Chemistry IMF

phase

a form of matter that is uniform throughout in both chemical composition and physical state.

condensed phases

solids and liquids, because when a gas is cooled it condenses (becomes more compact) to a liquid or a solid since attractive forces pull them together, however repulsion dominates at short seperations

temperature at which it condenses, shows strength of forces

potential energy diagram

the potential energy of a pair of molecules changes with the distance between their centers. The dip in the curve shows the bond formation, even when atoms or moelcule’s do not form a bond, attractive forces lower energies of particles 

as they become closer, attractive forces dominate, however as seperation becomes small repulsions dominate and the potential energy increases

dissolving of an ionic solid

ionic solid in water dissolves as water molecules surround each ion form its partial charge, separating ions form one another

hydration

the attachment of water molecule to solute particles due to polar character of H2O ions or molecule

it is an ion-dipole reaction, but they have to be close together,

the size of ion and charge influence the extent of hydration

• stronger ion-dipole interaction, the closer the dipole can come

• small cations are hydrated better than large ones

• higher ion charge better hydration

-|z|μ/r²

the potential energy of interaction between the full charge of an ion and the two partial charges of a polar molecule 

z = charge number of ion 

mu μ= dipole moment of polar molecule 

if it is negative it shows the potential energy of the ion and solvent is lowered by their attractive interaction 

why does potential energy of interaction between a point charge and a dipole fall off more quickly with distance than the interaction of two point charges does

At large distances, the two partial charges are at almost the same distance from the ion, and the cancellation is nearly complete. to point charges 1/r and point charge and dipole1/r²

dipole dipole interaction

the interaction betweend ipoles, between tehir partial charges

the potential energy for startionary polar moelcyles in a solid = -μ²/r³ if idneitical 

non idential = μ₁μ₂=r³

greater plarity , stirnger the interation

reason for rapid falling of the dipole dipole interaction potential enegry

as the distance between molecules increases, the opposite partial charges on each molecule appear to merge and cancel, whereas in the interaction between a point charge and a dipole, only the partial charges on the dipole appear to merge.

perfectly free rotation

found in gasses, the attractions between opposite partial charges and the repulsion between like partial charges cancel thus no net interaction between the neibouding molecules 

but in reality attractive interaction between oppostive partial charges slightly outweigh the repulsive, thus a weak net attraction between rotating neighboruing polar molecules in gas phase

potential energy proportional to 1/r^6

evidence for attractive interaction in non polar moelcules

monatomic polar noble gases can be liquified and other non polar compounds are liquids

electron distribution

the electron clouds of atoms and molecules are not uniform, creating fleeting partial charges and instantaneous dipole moment

instantaneous dipole moment

a momentary dipolar seperation of charge, distors the electron cloud on neighbouring molecule inducing a temporary dipole moment on that molecule

London dispersion interaction

acts between all molecules and atoms an dis the only interaction between non polar molecules in a monatomic gas, 

strength depends on polaribility 

bigger molecule stronger force, heavier atom sstronger force, longer molecule stronger force

polarizability

the easy with which their electron clouds can be distorted

highly polarizable molecules are those in which the nuclear charges have little control over the surrounding electrons, large molecules

dipole induced interaction

a polar molecule interacts with a non polar molecule, able to induce a dipole in molecule which is a permanent dipole moment

trend in group 14

boiling point increases down because strength of london interaction increases

However, ammonia, water, and hydrogen fluoride all show anomalous behavior. Their exceptionally high boiling points suggest that there are unusually strong attractive forces between their molecules. - due to hydrgen bodning 

hydrogen bonding

an intermolecular interaction in which a hydrogen atom bonded to a small, strongly electronegative atom, specifically N, O, or F, is attracted to a lone pair of electrons on another N, O, or F atom

dimers

pairs od identical molecules, they are linked by two hydrogen bonds

reoulsion

molecules close together repel one another, responsible for the steep rise in potential energy

when they get close together their orbitals overlap and form bonding and anti-bonding molecular orbitals. the antibonding orbital raises energy ore than  bonding orbital lowers energy, thus a net increase in energy as they get closer

rule of thumb for increasing boiling point

for every CH2 added about 10 increase