Unit 9 – Intermolecular Forces

Unit 9 – Intermolecular Forces

Liquids, Solids, and Phase Changes Chapters 11 and 12

Intermolecular Forces• Interaction of molecules with other

molecules• Weaker forces than ionic or covalent

bonds• Less energy is required to overcome

the force of an intermolecular attraction than a chemical bond.– Examples of “breaking” intermolecular

forces: – Melting Point– Boiling Point– The stronger the IMF, the higher the

melting point/boiling point.

Types of Intermolecular Forces• Ion-Dipole Forces• Dipole-Dipole Forces• London Dispersion Forces• Hydrogen Bonding

• All are considered electrostatic in nature (attractions between + and – species)

Ion-Dipole Forces• Exist between an ion and the

partial charge on the end of a polar molecule

• Remember that polar molecules are dipoles and have a dipole moment+

-

+-

+

Dipole-Dipole Forces• The positive end of one polar

molecule is attracted to the negative end on another polar molecule.

• Effective only when polar molecules are close together and in correct orientation

• Weaker than ion-dipole forces• Dipole-dipole interactions involve

both attractive and repulsive forces• Molecules of equal mass and size

exhibit stronger intermolecular attractions with increasing polarity

Hydrogen Bonding• Attraction between a hydrogen

atom in a polar bond and an unshaired electron pair on a nearby small electronegative ion or atom (generally F, O, or N)

• A strong intermolecular attraction that accounts for high m.p. and b.p., specific heat, and heat of vaporization

• Sometimes considered to be a dipole-dipole attraction

• Strong because protons are more massive in size

• Reason ice is less dense than water

CH4

SiH4

GeH4SnH4

PH3

NH3 SbH3

AsH3

H2O

H2SH2Se

H2TeHF

HIHBr

HCl

Boiling Points

0ºC

100

-100

200

Water

d+d-

d+

London Dispersion Forces• Attraction based on electron location

of one molecule to electron location on a different molecule

• Non - polar molecules also exert forces on each other – otherwise, there would be no solids or liquids.

• Instantaneous dipole moment can be induced then attraction of opposite dipoles

• More polarizable molecules (“squishy clouds”) have greater L-D forces

• This force increases with larger molecule size (thus larger molecular weights)

London Dispersion Forces

• Weak, short-lived.• Lasts longer at low temperature.• Explains the trend for bigger

molecules, higher melting and boiling points.

• Much, much weaker than other forces.

• Also called Van der Waal’s forces.

Example

• Vibrations of the nucleus within the electron charge cloud creates a temporary dipole, shown here as delta + and delta - charges.

• This causes sympathetic vibrations of the nucleus in neighboring atoms, so that the delta + and delta - charges become arranged opposite one another, plus near minus.

• There is now attraction between the opposite delta charges on neighboring atoms, creating Van der Waal's forces.

• Eventually all of the atoms or molecules vibrate in unison, making the partial charges enough to hold the structure together.

• When enough energy is given to a substance held together by only Van der Waal's forces, they are overcome and the compound melts.

Represented above are five identical balloons, each filled to the same volume at 25°C and 1.0 atm pressure with the pure gas indicated. For each question, be sure to explain your reasoning.(a) Which balloon contains the greatest mass of gas?(b) Compare the average kinetic energies of the gas molecules in the balloons.(c) Which balloon contains the gas that would be expected to deviate most from the behavior of an ideal gas? (d) Twelve hours after being filled, all the balloons have decreased in size. Predict which balloon will be the smallest.

Comparing Intermolecular Forces

Interacting molecules/ions

Are ions involved?

Are Polar Molecules Involved?

London DispersionAr, I2

Are polar molecules AND ions both present?

Ionic BondNaCl

Ion-dipole ForcesKBr in H2O

Are H atoms bonded to O, N, or F?

Dipole-DipoleH2S, CH3Cl

Hydrogen BondingH2O, NH3, HF

ON

YES

ON

YES YES ON

ON

YE

S

Properties of Liquids – Many of the properties are due to internal attraction of atoms.

