Intermolecular ForcesIntermolecular Forces

Intermolecular ForcesIntermolecular Forces

• The attractive forces holding solids and liquids together are called intermolecular forces. The covalent bond holding a molecule together is an intramolecular force

• Intermolecular forces are much weaker than intramolecular forces (chemical bonds) (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl).

• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).

• When a substance condenses intermolecular forces are formed.

Intermolecular ForcesIntermolecular Forces

Intermolecular ForcesIntermolecular Forces

Dipole-Dipole ForcesDipole-Dipole Forces• Molecules orient themselves to

maximaize the + --- - interactions and minimize the + --- + and - --- -forces.

• There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble.

• If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

Intermolecular ForcesIntermolecular Forces

London Dispersion ForcesLondon Dispersion Forces• Weakest of all intermolecular forces.• Exist primarily between noble gases and nonpolar

molecules• The nucleus of one molecule (or atom) attracts the

electrons of the adjacent molecule (or atom).• For an instant, the electron clouds become distorted.• In that instant a dipole is formed (called an

instantaneous dipole).• Relatively weak and short-lived

Intermolecular ForcesIntermolecular Forces

London Dispersion ForcesLondon Dispersion Forces

Intermolecular ForcesIntermolecular Forces

London Dispersion ForcesLondon Dispersion Forces• One instantaneous dipole can induce another instantaneous dipole in

an adjacent molecule (or atom).• Polarizability is the ease with which an electron cloud can be

deformed. • The larger the molecule (the greater the number of electrons) the

more polarizable.• London dispersion forces increase as molecular weight increases.• London dispersion forces exist between all molecules.• The greater the surface area available for contact, the greater the

dispersion forces.• London dispersion forces between spherical molecules are lower

than between sausage-like molecules.

Intermolecular ForcesIntermolecular ForcesHydrogen BondingHydrogen Bonding• Special case of dipole-dipole forces.• By experiments: boiling points of compounds with H-

F, H-O, and H-N bonds are abnormally high.• Intermolecular forces are abnormally strong.• H-bonding requires H bonded to an electronegative

element (most important for compounds of F, O, and N).– Electrons in the H-X (X = electronegative element) lie much

closer to X than H.

– H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X.

– Therefore, H-bonds are strong.

Intermolecular ForcesIntermolecular ForcesHydrogen BondingHydrogen Bonding

Intermolecular ForcesIntermolecular ForcesHydrogen BondingHydrogen Bonding• Hydrogen bonds are responsible for:

– Ice Floating• Solids are usually more closely packed than liquids;

• therefore, solids are more dense than liquids.

• Ice is ordered with an open structure to optimize H-bonding.

• Therefore, ice is less dense than water.

• In water the H-O bond length is 1.0 Å.

• The O…H hydrogen bond length is 1.8 Å.

• Ice has waters arranged in an open, regular hexagon.

• Each + H points towards a lone pair on O.

• Ice floats, so it forms an insulating layer on top of lakes, rivers, etc. Therefore, aquatic life can survive in winter.

Intermolecular ForcesIntermolecular ForcesHydrogen BondingHydrogen Bonding

Hydrogen bonds are responsible for:– Protein Structure

• Protein folding is a consequence of H-bonding.

• DNA Transport of Genetic Information

Intermolecular ForcesIntermolecular ForcesComparing Intermolecular ForcesComparing Intermolecular Forces

Some Properties of LiquidsSome Properties of Liquids

ViscosityViscosity• 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.

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

attracted to their neighbors. 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.

Some Properties of LiquidsSome Properties of Liquids

Surface TensionSurface Tension

Phase ChangesPhase Changes• Surface molecules are only

attracted inwards towards the bulk molecules.

• Sublimation: solid gas.

• Vaporization: liquid gas.

• Melting or fusion: solid liq.

• Deposition: gas solid.

• Condensation: gas liquid.

• Freezing: liquid solid.

Energy Changes Accompanying Phase Changes

•Sublimation: Hsub > 0 (endo)

•Vaporization: Hvap > 0 (endo)

•Melting or Fusion: Hfus > 0 (endo)

•Deposition: Hdep < 0 (exo)

•Condensation: Hcon < 0 (exo)

•Freezing: Hfre < 0 (exo)

Phase ChangesPhase Changes

Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes• All phase changes are possible under the right

conditions (e.g. water sublimes when snow disappears without forming puddles).

• The sequence

heat solid melt heat liquid boil heat gas

is endothermic.• The sequence

cool gas condense cool liquid freeze cool solid

is exothermic.

Phase ChangesPhase Changes

Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes

Phase ChangesPhase ChangesHeating CurvesHeating Curves• Plot of temperature change versus heat added is a heating curve.

• During a phase change, adding heat causes no temperature change.

– These points are used to calculate Hfus and Hvap.

Vapor PressureVapor Pressure

Explaining Vapor Pressure on the Molecular Explaining Vapor Pressure on the Molecular LevelLevel

• 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 PressureVapor Pressure

Explaining Vapor Pressure Explaining Vapor Pressure on the Molecular Levelon the Molecular Level

• Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface.

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

Vapor PressureVapor Pressure

Volatility, Vapor Pressure, and TemperatureVolatility, Vapor Pressure, and Temperature• If equilibrium is never established then the liquid evaporates.

• Volatile substances evaporate rapidly.

• The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

Vapor PressureVapor PressureVapor Pressure and Boiling PointVapor Pressure and Boiling Point• Liquids boil when the external pressure equals the

vapor pressure.• Temperature of boiling point increases as pressure

increases.• Two ways to get a liquid to boil: increase temperature

or decrease pressure.– Pressure cookers operate at high pressure. At high

pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked.

• Normal boiling point is the boiling point at 760 mmHg (1 atm).

Phase DiagramsPhase Diagrams• Phase diagram: plot of pressure vs. Temperature

summarizing all equilibria between phases.• Given a temperature and pressure, phase diagrams

tell us which phase will exist.• Features of a phase diagram:

– Triple point: temperature and pressure at which all three phases are in equilibrium.

– Vapor-pressure curve: generally as pressure increases, temperature increases.

– Critical point: critical temperature and pressure for the gas.

– Normal melting point: melting point at 1 atm.

Phase DiagramsPhase Diagrams• Any temperature and pressure combination not on a

curve represents a single phase.

Phase DiagramsPhase DiagramsThe Phase Diagrams of HThe Phase Diagrams of H22O and COO and CO22

•Triple point at 0.0098C and 4.58 mmHg.•Normal melting (freezing) point is 0C.•Normal boiling point is 100C.•Critical point is 374C and 218 atm

•Triple point at -56.4C and 5.11 atm.•Normal sublimation point is -78.5C. •(At 1 atm CO2 sublimes it does not melt.)•Critical point occurs at 31.1C and 73 atm.