquid in a closed container. As it is heated, it expands, so its density decreases. At the same time, more liquid vaporizes, so the density of the vapor increases.

• The liquid and vapor densities become closer and closer to each other until, at the critical temperature (Tc), the two densities are equal and phase boundary disappears.