• Viscosity – the resistance of a liquid to flow.– The greater the viscosity, the slower it

flows– Can be measured as a rate of flow

through a metal tube or the rate at which an object falls through the liquid.

– Viscosity increases with an increase in molecular weight

• Large IMFs = more viscous liquid• Large molecules can get tangled

up.• Cyclohexane has a lower viscosity

than hexane because it is a circle- more compact.

Surface Tension• Energy required to increase the

surface area of a liquid by a unit amount– Inward forces that must be overcome– Atoms at surface are closer together– Forces that bind similar molecules

(such as H-bonding) are called cohesive forces

– Forces that bind a substance to a surface are called adhesive forces

Surface tension

• Molecules in the middle are attracted in all directions.

Molecules at the the top are only pulled inside.

Minimizes surface area.

Example:• Water in a glass adheres to the

glass because the adhesive force is larger than the cohesive forces of the H2O molecules

• Mercury is the opposite

Capillary Action• Liquids spontaneously rise in a

narrow tube• Occurs when a small tube is placed

in H2O and the water rises into the tube

• Glass is polar.• It attracts water molecules

Beading

• If a polar substance is placed on a non-polar surface. – There are cohesive,– But no adhesive

forces.• And Visa Versa

Phase Changes• Every phase change is

accompanied by a change in the energy of the system

• ∆Hfus (heat of fusion) – molecules of a solid free up, having more ability to move.

• ∆Hvap (heat of vaporization) – molecules of a liquid break free

• ∆Hvap is almost always larger in comparison to ∆Hfus

• ∆Hsub (heat of sublimation) – molecules transition from a solid to a gas

• ∆Hsub = ∆Hfus + ∆Hvap

Phase Diagrams• A graphical way to summarize the

conditions under which equilibria exist between different states of matter

• Predicts the phase of a a substance at a given temperature and pressure

Change in melting point

Boiling Point Variance

All three phases are in equilibrium at this temperature and pressure

Interpreting a Phase Diagram• Critical Temperature: The highest

temperature at which a distinct liquid phase can form

• Critical Pressure: pressure required to bring about liquefaction at this critical temperature

• Nonpolar, low molar mass substances have lower values for critical temperature and pressure

Vapor Pressure• The pressure exerted by a vapor

in the space above a liquid• Weaker intermolecular forces =

higher vapor pressure• Dynamic Equilibrium: rates of

changing from gas to liquid and liquid to gas is the same

• Volatile Liquids: substances with higher vapor pressure that allow liquids to evaporate to dryness (no dynamic equilibrium exists)

Boiling• Liquids boil when the vapor

pressure equals the external pressure acting on the surface of the liquid.

• Bubbles of vapor form within the interior of the liquid

• Lower pressure = lower boiling point

• Vapor pressure and temperature are related through The Clausius-Clapeyron Equation

where C is a constant

Structures of Solids• Crystalline Solid – atoms, ions, or

molecules are ordered in well-defined arrangements– Have flat surfaces, or faces– Definite angles– Regular shapes – “structured”

• Amorphous Solid – solid whose particles have no orderly structure– Lack of well-defined faces and shapes– Mixtures of molecules that do not

stack together well– Generally composed of large,

complicated molecules

• Crystalline solids – melts at a specific temperature

• Amorphous solids – soften over a wide range of temperatures.

Warm-Up1. Identify the most important types

of inter-particle (between “molecules”) forces present in the solids of each of the following substances

• a) BaSO4

• b) H2S • c) Xe • d) CHCl3

• e) NH3

Warm-Up, cont.2. Predict which substance from the

following pairs would have the greater intermolecular forces. (you WILL need to draw the Lewis structures to answer correctly)

• a) CO2 or OCS

• b) PF3 or PF5

• c) SO3 or SO2

• d) HF or HBr • e) CH3OH or H2CO

• f) CH3CH2CH2NH2 or H2NCH2CH2NH2