Intermolecular Forces

- Average kinetic energy of particles
- Recall KE ∝ Tave
- Intermolecular Forces
- Energy of Inter-particle attraction

- How tightly molecules are packed together
- Strength of attractions between molecules

- example of intermolecular attraction
- Molecules must get closer together
- Cool or compress

- example of intermolecular attraction
- Requires molecules to move farther apart
- Heat or reduce pressure

At certain T, molecules don’t have enough energy to break away from one another’s attraction

- Covalent bonds within molecule
- Strong
ΔHbond (HCl) = 431 kJ/mol

- Attraction forces between molecules
- Weak
- ΔHvaporization (HCl) = 16 kJ/mol

- Measure of attractive force that one atom in a covalent bond has for electrons of the bond
**Two atoms with different electronegativity values share electrons unequally
**Electron density is uneven
Higher charge concentration around more electronegative atom
- Electronegatity is related to bond dipoles

- Indicated with delta (δ) notation
- Indicates partial charge has arisen

- Symmetrical molecules
- Asymmetrical molecules
- Molecular dipoles

- Even if they have polar bonds
- Are non-polar because bond dipoles cancel

- Are polar because bond dipoles do not cancel
- These molecules have permanent, net dipoles

- Cause molecules to interact
- Decreased distance between molecules increases amount of interaction

- When substance melts or boils
**Intermolecular forces are broken
**Not covalent bonds
- Responsible for non-ideal behavior of gases
- Responsible for existence of condensed states of matter
- Responsible for bulk properties of matter
**Boiling points and melting points
**Reflect strength of intermolecular forces

1. Dipole-dipole forces
- Hydrogen bonds
2. London dispersion forces
3. Ion-dipole forces
- Ion-induced dipole forces

- Occur only between polar molecules
** Possess dipole moments
- Molecules need to be close together
- Polar molecules tend to align their partial charges
** Positive to negative
- As dipole moment increases, intermolecular force increases
- also, contains London Dispersion Reactions

measure of seperation of two electrical charges
- has a magnitude and direction - polar molecules tend to align
their partial charges -- from pos to neg.
** OPS dipoles cancel!!

- Interactions between net dipoles in polar molecules
- About 1–4% as strong as a covalent bond
**Decrease as molecular distance increases
- Dipole-dipole forces increase with increasing polarity

- Special type of dipole-dipole Interaction
** Very strong dipole-dipole attraction
** ~10% of a covalent bond
- Occurs between H and highly electronegative atom (O, N, or F)
H—F, H—O, and H—N bonds very polar
** Electrons are drawn away from H, so high partial charges
** H only has one electron, so δ+H presents almost bare proton
** δ–X almost full –1 charge
** Element’s small size, means high charge density
** Positive end of one can get very close to negative end of another

- Responsible for expansion of water as it freezes
- Hydrogen bonding produces strong attractions in liquid
- Hydrogen bonding (dotted lines) between
water molecules in ice form tetrahedral configuration (triangular)

- When atoms near one another, their valence electrons interact
** example: electrons always moving to the left
- Repulsion causes electron clouds in each to distort and polarize
- Instantaneous dipoles result from this distortion
** Effect enhanced with increased volume of electron cloud size
** Effect diminished by increased distance between particles and compact arrangement of atoms

** London Forces
Dispersion forces
Operate between all molecules
Neutral or net charged
Nonpolar or polar

- Ease with which dipole moments can be induced and thus London Forces depend on
1. Polarizability of electron cloud
2. Points of attraction
** Number atoms
** Molecular shape (compact or elongated)

- Ease with which the electron 
cloud can be distorted
- Larger molecules often more polarizable
** Larger number of less tightly held electrons
** Magnitude of resulting partial charge is larger
** Larger electron cloud
** Larger molecules have stronger London forces and thus higher boiling points.

