Liquids notes

General Properties of Liquids

A liquid has a definite volume but no definite shape. It takes the shape of the container in which it is placed. Liquid particles are close together, but they can move past one another because the intermolecular forces are not strong enough to fix them in a rigid arrangement.

Liquids are intermediate between gases and solids. Their particles have more freedom of movement than solid particles but are much closer together than gas particles. The physical properties of a liquid depend on the strength and distance of its intermolecular forces.

• Liquids have definite volume and indefinite shape.
• Liquids are almost incompressible because their particles are already close together.
• Liquids flow because their particles can move past one another.
• The rate of diffusion in liquids is lower than in gases and higher than in solids.
• Intermolecular forces become weaker as the distance between molecules increases.
• Liquid particles possess kinetic energy and potential energy due to intermolecular attractions.
• CHCl3 is liquid at room temperature because of dipole-dipole forces and London dispersion forces.

Surface Tension, Viscosity and Other Properties

Surface tension is the tendency of a liquid surface to behave like a stretched elastic membrane. Molecules at the surface experience a net inward attraction because they are pulled by molecules below and beside them. Surface tension makes small liquid drops nearly spherical.

Viscosity is the resistance of a liquid to flow. It results from friction between layers of liquid. Stronger intermolecular forces usually produce greater viscosity. Honey is more viscous than water and alcohol.

Cohesive forces act between molecules of the same substance, while adhesive forces act between molecules of different substances. The balance between these forces explains the shape of a liquid surface in a container.

• Surface tension decreases when temperature increases.
• Surface tension is generally greater when intermolecular forces are stronger.
• Viscosity is the resistance offered by a liquid to flow.
• Honey is more viscous than water.
• Viscosity usually increases with stronger intermolecular attraction and larger molecular size.
• Cohesion is attraction between similar molecules, such as water-water attraction.
• Adhesion is attraction between different substances, such as water and glass.
• Water forms a concave meniscus in a glass tube because adhesion between water and glass is greater than cohesion between water molecules.

Evaporation and Cooling

Evaporation is the escape of molecules from the surface of a liquid into the vapour phase. It occurs at all temperatures, not only at the boiling point. Molecules with higher kinetic energy escape first, so the average kinetic energy of the remaining liquid decreases.

Because the average kinetic energy decreases, evaporation causes cooling. This is why sweating cools the body. Evaporation is a surface phenomenon, while boiling occurs throughout the liquid.

• Evaporation takes place only at the surface of a liquid.
• Evaporation occurs at all temperatures.
• Evaporation causes cooling.
• Evaporation increases with increase in temperature.
• Evaporation increases when surface area increases.
• Evaporation increases when wind speed increases because vapour is removed from the surface.
• Evaporation decreases when humidity increases.
• Liquids with weak intermolecular forces evaporate more rapidly.

Vapour Pressure and Boiling Point

Vapour pressure is the pressure exerted by the vapour of a liquid when the vapour and liquid are in dynamic equilibrium at a fixed temperature. In this state, the rate of evaporation equals the rate of condensation.

A liquid boils when its vapour pressure becomes equal to the external pressure. At standard atmospheric pressure, the boiling point is called the normal boiling point. A decrease in external pressure lowers the boiling point, while an increase in external pressure raises it.

The vapour pressure of a liquid depends on temperature and the nature of the liquid. It does not depend on the surface area when temperature and equilibrium conditions remain constant.

• Vapour pressure increases when temperature increases.
• Vapour pressure does not depend on surface area at equilibrium.
• A liquid boils when its vapour pressure equals external pressure.
• The normal boiling point is the boiling point at 1 atmosphere pressure.
• Boiling point increases when intermolecular forces increase.
• A liquid with high vapour pressure usually has a low boiling point.
• During boiling, the temperature remains constant until the liquid has changed into vapour.
• Evaporation and boiling differ because evaporation is a surface process, whereas boiling occurs throughout the liquid.

Hydrogen Bonding and Intermolecular Forces

A hydrogen bond is an attraction involving a hydrogen atom covalently bonded to a highly electronegative atom, usually F, O or N, and a lone pair on an electronegative atom of another molecule. It is weaker than a normal covalent bond but stronger than ordinary London dispersion forces.

In water, the oxygen atom of one molecule forms a hydrogen bond with the hydrogen atom of another water molecule. The oxygen uses one lone pair to form this bond. In the FSc description, this interaction is represented as a coordinate covalent bond because both bonding electrons are supplied by the oxygen atom.

Hydrogen bonding explains the unusually high boiling point of water. It also affects viscosity, surface tension, solubility and the structure of ice. Hydrocarbons do not normally form hydrogen bonds because they contain neither O, N nor F bonded to hydrogen.

• Hydrogen bonding occurs in substances such as H2O, HF and NH3.
• The hydrogen-bond donor commonly contains an O-H, N-H or F-H bond.
• The hydrogen-bond acceptor provides a lone pair of electrons.
• In water, an oxygen atom forms a bond with the hydrogen atom of another molecule by using one lone pair of electrons.
• Hydrogen bonding is stronger than ordinary van der Waals forces.
• Hydrocarbons generally do not have a tendency to form hydrogen bonding.
• Water has a higher boiling point than HF because the hydrogen bonding network in water is stronger overall.
• Propanone is miscible with water because its polar carbonyl group accepts hydrogen bonds, and its small organic part does not prevent mixing.

