IMF

fundamental difference between states of matter

distance b/w partciples

gas

total disorder

much empty space

liquid

disorder

partciles/clusters free to move relative to eachother

crystaline solid 

ordered arrangement 

partciles are essentially in fixed positions 

particles close together 

condense phases

liquids and solids

particles are close together

state that a substance is in at a particular temperature and pressure depends on

kinetic energy of the particles

strength of attraction b/w the particles

WHY do KE and IMFS affect phase 

IMFS stick molecules together 

Kinetic Energy Pulls them apart 

Solids -IMFS win 

Liquid - Tie 

Gas- KE wins  

IMFS

attractions b/w molecules are not nearly as strong as intramolecular attractions

BUT can control physical properties

• melting+boiling point

• vapor pressures

• viscosities

Dipole

Hydrogen Bonding

London Dispersion Forces

Ion Dipole

important in solutions of ions

make it possible for ionic substances to dissolve in polar solvents

Dipole 

molecules that have permanent dipoles attracted to eachother

• positive end of one is attracted to the end of the other and vice-versa 

• these forces are only important when the molecules are close to each other

the more polar the molecule

the higher is its boling point

LDF

attractions between an instantaneous dipole and an induced dipole.

These forces are present in all molecules whether they are polar or nonpolar.

Instantaneous Dipole:
A temporary dipole in a molecule caused by random electron movement.

Induced Dipole:
A dipole created in a molecule by the presence of a nearby instantaneous dipole.

polarizability 

refers to the tendency of molecules to generate induced electric dipole moments when subjected to an electric field

Factors affecting LDFS

long, skinny molecules - stronger LDFS

short, fat molecules - weaker LDFS

relationships b/w dispersion forces and increase # of electrons

LDFS increase w/ increased number of electrons because larger atoms have larger electron clouds which are easier to polarize

which have a greater effect - dipole or dispersion 

• if 2 molecules are comparable shape/size dipole is dominating 

• if 1 molecules is larger than other - dispersion forces determine physical properties 

Hydrogen Bonding

strong type of intermolecular force that occurs when a hydrogen atom bonded to a highly electronegative atom

N,O,F bonded DIRECTLY to H

hydrogen nucleus is highly exposed

Hydrogen Bonding cont.

Molecules can be H-bond donors or acceptors or both

A lone pair on FON can accept H-Bond

the hydrogen that is a donor has to be bonded to FON

Summarizing IMFS

Viscosity

resistance of a liquid to flow

increases with stronger IMFS and decreases with higher temperature

Surface Tension 

net inward force experience by the molecules on the surface of a liquid 

resistance of a liquid to an increase in its surface area 

large IMFS= high surface tension

polar molecules = high surface tension 

Chromatography

is used to separate mixtures of substances into their components.

Phases

• mobile

• stationary

How do components separate?

Different rates due to varying affinities for mobile/stationary phases.

Stationary and mobile phases in paper chromatography?

Stationary = paper, Mobile = solvent (water, ethanol, hexane)

Affinity : mobile vs stationary

High mobile phase affinity → ?

Moves far on chromatogram.

High stationary phase affinity → ?

Moves slowly; stays near start.

Column chromatography principle?

Components move through column at different speeds depending on affinity for stationary phase.

Phases are determine by three things

Kinetic Energy

Pressure

Strength of IMFs

Heat of Fusion

Energy needed to change a solid at its melting point to a liquid

Heat of Vaporization 

Energy needed to change a liquid at is boiling point to a gas 

Flat/horizontal points on phase change graphs

The substance is changing phase (e.g., melting or boiling) at constant temperature.

Heat energy is used to break/form intermolecular bonds instead of increasing kinetic energy.

• Melting (fusion): solid → liquid

• Boiling (vaporization): liquid → gas

Endothermic Phase Changes

• Melting (Fusion): solid → liquid

• Vaporization (Boiling/Evaporation): liquid → gas

• Sublimation: solid → gas

Exothermic Phase Changes

• Freezing (Solidification): liquid → solid

• Condensation: gas → liquid

• Deposition: gas → solid

Vapor Pressure

Vapor pressure is the pressure exerted by the vapor (gas) particles of a liquid when the liquid and its vapor are in dynamic equilibrium in a closed container.

Temperature and Vapor PRessure

As temperature increases, more molecules have enough energy to escape, so vapor pressure increases

when a liquid and its vapor reach dynamic equilibrium

 The rate of evaporation equals the rate of condensation

When does a liquid boil

When its vapor pressure equals the external (atmospheric) pressure

What is the normal boiling point

The temperature at which vapor pressure = 760 torr (1 atm).

Crystalline Solids

Solids with particles arranged in a highly ordered, repeating pattern.

amorphous solids

Solids with no specific or orderly particle arrangement

covalent-network solids

Solids where atoms are bonded by covalent bonds in a continuous network

Very hard, high melting points.

Diamond, SiO₂ (silicon dioxide), SiC (silicon carbide).

allotrope

Different structural forms of the same element (e.g., diamond and graphite are allotropes of carbon).

molecular solids

Solids where molecules are held together by intermolecular forces.

Softer, with lower melting points compared to network solids.