- London forces depend on number atoms in molecule
- Boiling point of hydrocarbons demonstrates this trend (larger atoms = larger bp)

1. Molecular Shape
2. Physical Origin of Shape Effect

- Increased surface area available for contact = increased London forces
- London dispersion forces between spherical molecules are lower than chain-like molecules
** More compact molecules
** Hydrogen atoms not as free to interact with hydrogen atoms on other molecules
- Less compact molecules
** Hydrogen atoms have more chance to interact with hydrogen atoms on other molecules

- small area for interaction = more compact - lower BP
- larger area for interation = less compact - higher BP

- Attractions between ion and charged end of polar molecules
- Attractions can be quite strong as ions have full charges

** AlCl3·6H2O
- Positive charge of Al3+ ion attracts partial negative charges δ– on O of water molecules
- Ion-dipole attractions hold water molecules to metal ion in hydrate
** Water molecules are found at vertices of octahedron around aluminum ion

- Often can predict physical properties (like BP, MP and many others) by comparing strengths of intermolecular attractions
- Larger, longer, and therefore heavier molecules often have stronger intermolecular forces
- Smaller, more compact, lighter molecules have generally weaker intermolecular forces

1. Compressibility
2. Diffusion

- Measure of ability of substance to be forced into smaller volume
- Determined by strength of intermolecular forces
- Gases highly compressible
** Molecules far apart
** Weak intermolecular forces
- Solids and liquids nearly incompressible
** Molecules very close together
** Stronger intermolecular forces

- Movement that spreads one gas though another gas to occupy space uniformly
- Spontaneous intermingling of molecules of one gas with molecules of another gas
** Occurs more rapidly in gases than in liquids
** Hardly at all in solids

In Gases:
- Molecules travel long distances between collisions
- Diffusion rapid
In Liquids:
- Molecules closer
- Encounter more collisions
- Takes a long time to move from place to place
In Solids:
- Diffusion close to zero at room temperature
- Will increase at high temperature

- Retention of volume and shape
** Solids retain both volume and shape
>>Strongest intermolecular attractions
>>Molecules closest
** Liquids retain volume, but not shape
>>Attractions intermediate
** Gases, expand to fill their containers
>>Weakest intermolecular attractions
>>Molecules farthest apart

- Decrease with increasing temperature
** Increasing kinetic energy overcomes attractive forces
** If allowed to expand, increasing temperature increases distance between gas particles and decreases attractive forces

- Causes a liquid to take the shape (a sphere) that minimizes its surface area
- Molecules at surface have higher potential energy than those in bulk of liquid and move to reduce the potential energy
- Liquids containing molecules with strong intermolecular forces have high surface tension
- Allows us to fill glass above rim: Gives surface rounded appearance, Surface acts as “skin” that lets water pile up,
Surface resists expansion and pushes back
- Surface tension increases as intermolecular forces increase
- Surface tension decreases as temperature increases

- Ability of liquid to spread across surface to form thin film
- Greater similarity in attractive forces between liquid and surface, yields greater wetting effect
- Occurs only if intermolecular attractive force between surface and liquid about as strong as within liquid itself

Ex. H2O wets clean glass surface as it forms
H–bonds to SiO2 surface
** Does not wet greasy glass, because grease is nonpolar and water is very polar
- Only London forces
- Forms beads instead
** Surfactants

- Substances that have both polar and non-polar characteristics
- Long chain hydrocarbons with polar tail
- Nonpolar end dissolves in nonpolar grease
- Polar end dissolves in polar H2O
- Thus increasing solubility of grease in water

- Resistance to flow
- Measure of fluid’s resistance to flow or changing form
- Related to intermolecular attractive forces
- Also called internal friction
** Depends on intermolecular attractions
example: dishwasher soap more viscosity (thicker) versus water with a low viscosity —- moves in a slow motion resisting flow

- Viscosity decreases when temperature increases
- Most people associate liquids with viscosity
** Syrup more viscous than water
- Gases have viscosity
** Respond almost instantly to form-changing forces
- Solids, such as rocks and glass have viscosity
** Normally respond very slowly to forces acting to change their shape

- “Like dissolves like”
**To dissolve polar substance, use polar solvent
**To dissolve nonpolar substance, use nonpolar solvent (** applies to ions)
- Compare relative polarity
**Similar polarity means greater ability to dissolve in each other
**Differing polarity means that they don’t dissolve, they are insoluble ** water and oil
- Surfactants ** like soup and detergent
** Both polar and non-polar characteristics
** Used to increase solubility

- Substances exist in one of three phases: solid, liquid or gas.
- Substances can change phase by the absorption or release of energy.

