Prepared by:Mrs Faraziehan Senusi

PA-A11-7CPhysical Transformation of Pure Substances

Chemical Equilibrium

Chapter 4Thermodynamic and Equilibria

First Law of Thermodynamics

Reference: Chemistry: the Molecular Nature of Matter and Change, 6th ed, 2011, Martin S. Silberberg, McGraw-Hill

Second Law of Thermodynamics

Simple Mixtures

Physical Transformation of Pure Substances

Viscosity

• Viscosity is the resistance of a liquid to flow.• A liquid flows by sliding molecules over each other.• The stronger the intermolecular forces, the higher the

viscosity.• As temperature increases and the molecules move more

rapidly, their kinetic energies are better able to overcome intermolecular attractions.

• Thus, viscosity decreases with increasing temperature, as long as no changes in composition occur.

Some Properties of Liquids

Surface Tension• Bulk molecules (those in the liquid) are equally attracted to

their neighbors.• Surface molecules are only attracted inwards towards the

bulk molecules.• Surface tension is a measure of the inward forces that must

be overcome to expand the surface area of a liquid.• Surface tension is the amount of energy required to increase

the surface area of a liquid.• Cohesive forces bind molecules to each other.• Adhesive forces bind molecules to a surface.

• Meniscus is the shape of the liquid surface. – If adhesive forces are greater than cohesive forces, the

liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e.g. water in glass).

– If cohesive forces are greater than adhesive forces, the meniscus is curved downwards. (e.g. mercury)

The meniscus, as observed in glass tubes with water and with mercury.

• Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid.

• These molecules move into the gas phase.• As the number of molecules in the gas phase increases, some of

the gas phase molecules strike the surface and return to the liquid.

• After some time the pressure of the gas will be constant at the vapor pressure.

• Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

Vapor Pressure

A representation of the measurement of vapor pressure of a liquid at a given temperature. • The container is evacuated before the liquid is added. • At the instant, the liquid is added to the container, there are no

molecules in the gas phase, so the pressure is zero. • Some of the liquid then vaporizes until equilibrium is established. • The difference in heights of the mercury column is a measure of the

vapor pressure of the liquid at that temperature.

Phase Changes

Hsub > 0 (endothermic)

Hvap > 0 (endothermic)

Hfus > 0 (endothermic)

Hdep < 0 (exothermic)

Hcon < 0 (exothermic)

Hfre < 0 (exothermic)

• Phase diagrams that show the equilibrium pressure–temperature relationships among the different phases of a given pure substance in a closed system.

• The curved line from A to C is a vapor pressure curve obtained experimentally by measuring the vapor pressures of water at various temperatures.

• Points along this curve represent the temperature–pressure combinations for which liquid and gas (vapor) coexist in equilibrium.

• At points above AC, the stable form of water is liquid; below the curve, it is vapor.

Phase Diagrams

• Line AB represents the liquid–solid equilibrium conditions.• We see that it has a negative slope. The negative slope (up and to the

left) indicates that increasing the pressure sufficiently on the surface of ice causes it to melt. This is because ice is less dense than liquid water in the vicinity of the liquid–solid equilibrium.

• The stable form of water at points to the left of AB is solid (ice). Thus AB is called a melting curve.

• There is only one point, A, at which all three phases of a substance-solid, liquid, and gas can coexist at equilibrium. This is called the triple point.

• At pressures below the triple-point pressure, the liquid phase does not exist; rather, the substance goes directly from solid to gas (sublimes) or the reverse happens (crystals deposit from the gas).

• At pressures and temperatures along AD, the sublimation curve, solid and vapor are in equilibrium.

• The critical temperature is the temperature at and above which a gas cannot be liquefied, that is, the temperature above which the liquid and gas do not exist as distinct phases.

• A substance at a temperature above its critical temperature is called a supercritical fluid.

• The critical pressure is the pressure required to liquefy a gas (vapor) at its critical temperature.

• The combination of critical temperature and critical pressure is called the critical point.