Intermolecular forces: Generalizing properties

Low boiling point = particles are more likely to leave liquid solution

Weaker IM forces = lower boiling point Lower boiling point = more vapor = higher vapor

pressure High boiling point = slow evaporation If IM forces are the same, look at formula weight.

Heavier molecules have higher boiling points. Strength of IM forces: Hydrogen bond>dipole-dipole>London dispersion

Intermolecular ForcesInteracting molecules

or ions

Polar molecules? Ions involved? Are polar molecules and ions both present?

Are H atoms bonded toN,O or F atoms?

No Yes

Yes

London ForcesonlyEx. Ar(l), I2(s)

No

Dipole-Dipole

Ex. H2S

No

Hydrogen BondingEx. NH3, H2O

Yes

Ion-dipole ForcesEx. KBr in H2O

Yes

Ionic bondingEx. NaCl

No

Water’s Properties

Hexagonal crystal shape Molecule is polar. Hydrogen bonding

Ice floats. Expands during freezing until -4.0 º C. Solid form is less dense than liquid Surface tension:

Water “beads” on smooth surfaces. Insects walk on water surfaces.

Surface tension Force that pulls adjacent parts of a liquid

surface together. The higher the attractive forces between

particles in the liquid, the higher the surface tension.

Hydrogen bonds make water have higher surface tension than most liquids.

Water droplet

Soap

Phases of matter: ComparisonProperty Solid Liquid Gas

Particles Closely packed

High density(But…Water is different!)

More densely packed than in gas

Most compressible-least densely packed

Particle movement

Vibrate weakly around fixed positions

Lowest kinetic energy

Can change positions with other particles

Can change positions with other particles

Highest kinetic energy

Intermolecular forces

Most effective (strongest)

Stronger than in gases

Least effective (weakest)

Shape and volume

Both definite Definite volume only

No definite shape or volume

Solids

Crystalline solids Particles are arranged in an orderly, geometric, repeating

pattern. Examples: Emerald, diamond, calcite

Amorphous solids: (Without shape) Particles are arranged randomly. Examples: Glass, plastic

Network solids Covalent bonds, usually single element arranged in orderly

pattern Examples: Diamond, graphite

Bonding in Solids Molecular solids

Most are liquids or gases at room temp. Ex. H2O, Ar

Covalent Network solidsCovalent Network solids Covalent bonds are stronger than IM forces, so substances

have relatively high melting points and are harder than molecular ones.

Ex: quartz, diamond, graphite, SiO2

Ionic solids Ionic bonds are the strongest of all Strength of bond depends on charge: Higher charges =

higher melting point. Crystal structures: Examples

Face-centered cubic, body-centered cubic, hexagonal close-packed structures.

Metallic Solids (metallic bonds)

8 Mullis

The Crystal Lattice

3-dimensional pattern that repeats itself over and over again.

Each ion is bonded with all oppositely charged ions that directly surround it.

NaCl forms a cube shape, called a body-centered-cubic structure.

There are 7 crystal shapes, determined by how the ions are arranged in the lattice.

9 Mullis

Crystal Growth

Crystals grow by adding ions to all sides. They grow equally in all directions from the

outside. Crystals form in 2 ways:

Solution containing a dissolved ionic compound evaporates.

An ionic solid is heated until it melts, then liquid is cooled. (Igneous rocks)

10 Mullis

Energetics of Ionic Bond Formation

Recall that heat of formation of NaCl was exothermic (∆H°f = -410.9kJ/mol)

Separation of NaCl is endothermic: (∆H = +788 kJ/mol) The energy required to separate 1 mol of ions

in an ionic lattice into gaseous ions is called lattice energy, ∆Hlattice .

Lattice energy depends on the charge on the ions and the size of the ions.

11 Mullis

Lattice Energy (Ionic bonds) Lattice energy depends on the charge on the ions

and the size of the ions. The stability of the compound comes from the

attraction between ions of unlike charge. The specific relationship is given by Coulomb’s

equation: E = kQ1Q2 Q is the charge on the particles, d is the distance

d between their centers and k is a constant.

