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Chemistry VSEPR Cheat Sheet (DRAFT) by

Molecular Geometry, Intermolecular Forces, Valence Bond Theory, etc.

This is a draft cheat sheet. It is a work in progress and is not finished yet.

Molecular Geometry

VSEPR
electron pairs in the valence shell of an atom repel one another and will arrange themselves to be as far apart as possible, minimizing the repulsive intera­ctions between them
electron domain
a lone pair or a bond, regardless of whether the bond is single, double, or triple
electron domain geometry (EG)
arrang­ement of electron domains (bonds and lone pairs) around the central atom
molecular geometry (MG)
arrang­ement of bonded atoms
when there are no lone pairs on any of the the central atoms, MG is the same as EG.
bond angle
angle between two adjacent bonds in a molecule or polyatomic ion
AB5 molecules contain 2 bond angles because positions occupies by bonds in a trigonal bipyramid are not all equal.
1) equato­rial: 3 bonds arranged in a trigonal plane 2) axial: 2 bonds that form an axis perpen­dicular to the trigonal plane

Valence Bond Theory

valence bond theory
bonds form between atoms when atomic orbitals overlap, thus allowing the atoms to share valence electrons.
 
each of the overla­pping atomic orbitals must contain a single, unpaired electron. the two electrons shared by the bonded atoms must have opposite spins.
 
the nuclei of both atoms are attracted to the shared pair of electrons.
 
the mutual attraction for the shared electrons holds the atoms together.
valence bond theory explains why covalent bonds form.
a covalent bond will form between two atoms if the potential energy of the resulting molecule is lower than the combined potential energies of the isolated atoms.
 
formation of a covalent bond gives off energy; energy must be supplies to a molecule to break covalent bonds.
summary of valence bond theory: a bond forms when slightly occupied atomic orbitals on two atoms overlap. the two electrons shared in the region of orbital overlap must be of opposite spin. formation of a bond results in a lower potential energy for the system.

Interm­ole­cular Forces

interm­ole­cular forces
electr­ostatic attrac­tions between opposite charges or partial charges.
particles in the condensed phases (solids and liquids) are held together by interm­ole­cular forces.
van der waals forces
interm­ole­cular forces acting between atoms or molecules in a pure substance.
include dipole­-dipole intera­ctions (which include hydrogen bonding) and dispersion forces.
dipole­-dipole intera­ctions
attractive forces that act between polar molecules.
partial positive charge on one molecule is attracted to the partial negative charge on the neighb­oring molecule.
the larger the dipole, the larger the attractive force.
hydrogen bonding
a strong type of dipole­-dipole intera­ction that occurs in molecules containing hydrogen bonded to a small, highly electr­one­gative atom, such as nitrogen, oxygen, or fluorine.
within a series of hydrogen compounds of group 14, the boiling point increases with increasing molar mass.
for groups 15-17, the same trend is observed for all but the smallest member of each series, which has an irregu­larly high boiling point.
dispersion forces
attractive forces that act between all molecules, nonpolar and polar.
forces between an instan­taneous dipole (fleeting, temporary dipole) and induced dipoles (insta­nta­neous dipoles can induce neighb­oring dipoles).
when a nonpolar molecule acquires an instan­taneous dipole, it is polarized.
magnitude of dispersion forces depends on how mobile the electrons in the molecule are.
in small molecules, the electrons are relatively close to the nuclei and cannot move about very freely. thus, the electron distri­bution is not easily polarized.
in larger molecules, the electrons are farther away from the nucleus and can move about more freely. thus, electron distri­bution is easily polarized, resulting in larger instan­taneous dipoles, larger induced dipoles, and larger interm­ole­cular forces overall.
the more valence electrons a compound has, the more easily polarized it is.
if a compound is polar, dipole­-dipole forces and dispersion forces are acting on it.
if a molecule is polar with a hydrogen bonded to fluorine, nitrogen, or oxygen, dipole­-dipole (including hydrogen) and dispersion forces are acting on it.
if the compound is nonpolar, only dispersion forces are acting on it.
ion-dipole intera­ctions
attraction between ions and polar molecules in solutions.
magnitude depends on the charge and size of the ion and on the dipole moment and size of the polar molecule.
cations interact more strongly with dipoles than anions because they tend to be smaller. an ion with higher charge and smaller size will interact more strongly with water molecules.
strongest force to weakest force:
ion-di­pole, hydrogen, dipole­-di­pole, dispersion

Hybrid­ization of Multiple Bonds

sigma bonds
first bond formed and stronger than pi bonds
shared electron density concen­trated directly along the intern­uclear axis
overlap of hybrid orbitals that point directly toward each other
contribute more to overall bond strength than pi bonds
pi bonds
second bond formed and weaker than sigma bonds
form when parallel, unhybr­idized p orbitals sideways overlap.
electron density concen­trated above and below plane of molecule.
less effective than sigma bonds and contribute less to the overall bond strength than sigma bonds.
singly occupied p orbitals give rise to multiple bonds.
one sigma and one pi bond together constitute a double bond.
one sigma and two pi bonds together constitute a triple bond.
pi bonds restrict the rotation of a molecule in a way that sigma bonds do not.
double bonds consist of one sigma bond and one pi bond; the pi bond restricts rotation about the sigma bond, making the molecules rigid, planar, and not interc­han­geable.
 