Ice (solid H₂O).

metallic solid

metal cations surrounded by delocalized valence electrons.

What type of bonding holds metallic solids together?

Metallic bonding — attraction between metal cations and delocalized electrons.

metals conduct electricity- Because delocalized electrons move freely throughout the solid.

metals :malleable, ductile,volatile

are malleable and ductile because : Metal ions can slide past each other without breaking the metallic bond

low volatility- Strong metallic bonding keeps atoms tightly bound, making them hard to vaporize.

alloy

A mixture of a metal with another element (metal or nonmetal)

interstitial alloy

Formed when smaller atoms fill spaces between larger metal atoms

More rigid; decreased malleability and ductility.

substitutional alloy

Formed when atoms of similar size substitute for each other in the lattice

Density between component metals; slightly less malleable and less ductile than pure metals

Do alloys conduct electricity

Yes, most alloys remain good conductors.

What is doping?

Adding a small amount of another substance to improve properties (e.g., conductivity).

: What is n-type doping

Replacing an atom with one that has more electrons, adding extra free electrons for conduction

What is p-type doping?

Replacing an atom with one that has fewer electrons, creating “positive holes” that allow electrons to move and conduct.

How does doping affect conductivity?

It increases conductivity by adding free electrons (n-type) or holes (p-type).

Solutions

homogeneous mixtures of two or more pure substances.

In a solution, the solute is dispersed uniformly

throughout the solvent.

How Does a Solution Form?

Solvent molecules attracted to surface ions.

Each ion is surrounded by solvent molecules.

Enthalpy (ΔH) changes with each interaction broken or

formed.

Solvated

a solute, or dissolved substance, has been surrounded by a solvent
if the solvent is water the ions are hydrated

interparticle force in solutions 

ion-dipole.

Enthalpy Changes in Solutions

1) Separation of solute

particles.

2) Separation of solvent

particles to make

‘holes’.

3) Formation of new

interactions between

solute and solvent.

Entropy

Dispersal of energy in

the system.

Number of microstates

(arrangements) in the

system (i.e. disorder)

more spread out - higher

less spread out - lower

Dissolution 

usually considered a physical change —you

can get back the original solute by evaporating the

solvent.

If you can’t, the substance didn’t dissolve, it reacted.

Saturated Solution

• Solvent holds as much

solute as is possible at

that temperature.

• Undissolved solid

remains in flask.

• Dissolved solute is in

dynamic equilibrium

with solid solute

particles.

Unsaturated Solution

Less than the

maximum amount of

solute for that

temperature is

dissolved in the

solvent.

No solid remains in

flask.

Supersaturated 

Solvent holds more solute than is normally

possible at that temperature.

These solutions are unstable; crystallization can

often be stimulated by adding a “seed crystal” or

scratching the side of the flask.

Factors Affecting Solubility

Polar substances tend to

dissolve in polar solvents.

Nonpolar substances tend

to dissolve in nonpolar

solvents.

The stronger the

intermolecular

attractions between

solute and solvent,

the more likely the

solute will dissolve.

Gases in solutions 

solubility of gases in water increases with increasing mass.

solubility of liquids and solids with pressure

does not change with pressure

solubility of a gas in a liquid

directly proportional to its pressure.

How does temperature affect the solubility of most solid solutes in liquids

Solubility increases with increasing temperature.

: How does temperature affect gas solubility in liquids?

Solubility decreases as temperature increases.

Why do gases become less soluble at higher temperatures?

: Higher temperature gives gas molecules more kinetic energy, allowing them to escape the solution.

Mass Percentage

Mass % of A =

mass of A in solution/

total mass of solution

× 100%

Volume Percentage

Volume% of A =

(vol of A in solution/

total vol of solution)

× 100%

Mole Fraction

Moles of solute/moles of solute+solvent

Colloids 

Suspensions of particles larger than

individual ions or molecules, but too small to

be settled out by gravity.

Tyndall Effect

Colloidal suspensions

can scatter rays of light.

Colloids in Biology

Some molecules have

a polar, hydrophilic

(water-loving) end and

a nonpolar,

hydrophobic

(water-hating) end.

Intermolecular Forces are NOT 

bonds. Avoid

using the word bonding when referring to IMFs.

(i.e. don’t say something like “both molecules are

polar so they have strong bonds”)

IMFS are not found in

ionic compounds

what not to use as an explanation for increasing LDFS

increased mass as an explanation for

increasing LDFs. Instead use “more polarizable”

(or more electrons, so it is more polarizable)

single molecule does NOT have 

have hydrogen

bonds. H-Bonds are between 2 (or more)

molecules. (Don’t say things like “A water molecule

has hydrogen bonding” or “a water molecule has

hydrogen bonds between the O and H”)

if molecules are of similar size( number of electrons) IMFS rank

LDF < Dipole-dipole < H-bonding

when dissolve in water

Molecules do not ionize or break apart

sugar molecule is still a

sugar molecule when dissolved in water. Ionic

compounds ionize (or dissociate) into their

component ions.

Volatile refers to 

vapor pressure

High volatility = high vapor pressure

Low volatility = low vapor pressure

Which substance is more volatile = which

substance has high vapor pressure