As Q1 and Q2 increase, E increases and as d increases, E decreases.

Lattice Energies for Some Ionic Compounds

Compound

Lattice Energy (kJ/mol)

Compound Lattice Energy (kJ/mol)

LiF 1030 MgCl2 2326

LiCl 834 SrCl2 2127

LiI 730

NaF 910 MgO 3795

NaCl 788 CaO 3414

NaBr 732 SrO 3217

NaI 682

KF 808 ScN 7547

KCl 701

KBr 671

CsCl 657

CsI 600

CrystalsImages created by Daniel Mayer or Wikimedia Commons and licensed under terms of the GNU FDL.

Can examine structure using X-ray diffraction Uses Bragg’s Law to determine

distance between planes of atoms. Computer instrumentation is used to translate wave functions into photographic images.

Bragg’s Law: nBragg’s Law: nλλ = 2dsin = 2dsinΘΘ n is an integer (1), λ is wavelength, d is

distance between atoms, Θ angle of incidence

3D link:http://www.le.ac.uk/eg/spg3/atomic.html

cbc

cfc

hexagonal

Allotropes (different forms of same element)

Carbon (C) Diamond Graphite (pencil “lead”) Charcoal

Sulfur (S) Rhombic (puckered ring)= S8

Phosphorous (P) White phosphorous, P4 is most reactive,

tetrahedral Red phosphorous is more stable.

(1) http://www.green-planet-solar-energy.com/the-element-sulfur.html(2) http://www.avogadro.co.uk/structure/chemstruc/network/g-molecular.htm(3) http://www.enmu.edu/services/museums/miles-mineral/images/diamond_large.jpg

(1) Sulfur rings

(2) graphite’s hexagonal crystal structure (3) Diamond

Silicon Doping(N-type is more conductive when voltage is applied.)

O:O:O:O: O:O:O:O: O:O:O:O:

O:O:O:O: O:B.O:O: O:P:O:O:

Silicon (4 e-) P-type N-type

semiconductor hole created extra e- in lattice

p = positive n= negative

To customize

conductive properties, add a dopant such as B (p-type), As or P (n-type)

.. .. .. .. .. .. .. .. .. … .. ..

.. .. .. .. .. .. .. .. .. .. .. ..

Glass: SiO2 High melting point (~ 1700°C) but may vary depending on the particular

structure . Very strong Si-O covalent bonds arranged in a continuous lattice have to be broken throughout the structure before melting occurs.

Is also very hard due to the need to break the very strong covalent bonds. Doesn't conduct electricity because there aren't any delocalized electrons. All

the electrons are held tightly between the atoms, and aren't free to move. Insoluble in water and organic solvents because there are no possible

attractions which could occur between solvent molecules and the silicon or oxygen atoms which could overcome the covalent bonds in the giant structure.

(Water is a simple covalent structure: Each molecule is made of 3 atoms so its melting point is low compared to the giant lattice structure of solids like SiO2. )

Allotrope: Two or more forms of the same element that have distinctly different physical or chemical properties.

Fullerenes include C60, buckminsterfullerene, a hollow sphere resembling a soccer ball.

Graphite is a black solid that feels soft and “greasy” to the touch. Planar sheets of molecules can slip by one another easily. It is used as a lubricant and leaves black marks if rubbed on a lighter-colored surface. It conducts electricity. Selling price: < $0.01/gram.

Diamond is one of the hardest substances known (Mohs hardness = 10). Its hardness is due to rigid networks of tetrahedrons, carbon atoms covalently bound. It does not conduct electricity. Selling price = $50.00 - $20,000.00/gram.

Material modification Pencil “lead” is softened by adding clay to graphite. Gold jewelry is strengthened by adding copper or other

metal. 14 karat gold means that 14/24’s of the material is Au. (The relative proportion of gold originated with a medieval coin called a mark; a mark weighed 24 karats.)

Ceramics: Developed from conventional clay (Si, O, Al) and the addition of other minerals to improve strength, melting point and brittleness. Ceramics can often get much hotter before they melt than metals.

