Intermolecular Forces and Liquids

Intermolecular Forces and Liquids

Kinetic Molecular Theory

• According to the Kinetic Molecular Theory, ALL particles of matter are in constant motion.

• This theory helps explain the behavior of solids, liquids, and gases.

Behavior of gases

• Particles in a gas are never at rest.

• Gaseous atoms travel in a straight line until it collides with either another atom or the wall of the container.

• The constant motion of gas particles allow it to fill a container of any shape or size.

Behavior of liquids

• Particles in a liquid are more closely packed than the particles in a gas.

• Therefore, attractions between the particles in a liquid do affect the movement of the particles.– It slows them down – less kinetic energy

• A liquid takes the shape of its container because particles can flow to new locations.

• The volume is constant because forces of attraction keep the particles close together.

Behavior of solids

• Solids have a definite volume and shape because their particles vibrate around fixed locations.

• Strong attractions restrict motion and keep each atom in a fixed location relative to its neighbors.

• Atoms vibrate around its location but it does not exchange places with neighboring atoms.

Phase changes

• The reversible physical change that occurs when a substance changes from one state of matter to another.

Forces that hold atoms and molecules together

• Ionic Bond: due to electrostatic attraction between opposite charges (ionic compounds).

• Covalent Bond: due to combining of atomic orbitals when electrons are shared.

• Intermolecular Forces: due to electrostatic attraction between opposite charges (ionic and covalent compounds).

Summary of Intermolecular Forces

Van der Waals intermolecular forces:• Ion-dipole forces• Dipole-dipole forces

– Special dipole-dipole force: hydrogen bonds• Forces involving nonpolar molecules: induced

forces• Dispersion or London forces.

Two factors affect the strength of ion or dipole force (Coulomb’s law)

• Magnitude of charge• Distance

Generally – order of strength

Strongest: Ionic Bonds Ion-dipole bonds

Hydrogen Bonding Dipole forces Induced dipole

Weakest: Dispersion forcesReflects in properties

Comparing Properties

Melting Point, Boiling Point, Heat of Fusion, Heat of Vaporization, Surface tension, viscosity - - - all go up as strength of IMF increases

Vapor pressure – goes down as strength of IMF increases (higher IMF makes evaporation less likely and less gas molecules mean lower vapor pressure)

Ion-Ion Forcesfor comparison of

magnitude

Na+—Cl- in saltThese are the

strongest forces. Lead to solids with

high melting temperatures.

NaCl, mp = 800 oCMgO, mp = 2800 oC

Attraction Between Ions and Permanent Dipoles

Attraction between ions and dipole depends on ion charge and ion-dipole distance.

Dipole-Dipole ForcesDipole-Dipole Forces

Dipole-dipole forces bind molecules having permanent dipoles to one

another.

Dipole moment

• Polar molecules or dipoles have positive and negative side.

• A measure for the polarity is the dipole moment (p. 376).

• Dipole moment is given in Debye (D).

Hydrogen BondingHydrogen BondingA special form of dipole-dipole attraction, which enhances dipole-dipole attractions.

H-bonding occurs when X and Y are N, O, or F

Watch out!

Hydrogen must be connected to F, O, or N. Not all hydrocarbons with these atoms present contain hydrogen bonding.

Example:Dimethyl ether: CH3 – O – CH3

does NOT have hydrogen bonding.Methanol: CH3OH

does have hydrogen bonding.

H-Bonding Between Methanol and WaterH-Bonding Between Methanol and Water

H-bondH-bondH-bondH-bond

Hydrogen Bonding in H2O

Hydrogen Bonding in H2OH-bonding is

especially strong in water because

• the O—H bond is very polar

• there are 2 lone pairs on the O atom

Accounts for many of water’s unique properties.

Hydrogen Bonding in H2O

Hydrogen Bonding in H2O Ice has open

lattice-like structure.

Ice density is < liquid.

And so solid floats on water.

A consequence of

hydrogen bonding

H-bonding leads to abnormally high specific heat capacity of water (4.184 J/g•K)

This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

DNA Base-Pairing through H-BondsDNA Base-Pairing through H-Bonds

Boiling point of water

• Water has a very high boiling point compared to other simple hydrogen containing compounds, such as CH4, H2S, and even NH3 and HF (which exhibit H-bonding as well, but less “extreme” as H2O

Boiling Points of Simple Hydrogen-

Containing Compounds

FORCES INVOLVING INDUCED DIPOLESFORCES INVOLVING INDUCED DIPOLES

How can non-polar molecules such as O2 and I2 dissolve in water?

The water dipole INDUCES a dipole in the O2 electric cloud.

Dipole-induced dipoleDipole-induced dipole

FORCES INVOLVING INDUCED DIPOLES

FORCES INVOLVING INDUCED DIPOLES

Solubility increases with the mass of the gas

Process of inducing a dipole is polarizationDegree to which electron cloud of an atom or molecule can be distorted in its polarizability.

Dispersion forces

• Induced dipole – induced dipole• Between 2 non-polar molecules.• Caused by the movement of the electron

cloud.• The more electrons (higher molar mass) the

stronger the force (higher MP and BP).

FORCES INVOLVING INDUCED DIPOLES

FORCES INVOLVING INDUCED DIPOLES

Formation of a dipole in two nonpolar I2 molecules.

