Chem unit 3

Bonding orbitals

in phase overlap of atomic wave function

lower in energy than parent atom

increased electron density between nuclei

no astri

antibonding orbitals

atoms overlap out of phase

higher in eneergy than og atom

electron density pushed outside nucleii

bond order

Bond Order = 1/2 (# of e– in bonding MOs – # of e– in antibonding MOs)

number of bonds shared between two atoms bonding

magnitude of bonding interaction between adjacent orbitals is governed by

energy difference

orbital overlap

bond length vs strength

shorter bonds mean stronger bond

electronegativty

tendency of an atom to attract shared electrons in a mokecule to itself

electronegatvity across a row

increases due to higher effective nuclear charge and higher proton number that attracts the shared electrons

electronegativity down a group

decreases because lower effective nuclear charges and larger atomic radius

permanent dipole

due to uneven electron distribution

dipole = charge magnitude x distance between atoms.

electron group geometry

2 groups: linear (180)

3 groups: trigonal planar: 120 degrees

4 groups: tetrahedral (109.5)

formal charge

hypothetical charge of a given atom in

a molecule determined by assuming that

• nonbonding electrons belong solely to the atom

• bonding electrons are shared evenly, regardless of

electronegativity differences

= valence electrons - non bonding electrons - ½ bonding electrons

lewis structures provide esmallest formal charge

valence electrons- individual dots - bonds

structural isomers

identical chemical formulas but differ in how the atoms are connected to eachother

sigma bond

first bond formed

single line

head on overlap

allows atoms to be rotated around the axis

sp hybridization

one s orbital and one p orbital (2 total)

linear geometry

sp² hybridization and sp³

one s orbital and 2 p orbitals

trigonal planar

vs.

4 total (including lone pairs)

tetrahedral

one s adn 3 p

molecular vs electron geometry

electron geometry is of all electron groups including the lone pairs while molecualr geometry only focuses on the bonded pairs

linear molecular and electron geometry

2 groups 0 lone pairs

bent molecular geometry

3 groups, one lone pair

trigonal planar electron geometry

3 dmoains, 1 lone pair or 3 domains 0 lone pair

geometric isomers

same connectivity but different spatial awareness

does not have rotation (a bond would have to be broken)

polar molecule

uneven distribution of electron charge

creates permanent dipole moment

slightly pos and slightly neg end

no unpaired electrons

diagmagnetic (not attracted to the magnetic field)

hydrogen bond

one h-x bond where x is F O N usually,

x also has at least one lone pair

pi conjugation

sharing of pi electrons accross 3 or more atoms when p orbitals overlap

lowers potential energy bc negative charge of electrons is distrubited over larger area, minimizes electron electron repulsion and electrons are attracted to multiple nuclei

look for lone pairs in p orbitals next to or separated by one sigma bond from a pi orbital