Chemistry: Liquids and Intermolecular Forces

11.1 Molecular Comparison of States of Gases, Liquids, and Solids

• The state of a substance depends on the balance of particle kinetic energy and the energy of attraction.
  • Kinetic energy keeps particles apart and moving.
  • Strong attractive forces bring molecules close together.
• Temperature is related to the average kinetic energy.
• Decreasing temperature brings the particles closer together.

Properties and State of Matter

• Gas
  • Assumes both volume and shape of its container
  • Expands to fill its container
  • Is compressible
  • Flows readily
  • Diffusion within a gas occurs rapidly
• Liquid
  • Assumes shape of portion of container it occupies
  • Does not expand to fill its container
  • Is virtually incompressible
  • Flows readily
  • Diffusion within a liquid occurs slowly
• Solid
  • Retains own shape and volume
  • Does not expand to fill its container
  • Is virtually incompressible
  • Does not flow
  • Diffusion within a solid occurs extremely slowly

11.2 Intermolecular Forces

• The attractions between molecules, intermolecular forces, are not nearly as strong as the intramolecular attractions, covalent bonds, that hold compounds together.
• Many physical properties reflect intermolecular forces, like boiling points, melting points, viscosity, surface tension, and capillary action.

Relative Strength of Attractions

• Chemical bonds
  • Ionic bonds: Lithium fluoride (LiF), Melting Point = 1118 K, Boiling Point = 1949 K
  • Metallic bonds: Beryllium (Be), Melting Point = 1560 K, Boiling Point = 2742 K
  • Covalent bonds: Diamond (C), Melting Point = 3800 K, Boiling Point = 4300 K
• Intermolecular forces
  • Dispersion forces: Nitrogen (N_2), Melting Point = 63 K, Boiling Point = 77 K
  • Dipole–dipole interactions: Hydrogen chloride (HCl), Melting Point = 158 K, Boiling Point = 188 K
  • Hydrogen bonding: Hydrogen fluoride (HF), Melting Point = 190 K, Boiling Point = 293 K
• Intermolecular attractions are weaker than bonds.
• Hydrogen bonds are NOT chemical bonds.

Types of Intermolecular Force Between Neutral Molecules

• Weakest to strongest forces:
  • Dispersion forces (also known as London dispersion forces or induced dipole-induced dipole interactions)
  • Dipole–dipole forces
  • Hydrogen bonding (a special dipole–dipole force)
• The first two types are also referred to collectively as van der Waals forces.

Dispersion Forces

• A nonpolar particle (helium atoms below) can be temporarily polarized to create dispersion force.
• The tendency of an electron cloud to distort is called its polarizability.
• The easier an electron cloud can be distorted, the stronger the dispersion forces.
• Factors include:
  • Number of electrons in an atom (more electrons, more dispersion force).
  • Size of atom or molecule/molecular weight (larger, more dispersion force).
  • Shape of molecules with similar masses (more compact, less dispersion force).
• If a substance is easy to polarize, it has a lower boiling point, i.e., Ne (smaller) versus (larger).
• The lower boiling point indicates weaker intermolecular forces.
• You can predict the relative order of physical properties of neutral substances by molecular weight.

Dipole–Dipole Interactions

• Polar molecules have a more positive and a more negative end—a dipole
• The oppositely charged ends attract each other.
• For molecules of approximately equal mass and size, the more polar the molecule, the higher its boiling point.

Which Have a Greater Effect: Dipole–Dipole Interactions or Dispersion Forces?

• If two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force.
• If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

Hydrogen Bonding

• The boiling points of binary compounds between hydrogen and Groups 4, 5, 6, and 7 elements follow a trend.
  • As molecular weight increases, boiling point increases.
• However, Period 3 elements have much higher boiling points than expected.
• Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.
• These atoms interact with a nearly bare hydrogen nucleus (which contains one proton but NO inner electrons).
• The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
• These interactions are called hydrogen bonds.
• A hydrogen bond is an attraction between a hydrogen atom attached to a highly electronegative atom and a nearby small electronegative atom in another molecule or chemical group.

Ice Compared to Liquid Water

• Hydrogen bonding makes the molecules farther apart in ice than in liquid water.
• The density of ice is less than water so ice floats on water.

Ion–Dipole Forces

• Ion–dipole interactions are found in solutions of ions.
• The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents, like water.

Determining Intermolecular Forces in a Substance

• Note that ALL substances exhibit dispersion forces.
• The strongest force dictates the extent of attractions between molecules.

Generalizations About Relative Strengths of Intermolecular Forces

• When two molecules have comparable molar masses and shapes, dispersion forces are roughly equal.
• When two molecules have very different molar masses and there is no H-bonding, dispersion force determines the substance with stronger attractions.
• For more complex molecules, there is often a weighting of individual intermolecular forces.