Induced dipole-induced dipole(dispersion)

Induced dipole-induced dipole(dispersion)

FORCES INVOLVING INDUCED DIPOLES

FORCES INVOLVING INDUCED DIPOLES

The induced forces between I2 molecules are

very weak, so solid I2 sublimes (goes from a solid to gaseous molecules).

FORCES INVOLVING INDUCED DIPOLES

FORCES INVOLVING INDUCED DIPOLES

The magnitude of the induced dipole depends on the tendency to be distorted.

Higher molar mass = stronger forces Molecule Boiling Point (oC) CH4 (methane) - 161.5

C2H6 (ethane) - 88.6

C3H8 (propane) - 42.1

C4H10 (butane) - 0.5

Boiling Points of Hydrocarbons

CH4

C2H6

C3H8

C4H10

Note linear relation between bp and molar mass.

LiquidsLiquids In a liquid• Molecules are in

constant motion• There are

appreciable intermolecular forces

• Molecules are close together

• Liquids are almost incompressible

• Liquids do not fill the container

Liquids: Energy and Phase changes

Melting and Freezing

• Melting: add energy to break bonds that keep molecules at fixed position (kinetic energy goes up): endothermic

• Freezing: energy released (kinetic energy goes down) as particles “settle” in fixed positions: exothermic

LIQUID VAPOR

Evaporation: Add energy to break IM bonds

Condensation: Remove energy to form IM bonds

Evaporation and condensation

To evaporate, molecules must have sufficient energy to break IM forces. This breaking requires energy, so the process of evaporation is endothermic.

Evaporation

Condensation

• When a gas or vapor condensates, the kinetic energy of molecules gets lower, while IM forces get stronger. Energy is released: this process is exothermic.

Distribution of Energy in a Liquid

.

0

Nu

mb

er o

f m

olec

ule

s

Molecular energy

minimum energy neededto break IM forces and evaporate

higher Tlower T

At higher T a much larger number of molecules has high enough energy to break IM forces and move from liquid to vapor state.

High E molecules carry away E. You cool down when sweating or after swimming.

When molecules of liquid are in the vapor state, they exert a VAPOR PRESSURE.

EQUILIBRIUM VAPOR PRESSURE is the pressure exerted by a vapor over a liquid in a closed container when the rate of evaporation = the rate of condensation.

Liquid in flask evaporates and exerts pressure on manometer.

Measuring Equilibrium Vapor Pressure

Measuring Equilibrium Vapor Pressure

HEAT OF VAPORIZATION

HEAT OF VAPORIZATION is the heat required (at constant P) to vaporize the liquid.

Compd. ∆vapH (kJ/mol) IM Force

H2O 40.7 (100 oC) H-bonds

SO2 26.8 (-47 oC) dipole

Xe 12.6 (-107 oC) induced dipole

Heat of vaporization

• How much heat is needed to vaporize 6.51 mL of water? (ΔvapH = 40.7 kJ/mol)

• Density of water = 1.0 g/mL, so 6.51 mL = 6.51 g• 6.51 g/18.02 g/mol = 0.361 mol H2O• 0.361 x 40.7 = 14.7 kJ

Boiling Liquids

Liquid boils when its vapor pressure equals atmospheric pressure.

Liquid boils when its vapor pressure equals atmospheric pressure.

Then bubbles of vapor form within the liquid.

Then bubbles of vapor form within the liquid.

Boiling point

Boiling occurs when a liquid turns to a gas inside the liquid◦bubbles are produced

Liquid boils when its Vapor Pressure = Atmospheric Pressure◦Normal boiling point

Larger IMF = lower vapor pressure = high BPWeaker IMF = high vapor pressure = lower BP

Molecules at surface behave differently than those in the interior.

Molecules at surface experience net INWARD force of attraction. This leads to SURFACE TENSION — the energy required to break the surface.

Surface Tension

Surface Tension

Surface Tension also leads to spherical liquid droplets

Surfactants• Surface tension can be decreased by adding

surfactants (soap, detergents).• They interfere with hydrogen bonding.

Capillary action

IMF also lead to capillary action and to the existence of a meniscus for a water column.

This is caused by ADHESIVE

FORCES between water and glass

Capillary Action

Movement of water up a piece of paper depends on H-bonds between H2O and the OH groups of the cellulose in the paper.

Viscosity

• Resistance to flow – • Goes up as IMF increases• Also goes up as length of

hydrocarbon chain increases as molecules get tangled up and don’t flow easily

Phase diagram

• We use these diagrams to relate the process that occur when a substance changes from one phase to another.

• Substances are in the following states when in certain locations on the diagram:– Solid – left side of diagram– Liquid – middle of diagram– Gas/Vapor – right side of diagram

• When either the temp or pressure is changed, you can identify the process that is taking place and identify the phase change.– Ex (from diagram on last slide) – At 1 atm if you

increase the temperature from 90oC to 200oC, the process you are undergoing is vaporization or boiling (liquid to gas).

Triple point• The change of state occurs right on the

equilibrium line.• Triple point identifies the conditions

when you have all 3 states in dynamic equilibrium with one another.

Normal MP and BP

• Tm normal melting point–The point at 1 atm or 101.3 kPa

when solid turns to liquid.• Tb normal boiling point

–The point at 1 atm or 101.3 kPa when a liquid turns to a vapor

Critical point

• Critical point – you are no longer able to distinguish between gas and liquid phases past this point.

CO2 Phase diagram