CHEM1211 - exam 4

pauli exclusion principle

no two electrons can have the same set of quantum numbers

aufbau principle

electrons fill the lower energy orbitals first

hund’s rule

when multiple degenerate orbitals are avalible, electrons fill with paralelle spins before pairing up

max number of electrons in S

2

max number of electrons in p

6

max number of electrons in d

10

max number of electrons in f

14

electron configuration exception for group 6 and 11 (Cr and Cu)

instead of having a full ns orbital and (n-1)d⁴ / (n-1)d⁹, move one electron from ns to (n-1)

excited state electron configuration

electron will be missing from a lower energy subshell and move into a higher energy subshell

anion electron configuration

continue the same filling pattern

cation electron configuration

lose electron from the highest principle energy level first (s > p > d > f)

paramagnetic

unpaired electrons are attected to magnetic feild

diamagnetic

paired electrons are slightly repelled by a magnetic feild

effective nuclear charge

net charge of experienced by one electron in a multi-electron atom

sheilding

electrons that are close to the nucleus sheild the electrons further away from the full nuclear attraction

effective nuclear charge equation

Z_eff = Z - S

Z: atomic number

S: # of core electrons

effective nuclear charge periodic trend

increasing going right across period table

atomic radius periodic trend

increasing from right ot left, top to bottom

why does atomic radius increase from right to left?

the amount of protons changes pulling electrons towards it

why does atomic radius increase going down?

“n” increases so electrons become further from the nucleus

cation atomic radius

smaller then their parent ion because of increased positive charge, which pulls electrons closer.

anion atomic radius

larger then their parent ion because the addition of electrons increases electron-electron repulsion, causing the electron cloud to expand.

polarizability

how easily electrons move away from the nucleus

polarizability trend

larger atoms have higher polarizability

ionization energy

energy required to remove an electron from a gaseoous atom or ion, always exothermic

ionizaion energy periodic trend

increases from left to right, bottom to top

ionization energy group 2/13

requires less energy to remove an electron from group 13 than 2

ionization energy excpetion group 15/16

requires less energy to remove an electron from group 16 than 15

electron attachement enthalpy

the amount of energy released when a neutral atom gains an electron in the gaseous phase

electron attachement enthalpy periodic trend

increaseing from left to right, bottom to top

electron affinity

the energy change associated with the detachment of an electron from an ion in the gaseous phase, endothermic

electron attachement enthalpy exception row 2/3

more favorable to add electrons to period 3 than period 2

electron attachment enthalpy excpetion group 1/2

more favorable to add an electron to group 1 than 2

electron attachement enthalpy excpetion group 14/15

more favorable to add an electron to group 14 than 15

electron attachment enthlply group 18

electron attachement enthalpies are positive

electronegativity

the tendancy of an atom to atract electrons towards itself to form bonds

electronegativity periodic trend

increasing from left to right, bottom to top

lewis symbol

shows the valence electrons of a single atom or ion

lewis structure

shows the bonds and electrons of a molecule

how many bonds an atom can form

max # of valence electrons - # of valence electrons atom has

lewis symbol for anion

have 8 valence electrons

lewis symbol for cations

no valence electrons

lattice energy

the energy change when sepaarated gaeous ions are packed together to form an ionic solid

lattice energy periodic trend

increased from right to left, top to bottom

as lattice energy increases…

charge increases, size decreases

formal charge

tells us if one strucure is better than another, minimize formal charge, negative formal charges on the mose electronegative elements

formal charge equation

FC = (#valence electrons) - (#nonbonding electrons) - (# bonds)

resonance

when more then one valid lewis struture can be written for a molecule

odd electron species

compounds with an odd number of electrons can never achieve an octet for each element

incomplete octets

boron, beryllium, and aluminum form a compound with an incomplete octet

boron tends to form compounds that give it 6 valence electrons

expanded valence shell

elements in row 3 or higher can have more than 8 valence electrons

electron geometry

geometry of atoms and electrons

molecular geometry

geometry of only an atom’s shape

VSEPR theory

structure around a given atom is determined principally by minimizing electron pair repulsions

2 electron groups

linear

3 electron groups

trigonal planar

4 electron groups

tetrahedral

5 electron groups

trigonal bipyramidal

6 electron groups

octahedral

linear, 0 lone pairs

linear

trigonal planar, 0 lone pairs

trigonal planar

trigonal planar, 1 lone pair

bent

tetrahedral, 0 lone pairs

tetrahetral

tetrahedral, 1 lone pair

trigonal pyramidal

tetrahedral, 2 lone pairs

bent “V shaped”

trigonal bipyramidal, 0 lone pairs

trigaonal bipyramidal

trigonal bipyramidal, 1 lone pair

seesaw

trigonal bipyramidal, 2 lone pairs

T shaped

trigonal bipyramidal, 3 lone pairs

linear

octahedral, 0 lone pairs

octahedral

octahedral, 1 lone pair

square pyramidal

octahedral, 2 lone pairs

square planar

linear angle

180

trigonal planar angle

120

bent angle

<120

tetrahedral angle

109.5

trigonal pyramidal angle

<109.5

bent “V shaped” angle

<109.5

trigonal bipyramidal angle

180, 120, 90

seesaw angle

>180, <120

t shaped angle

>180, <120

octahedral angle

90, 180

square pyramidal angle

<90

square planar angle

90

bond order

number of bonding pairs of electrons shared by two atoms in a molecule

bond length

distance between the nuclei of two bonded atoms

bond strength

amount of energy required to break a covalent bond

bond length trend

decreases with bond order

bond strength trend

increases with bond order

averaage bond energy

average amount of energy that must be put into a bond to break it, also endothermic

bond energy equation

bonds broken - bonds formed

exothermic

reactants: weak bonds, products: strong bonds

endothermic

reactants: strong bonds, products: weak bonds

sigma bonds

single bonds

pi bonds

double and triple bonds, contain sigma bonds