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UNIT 3 AP CHEMISTRY REVIEW 
 
INTERMOLECULAR FORCES 
IMFs are forces between molecules in ​covalently bonded​ substances 
Need to be ​broken apart​ so substances may ​change phases 
➔ In comparison, when ionic substances change phases, bonds between ions are 
broken 
➔ For CV substances, bonds between atoms remain intact 
 
Dipole-dipole forces​ ​occur between polar molecules, attraction between positive end 
of a polar molecule to negative end of another 
Dipole-dipole attractions are ​relatively weak​, so substances melt and boil at very low 
temperatures 
➔ Most are gases or liquids at room temperature 
The greater the polarity, the greater the dipole-dipole attraction, so molecules with 
larger dipole moments​ have ​higher melting and boiling points 
 
Hydrogen bonds​ are a ​special type of dipole-dipole attraction 
Positively charged hydrogen end of molecule is attracted to negatively charged end of 
a molecule containing an extremely electronegative element ​(fluorine - F, oxygen - O, 
nitrogen - N) 
Much stronger than normal dipole-dipole forces bc hydrogen has a single electron, so 
its positive nucleus is left unshielded after giving that up to a bond 
Substances like ​water and ammonia​ have ​high melting and boiling points​ bc of 
hydrogen bonds  
Water’s solid form is less dense than liquid form bc in the crystal structure, hydrogen 
bonds force molecules farther apart 
 
London Dispersion forces​ ​occur between all molecules, very weak attractions caused 
by random movement of electrons  
Temporary induced dipole-induced dipole attraction​ when nonpolar molecule may 
have more electrons on one side than the other at a given moment, creating an 
instantaneous dipolarity 
➔ Molecule acts as a ​very weak dipole 
Molecules with ​more electrons​ experience ​greater London dispersion forces 
➔ When comparing substances with only dispersion forces, one with more 
electrons will have higher melting and boiling points - look at molar mass 

London dispersion forces are even weaker than dipole-dipole forces, so substances 
that only experience them will ​melt and boil at extremely low temperatures​ and tend 
to be ​gases at room temperature 
 
Bond Strength 
Ionic substances tend to be ​solids at room temperature​, so melting them into liquids 
requires bonds holding lattice together to be broken 
Amount of energy needed to do so depends on the ​Coulombic attraction​ between the 
molecules 
Relative strength of IMFs within molecules 
1
...
Dipole-dipole forces 
a
...
London dispersion forces 
a
...
08206 L atm mol​-1​ K​-1 
T​ = absolute temperature of gas (K) 

 
Can be manipulated, ​Combined Gas Law ​for when number of moles is held constant 
P ​= pressure of gas (atm) 
V​ = volume of gas (L) 
T ​= absolute temperature of gas (K) 
 
 
Simple relationships: 
➢ If the volume is constant: As pressure increases, temperature increases; as 
temperature increases, pressure increases 
➢ If the temperature is constant: As pressure increases, volume decreases; as 
volume increases, pressure decreases 
○ Boyle’s Law 
➢ If the pressure is constant: As temperature increases, volume increases; as 
volume increases, temperature increases 
○ Charles’s Law 
 
DALTON’S LAW 
Total pressure of a mixture of gases is the sum of all of the partial pressures of the 
individual ases 
 
Partial pressure of a gas is directly prop to the number of moles of gas present in the 
mixture, X​Gas 1 is
often called the mole fraction 
​

 
 
DERIVATIONS FROM IDEAL BEHAVIOR 
When the temperature is low and/or the pressure is high, gases behave in a 
less-than-ideal manner 
➔ Assumptions of kinetic molecular theory become invalid when gas molecules 
are packed too tightly together 
When packed too tightly together, 
➢ Volume of gas molecules becomes significant and the actual volume of non 
ideal gas is larger than the volume predicted by the ideal gas equation without 
taking gas molecules into account 

➢ Gas molecules attract one another and stick together, as the IMFs become 
significant, and there are fewer particles bouncing around and creating 
pressure, making the pressure of a nonideal gas smaller than the pressure 
predicted by the ideal gas equation 
Stronger IMFs will lead to more deviations 
➔ H​2​O, with hydrogen bonding, is more likely to deviate from ideal behavior than 
CH​4​, with only London dispersion forces 
When gases have similar IMFs, gases with more electrons are more polarizable and 
more likely to deviate from ideal behavior 
➔ For noble gases with only LDFs, argon is more likely to deviate than helium, but 
less likely to deviate than xenon 
 
DENSITY 
Measured in the same way of a liquid or solid, D = m/v 
D​ = density, ​m​ = mass of gas (grams), ​V​ = volume occupied by gas (liters) 
 
Determine density of a gas sample by combining density equation w/ Ideal Gas Law 
➔ Substitute V = m/D and rearrange until D isolated 
◆ m/n determines mass per mole, or MM, and that can also be rearranged 

➔

 

 
SPECTROSCOPY AND THE EM SPECTRUM 
Differences in absorption or emission of photons in different regions are related to 
different types of molecular motion or electronic transition: 
➔ Microwave radiation​ - transitions in molecular rotational levels 
➔ Infrared radiation​ - transitions in molecular vibrational levels 
➔ Ultraviolet and visible radiation​ - transitions in electronic energy levels 
 
PHOTOELECTRIC EFFECT 
When an atom or molecule absorbs/emits a photon, the energy of the species is 
increased/decreased by an amount equal to the photon’s energy 
Equation relates wavelength of electromagnetic wave to its frequency and the speed of 
light 

Planck’s equation relates energy of a photon to frequency of electromagnetic wave 
 
BEER-LAMBERT LAW 
Relates absorption of light by solution to three variables w/ equation 

 
In most experiments, path length and wavelength of light are constant 
➔ Absorption would only be proportional to concentration of absorbing molecules 
or ions 
➔ If you leave fingerprints on the sides of the cuvette, it will scatter the detector 
and absorb light, causing higher absorbance and higher concentration results 
➔ If the student orients the cuvette so that the path of the light is through the 
frosted sides of the cuvette, little light will be able to reach the detector
...
 
➔ If the student forgets to rinse the cuvette with blue dye solution first so that 
some distilled water droplets are still in the cuvette, the measured absorbance 
will be lower than the real absorbance of the dye solution
...
Notice that a cuvette has a volume of 
approximately 1-2 mL, so a few droplets are significant in volume change
...
 
 



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