Plastics: Synthetic polymers primarily from carbon. Disadvantage is that most are made from nonrenewable petroleum resources.

Changing states

Equilibrium: When there is no net change in a system.

Dynamic equilibrium: When a vapor is in equilibrium with its liquid

as one molecule leaves the liquid to become a vapor, another molecule leaves the vapor to become a liquid. In other words, an equal number of molecules will be found moving in both directions.

Boiling Point

Vapor pressure: Pressure exerted by a vapor = Pressure of the liquid at given temperature

Liquid boils when its vapor pressure equals pressure of the atmosphere.

Boiling is the conversion of a liquid to vapor within the liquid as well as at its surface.

Boiling point is the temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure.

Volatile liquids are liquids that evaporate readily.

Boiling Point, cont.

High elevation: Low atmospheric pressure Low atmospheric pressure = lower boiling

point High pressure in pressure cooker = increased

boiling point, faster cooking If pressure above liquid increases, the liquid

temperature rises until it matches the new pressure and boils again.

Separation by Distillation

Distillation is the separation of liquid substances according to their different boiling points.

As a liquid mixture is heated, the substance with the lower boiling point will vaporize first.

Distillate: Condensed liquid substance

Kinetic Energy and Equilibrium Vapor Pressure

In the beginning: # particles condensing to liquid phase =

# particles evaporating to gas phase Increase temp Increase kinetic energy Now, more molecules have enough energy to

leave the liquid. More vapor molecules = higher vapor pressure Equilibrium will soon be established, but at a

higher vapor pressure.

Heat of Vaporization

Amount of heat necessary to boil (or condense) 1.00 mole of a substance at its boiling point

***1) 1.00 mole of a substance***2) There is no temperature change

The molar heat of vaporization (ΔHvap ) for water is 40.7 kJ/mol. It comes from a table.

q = ΔHvap (mass / molar mass) (q = ΔHvapn) q is total amount of heat involved.

Heat of Vaporization

Vapor pressure increases nonlinearly for liquids. Mathematically, the relationship is

ln(Pvap) = -ΔHvap (1/T) +C

RWhere C = constant characteristic of a given liquid.

m = slope = -ΔHvap and x = 1/T and b = intercept = C

R

Heat of Fusion aka standard enthalpy change of fusion Amount of thermal energy absorbed or lost for 1 gram of a

substance to change states from a solid to a liquid or vice versa. Temperature at which it occurs is called the melting point. Temperature falls if thermal energy is removed from a liquid or

solid At the transition point between solid and liquid (melting point),

EXTRA energy is required to go from liquid to solid and increase order. For molecules to maintain the order of a solid, extra heat must be withdrawn.

In the other direction, to create the disorder from the solid crystal to liquid, extra heat must be added.

The molar heat of fusion for water is 6.02 kJ/mol. q = ΔHvap (mass / molar mass)

Phase DiagramA phase diagram is a graph of pressure vs. temperature that shows the conditions under which phases of matter exist.

Critical temp (Tc): Above this, the substance cannot exist in the liquid state.

Phase Diagrams: Density

Negative liquid/solid slope shows density of solid is LESS than

liquid (like H2O). See previous slide. Most substances will have a positive slope of this line since most solids are more

dense than the liquid:

http://wine1.sb.fsu.edu/chm1045/notes/Forces/Phase/Forces06.htm

Four major "points" on a phase diagram

1. Triple point, TP - All three phases can exist in equilibrium at this temperature and pressure. (The solid-liquid line and the liquid-vapor line meet.)

2. Normal boiling point, Tb - The temperature at which the vapor pressure of a liquid is equal to standard atmospheric pressure. (Standard atmospheric pressure line crosses the liquid-vapor line.)

3. Normal melting point, Tm - The temperature at which the vapor pressure of the solid and the vapor pressure of the liquid are equal.

(Standard atmospheric pressure line crosses the solid-liquid line.)

4. Critical temperature, Tc - The temperature above which no amount of pressure will liquefy a vapor.

(The liquid-vapor line becomes vertical.)