Intermolecular Forces

Mr. Nelson AP Chem

Intermolecular Forces

Forces that exist between moleculesAlso called Van der Waals forces

Compare to intramolecular forces, which are the forces that hold atoms togetherHolds individual molecules together, but does not determine macro properties (aka boiling point)

Intermolecular Forces

Largely determines physical properties of solids and liquids

Intermolecular forces are weaker than intramolecular forces (chemical bonding)

Boiling points and melting points are an indication of the strength of intermolecular forces

Ion-Dipole forces

Interaction between an ion and a partial charge in a polar moleculeNot technically a van der Waals force

Positive ions are attracted to negative end of a dipoleForce depends on dipole moment of polar molecule and charge of the ion

Example: NaCl in Water

Dipole-dipole force

Attractive forces between polar moleculesWeaker than ion-dipole forces

Positive end of one molecule is near negative end of another

In liquids, molecules of equal size and mass have increasing intermolecular attractions with increasing polarity

London Dispersion Forces

London proposed that the motion of electrons in an atom or molecule can create an instantaneous dipole momentExample: If both electrons on an He atom are on the same side of a nucleus, you have an instantaneous dipole moment

London Dispersion Forces

A temporary dipole on one atom can induce a similar dipole on an adjacent atom (called an induced dipole)

London Dispersion Forces

Because of induced dipoles, molecules attract one another

Dispersion forces exist in all types of molecules (polar or non-polar, charged or not)

In many cases, dispersion forces can be stronger than dipole-dipole

Example: CH3F (B.P. -78.9 °C) vs. CCl4 (B.P. 76.5 °C)

London Dispersion Forces

Dispersion forces increase with molar mass of the moleculeMore electrons allows for greater chances of dispersion

When comparing molecules of similar weights and shapes, dipole-dipole forces tend to be the decisive factor

When comparing differing weights, dispersion forces is the decisive factor

Intermolecular Forces Recap

Example

Of Br2, Ne, HCl, HBr, and N2, which has:

The largest intermolecular dispersion forces?

The largest dipole-dipole attractive forces?

Hydrogen Bonding

A special type of dipole-dipole attraction that exists between hydrogen atoms in a polar bond and small electronegative ions

Can only form between hydrogen and N, O, or F

Hydrogen Bonding

Water has a lower molar mass and a higher boiling point

Dispersion forces do not account for this

Hydrogen Bonding

Explains density of ice

Solid is less dense than liquid (not common)

Review of IM forcesIon-Dipole Forces•Exist when ions and dipole present•Stronger than simply dipole-dipole

Dipole-Dipole Forces•Between polar molecules only•Larger dipole = larger force

Dispersion Forces•Between all molecules•Higher molar mass = larger force

Hydrogen Bonding•Strongest IM Force•Exists only between H with either N, O, or F

Phase Changes

Changes in state of matterIn general, each state of matter can change into either of the two other states

Phase changes require energy

When becoming a more disordered state, requires energy to overcome intermolecular forces that hold them together

Phase Changes

Phase Changes

Melting, vaporization, and sublimation are all endothermic processes

Freezing, condensation, and deposition are all exothermic processes

Phase Changes

FusionMelting of a solidMolar heat of fusion or enthalpy of fusion (∆Hfus)

∆Hfusion = -∆Hsolidification

VaporizationVaporizing of a liquidMolar heat of vaporization or enthalpy of vaporization (∆Hvap)

∆Hvaporization = -∆Hcondensation

Heating Curves

Graph of temperature of the system versus heat added to the system

Some segments of the graph are heating as single phase others are converting one phase to another

Heat change when temp. inc. is given by q=mC∆T. During phase change, it is q=n∆H (n is moles of substance)

Heating Curve of Water

Phase Change Problem

Freon-11 (CCl3F) has a normal boiling point of 23.8 °C. The specific heat of CCl3F (l) is 0.87 J/g K and CCl3F (g) is 0.59 J/g K. The heat of vaporization is 24.75 kJ/mole. Calculate the heat required to convert 10.0 g of Freon-11 from liquid at -50.0 °C to gas at 50.0 °C.

Vapor Pressure

In a closed system, liquid will initially evaporate and then condense back into its liquid form

The pressure exerted by this liquid/gas equilibrium is called vapor pressure

When evaporation occurs in an open system, no equilibrium can be established

Vapor Pressure

Substances with high vapor pressure evaporate more easily than those with a low vapor pressure

Volatile liquids (aka liquids that evaporate easily) have high vapor pressures

Water can vaporize at room temp. because molecules in liquid move at different speeds and some can overcome IM forces (think space shuttles/escape velocity)

Vapor Pressure

As the number of gas molecules increases in a closed system, the probability that gas molecules will strike the liquid increases.

This allows the molecules to become “trapped” by the IM forces of the liquid

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/vaporv3.swf

Temperature & Vapor Pressure

Vapor pressure increases nonlinearly with temperature

More molecules have enough KE to escape into gas phase

Vapor Pressure & Boiling Point

Boiling point is when the vapor pressure equals the external pressure acting on the surfaceA “normal” boiling point is measured at 1 atm

A more volatile liquid has a lower boiling point and a higher vapor pressure (and vice versa)

Phase Diagrams

A graph that displays the conditions under which an equilibrium exists between different states of matter

Allows prediction of a substance’s state of matter at a given temperature and pressure

Phase Diagrams

Critical Temperature (Tc) is the highest temperature a substance can exist as a liquid

Meaning above Tc molecular motion breaks IM attraction (Tc measures strength of IM forces)

Critical Pressure (Pc) is pressure required to bring about liquefaction at the critical temperature

Phase Diagrams

Triple Point is the place at which all three phases exist at equilibrium