Deviation from Ideal Bond Angles

lone pairs take up more space than bonding pairs.
a lone pair on the central atom is attracted only to the nucleus of that atom. a bonding pair is simult­ane­ously attracted to the nuclei of both the bonding atoms. lone pairs have more freedom to spread out and greater capacity to repel other electron domains.
multiple bonds repel more strongly than single bonds because they contain more electron density.
multiple bonds are shorter than single bonds.

Bonding Theory

in species that can be repres­ented by two or more resonance struct­ures, the pi bonds are deloca­lized, meaning that they are spread out over the molecule and not constr­ained to just two atoms.
localized bonds are those constr­ained to two atoms.

Hybrid­ization of Atomic Orbitals

hybrid­ization
atomic orbitals mix to form hybrid orbitals.
hybrid orbitals
orbitals formed by hybrid­ization of some combin­ation of s, p, or d atomic orbitals.
each sp hybrid orbital has one small lobe and one large lobe.
they are oriented in opposite directions with a 180 degree angle between them.
number of electron domains is equal to the number of hybrid orbitals which is also equal to the number of atomic orbitals needed to mix.
elements in the third period of the periodic table and beyond do not obey the octet rule because they have d orbitals that can hold additional electrons.
in molecules where there are more than 4 electron domains on the central atom, include d orbitals in hybrid­iza­tion.
 
sp3d have shapes similar to sp, sp2, and sp3 orbitals - one large lobe and one small lobe.
do not use hybrid orbitals to predict molecular geometries but to explain geometries that are already known.
  

Molecular Geometry and Polarity

polar molecule
a bond between two atoms of different electr­one­gat­ivities
a molecule bonded with a highly electr­one­gative element will be polar.
highly electr­one­gative elements: fluorine, nitrogen, oxygen, or chlorine
if the central atom has lone pairs, molecule is most likely polar
if vectors cancel each other, molecule is nonpolar.
if vectors do not cancel each other (are not equal) the molecule is polar.
structural isomers
molecules that have the same chemical formula but have different arrang­ement of atoms

Molecular Orbital Theory

molecular orbital theory
atomic orbitals involved in bonding combine to form new molecular orbitals that are associated with the whole molecule rather than with individual atoms.
molecular orbital theory predicts that the more effective the intera­ction or overlap of the atomic orbitals, the lower in energy the resulting bonding molecular orbitals will be and the higher in energy the resulting antibo­nding molecular orbitals will be.
fill in order of increasing energy.
molecular orbitals
orbitals that result from the intera­ction of atomic orbitals of bonding atoms.
 
similar to atomic orbitals - they have specific shapes and energies and can accomm­odate a maximum of two electrons.
 
two electrons residing in the same molecular orbital must have opposite spins.
 
number of molecular orbitals produced is equal to the number of atomic orbitals that combined.
bonding molecular orbital
lower in energy than the atomic orbitals that combined to produce it.
concen­trated between the nuclei, along the intern­uclear axis.
result from the intera­ction of atomic orbitals that are in phase.
antibo­nding molecular orbital
higher in energy than the atomic orbitals that combined to produce it.
most of its electron density is located outside the intern­uclear region.
result from the intera­ction of atomic orbitals that are out of phase.
sigma molecular orbitals
lie along the intern­uclear axis
bonding molecular orbitals are designated sigma1s; antibo­nding orbital is designated sigma1s*
bond order
indicated how stable a molecule is
higher bond order, the more stable the molecule; higher bond order, shorter bond length.
bond order = # electrons in bonding molecular orbitals - # electrons in antibo­nding molecular orbitals/2
pi molecular orbitals
regions of electron density affect both nuclei but do not lie along the intern­uclear axis.
2px orbitals lie along the intern­uclear axis and point directly toward each other, and combine to form sigma molecular orbitals.
2py and 2pz orbitals are aligned parallel to each other and combine to form pi molecular orbitals.
designated pi2py,z and pi2py,x*
molecular orbitals resulting from the combin­ation of p atomic orbitals are higher in energy than the molecular orbitals resulting from the combin­ation of s atomic orbitals.
px orbitals, which lie along the intern­uclear axis, overlap most effect­ively, giving the lowest­-energy bonding molecular orbital and the highes­t-e­nergy antibo­nding molecular orbital.

Hybrid­ization Chart

Number of Electron Domains on Central Atom
2
3
4
5
6
hybrid orbitals
sp
sp2
sp3
sp3d
sp3d2
geometry
linear
trigonal planar
tetrah­edral
trigonal bipyra­midal
octahedral