Liquid Crystals, Solutions and Distillation

Liquid crystals are substances that flow like liquids but possess some ordered arrangement like crystals. Their properties are intermediate between those of ordinary liquids and crystalline solids. They show anisotropy, meaning that some physical properties vary with direction, whereas ordinary liquids are generally isotropic.

Azeotropic mixtures are mixtures that distil without change in composition. Their vapour has the same composition as the liquid mixture, so they cannot be separated completely by ordinary fractional distillation. Dilute hydrochloric acid forms a constant-boiling mixture at about 22 percent HCl, so boiling does not increase its concentration beyond this composition.

• Liquid crystals have fluidity like liquids and directional order like crystals.
• The properties of liquid crystals are intermediate between crystals and ordinary isotropic liquids.
• Anisotropic properties depend on the direction of measurement.
• Azeotropic mixtures distil without change in composition.
• The composition of an azeotropic vapour is the same as that of the liquid mixture at the azeotropic point.
• Dilute HCl cannot be concentrated beyond about 22 percent by ordinary boiling because it forms a constant-boiling mixture.
• Azeotropic mixtures are not completely separated by simple fractional distillation.
• Ordinary liquids generally show isotropy, meaning their properties are the same in all directions.

Solubility, Distribution Coefficient and Henry's Law

Solubility is the maximum amount of a solute that dissolves in a specified amount of solvent at a given temperature. A substance dissolves best in a solvent with similar polarity. Polar substances generally dissolve in polar solvents, while non-polar substances dissolve in non-polar solvents.

The distribution coefficient is the ratio of the concentration or amount of a solute in two immiscible solvents at equilibrium. In extraction, it may be expressed as the ratio of the amount of solute in an organic solvent to its amount in an aqueous solvent.

Henry's law describes the solubility of a gas in a liquid. At constant temperature, the solubility of a gas is directly related to its partial pressure above the liquid. It can be written as p = KH x, where p is the partial pressure, KH is Henry's law constant and x is the mole fraction of the gas in solution.

• The distribution coefficient is the ratio of solute concentration in two immiscible solvents.
• For organic and aqueous solvents, it is commonly written as amount in organic solvent divided by amount in aqueous solvent.
• Henry's law: p = KH x.
• At constant temperature, increasing gas pressure increases its solubility in a liquid.
• A larger Henry's law constant generally indicates lower solubility for the gas under the same conditions.
• For nitrogen at 20 °C, p = 742.5 torr and KH = 5.75 x 10^7 torr, the dissolved amount is approximately 16 mL at STP per litre of water.
• Gas solubility usually decreases when temperature increases.

Anomalous Behaviour of Water

Water shows unusual behaviour because of extensive hydrogen bonding. Most liquids contract continuously on cooling, but water expands when cooled below 4 °C. Water has its maximum density at 4 °C, so ice is less dense than liquid water and floats on it.

At 0 °C, ice and liquid water can coexist in equilibrium. The freezing point and melting point of pure water at standard pressure are both 0 °C. During melting or freezing, the temperature remains constant while the phase change takes place.

The open structure of ice contains empty spaces produced by hydrogen bonding. When ice melts, some of these spaces disappear, so the molecules become more closely packed. Further heating above 4 °C causes ordinary thermal expansion, and the density decreases.

• The freezing point of pure water at standard pressure is 0 °C.
• The melting point of pure ice at standard pressure is 0 °C.
• At 0 °C, water may exist as both ice and liquid water.
• Water has maximum density at 4 °C.
• Water expands when cooled from 4 °C to 0 °C.
• Ice is less dense than liquid water and therefore floats on water.
• The anomalous expansion of water is due to its hydrogen-bonded open structure.
• Water boils at 100 °C at 1 atmosphere pressure, and its temperature remains constant during boiling.

Humidity and Numerical Relationships

Humidity is the amount of water vapour present in air. Relative humidity compares the actual partial pressure of water vapour with the vapour pressure of water at the same temperature. It is expressed as a percentage.

Pressure values must be converted into the same units before calculation. Since 1 atmosphere equals 760 mm Hg, 0.01876 atmosphere is approximately 14.26 mm Hg. Dividing this by 23.76 mm Hg and multiplying by 100 gives approximately 60 percent relative humidity.

• Relative humidity = (partial pressure of water vapour / vapour pressure of water) x 100.
• At 25 °C, if water vapour pressure is 23.76 mm Hg and actual vapour pressure is 0.01876 atmosphere, relative humidity is approximately 60 percent.
• Higher humidity reduces the rate of evaporation.
• Saturated air has relative humidity of 100 percent.
• Relative humidity is affected by temperature because the vapour pressure of water changes with temperature.
• A liquid evaporates more slowly when the surrounding air already contains more water vapour.