- a phase change directly from solid to gas w/o ever entering the liquid phase
** gas at -78 C it will become dry ice allowing it to absorb heat
from its surroundings to sublime going from a SOLID TO A GAS
It is a physical change !!

Depends on:
** Temperature
** Surface area
** Strength of intermolecular attractions
- Molecules that escape from liquid have larger than minimum escape KE
** When they leave
** Average KE of remaining molecules is less and so T lower

- For given liquid
** Rate of evaporation per unit surface area increases as T increases
Why?
- At higher T, total fraction of molecules with KE large enough to escape is larger
- Result: rate of evaporation is larger

- Liquid is placed in empty, closed, container
** Begins to evaporate
- Once in gas phase
** Molecules can condense by
** Striking surface of liquid and giving up some kinetic energy

Rate of evaporation = rate of condensation
** Occurs in closed systems where molecules cannot escape

- Solid begins to change into liquid as heat added
- Dynamic equilibria exists between solid and liquid states
This is true if you are making ice cubes in a freezer that is
EXACTLY 0 degrees celsius and so is the surroundings ! = PHASE EQUILIBRIUM

example: Temp S - L: M.P (add heat to melt (vibrating), but temp stays the same - because this heat is breaking the intermolecular forces - now the molecules will flow at the SAME kinetic energy; Meaning, average kinetic energy is CONSTANT - During a phase change a kinetic energy is being changed to potential energy.

Heat removed at a constant rate
- Supercooling
** Temperature of liquid dips below its freezing point
- Remove heat from a liquid = temp will decrease

below freezing point, but still liquid - did not have time to orient themselves to crystallize - However, this does not last so long

Must add heat
Energy entering system (+)
Sublimation: ΔHsub > 0
Vaporization: ΔHvap > 0
Melting or Fusion: ΔHfus > 0
** condensation = - (same amount of jules of vaporization just the negative) = same value

- Variation of vapor pressure with T
Ether:
** Volatile
** High vapor pressure near RT
Propylene glycol:
** Non-volatile
** Low vapor pressure near RT

As temp increase - vaporization increase = This is because there is more collisions with higher pressure

depends on temp (not surface area or mass) —- Meaning, it does not matter if you have 1 gram of water or 100 grams of water, just must be PURE water.

The temp. at which the vapor pressure of a liquid equals normal atm P (1 atm).

The temp. at which a liquids
v.p. equals the external pressure.
Ex. At high altitudes, P is lower. Thus, water will boil at a lower temp.

Initial V:
** Liquid – vapor equilibrium exists
Increase V:
** Pressure decreases
** Rate of condensation decreases
More liquid evaporates:
** New equilibrium established

Yes
At given temperature
Some solid particles have enough KE to escape into vapor phase
When vapor particles collide with surface
They can be captured
Equilibrium vapor pressure of solid
Pressure of vapor in equilibrium with solid

- gradually get higher when you go down a group, more and more electrons
DOWN GROUP: central atom gets larger and larger more protons neutrons and electrons
non polar molecules
- INT attractions increase, but not by much
MAJOR EXCEPTION: extra strong intermolecular force - hydrogen bonding (all the way to the left on the graph — more then just dispersion forces, more then just dipole-dipole)

Heat absorbed by one mole of solid when it melts to give liquid at constantT and P

Heat absorbed when one mole of liquid is changed to one mole of vapor at constant T and P
- Molar heat of vap: code for boiling point; if you boil water(100C) - if you countiue to add heat it stays at the same heat and is used to break the forces (a medium to cook in - cannot burn anything in boiling water)

Heat absorbed by one mole of solid when it sublimes to give one mole of vapor at constant T and P
- stronger intermolecular forces needs more energy

All of these quantities tend to increase with increasing intermolecular forces

positive and negative = SAME AMOUNT just negative or positive (20 dollars out of you bank account) - 20 absorbed 20 released same amount of J

- Show the effects of both pressure and temperature on phase changes
Boundaries between phases indicate equilibrium
** Triple & Critical Points

The temperature and pressure at which s, l, and g are all at equilibrium.
- one temp and pressure for each substance

The temperature and pressure at which a gas can no longer be condensed
TC = temperature at critical point
PC = pressure at critical point