11.3 Select Properties of Liquids

• Liquid properties affected by intermolecular forces:
  • Boiling point (discussed above)
  • Melting point
  • Viscosity (discussed here)
  • Surface tension (discussed here)
  • Capillary action
  • Heats of vaporization
  • Heats of fusion

Viscosity

• Resistance of a liquid to flow is called viscosity.
• It is related to the ease with which molecules can move past each other.
• Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

Viscosities of a Series of Hydrocarbons at 20°C

Surface Tension

• Water acts as if it has a “skin” on it due to extra inward forces on its surface.
• It causes water to “bead up” when in contact with nonpolar surfaces.
• Those forces are called surface tension.

Cohesion and Adhesion

• Intermolecular forces that bind similar molecules to one another are called cohesive forces.
• Intermolecular forces that bind a substance to a surface are called adhesive forces.
• These forces are important in capillary action.
  • Enable plants to move nutrients upward against gravity.
  • Paper towels (cellulose based) absorb liquid.

Capillary Action

• The rise of liquids up narrow tubes is called capillary action.
• Water has stronger adhesive forces with glass. Adhesive forces attract the liquid to the wall of the tube, i.e., concave meniscus.
• Mercury has stronger cohesive forces with itself. Cohesive forces attract the liquid to itself, i.e. convex meniscus.

11.4 Phase Changes

• Conversion from one state of matter to another is called a phase change.
• Energy is either added (endothermic) or released (exothermic) in a phase change.
• Phase changes: melting/freezing, vaporizing/condensing, subliming/depositing.

Energy Change Accompanies Phase Change

• The heat of fusion is the energy required to change a solid at its melting point to a liquid.
• The heat of vaporization is the energy required to change a liquid at its boiling point to a gas.
• The heat of sublimation is the energy required to change a solid directly to a gas.

Heating Curves

• A graph of temperature versus heat added is called a heating curve.
• When a phase changes (solid, liquid, or gas), temperature changes. The enthalpy change (heat, q) is the product of specific heat, sample mass, and temperature change.
• During the phase change, the temperature of the substance does not change, i.e., melting pointing and boiling point. The enthalpy change (heat, q) is the product of mass or moles and the heat of fusion or vaporization.

Supercritical Fluids

• Gases normally liquefy when pressure is applied.
• The temperature beyond which a gas cannot be compressed is called its critical temperature. The pressure needed to compress the liquid at critical temperature is called critical pressure.
• The state beyond this temperature is called a supercritical fluid.
• Carbon dioxide is commonly used in the supercritical state as a solvent.

Critical Temperatures and Pressures of Selected Substances

11.5 Vapor Pressure

• At any temperature, some liquid molecules have enough energy to escape the surface and become a gas.
• As the temperature rises, the fraction of molecules that have enough energy to break free increases.
• As more molecules escape the liquid, the pressure molecules exert increases.
• The liquid and vapor reach a state of dynamic equilibrium: the vapor pressure. Liquid molecules evaporate and vapor molecules condense at the same rate, i.e., the steady state.
• The boiling point of a liquid is defined as the temperature at which its vapor pressure equals atmospheric pressure.
• The normal boiling point is the temperature at which its vapor pressure is 760 torr (1 atm).
• Pressure is inversely proportional to temperature.
• Plotting the natural log of the vapor pressure (K) yields a straight line.
• This relationship is quantified in the Clausius Clapeyron equation.

11.6 Phase Diagrams

• A phase diagram is a graph showing states of matter under conditions of temperature and pressure.
• It also shows changes of state and the triple point (all phases present) and critical point (supercritical fluid).

Phase Diagram of Water

• The melting curve slants left.
  • As pressure increases, melting point decreases.
  • Density: (s) < (l). Ice floats.
• High critical temperature and critical pressure due to hydrogen bonding.
• Freeze-dry foods: pressure below 0.0063 atm.

Phase Diagram of Carbon Dioxide

• Solid sublimes at room temperature and normal atmospheric pressure.
• The melting curve slants right.
  • As pressure increases, melting point increases.
  • Density: (s) > (l).

11.7 Liquid Crystals

• Some substances do not go directly from the solid state to the liquid state.
• In this intermediate state, liquid crystals have some traits of solids and some of liquids.
• Molecules in liquid crystals have some degree of order.
• Used as sensors and device displays (LCD).
• In nematic liquid crystals, molecules are ordered in only one dimension, along the long axis.
• In smectic liquid crystals, molecules are ordered in two dimensions, along the long axis and in layers.
• In cholesteric liquid crystals, nematic-like crystals are layered at angles to each other.