CHEM 20A - Midterm 2

Ground State Electron Configuration

allowed and lowest energy, lower levels filled, parallel spin

Excited State Electron Configuration

allowed but at higher energy, lower levels NOT filled, NOT parallel spin, (violates Hund’s or Aufbau’s)

Forbidden Electron Configuration

not allowed (violates Pauli exclusion)

Column - changes in n (size of orbital)

going down a column → n inc, e are further away from nucleus

Row - Changes in effective nuclear charge

going to right → nuclear charge inc, more protons and neutrons, not good sheidling

atomic radius

distance of e from nucleus

atomic radius change

inc. down a group, inc left (bigger orbital, bigger atom)

electron affinity 

energy released when e is added 

electron affinity change

inc up, inc right, (stronger nuclear attraction, want e)

larger electron affinity

easier to add e (wants e)

smaller electron affinity

harder to add e (doesn’t want it)

ionization energy

energy required to move e

ionization energy change

inc up, inc right (same as electron affinity) (harder to remove e)

higher ionization

harder to remove e (want to keep e)

lower ionizaton

easier to remove e (want to give e away)

half filled subshell

extra stability, so hard to remove e 

valence electrons

number of e in orbitals of the highest n

core electrons

all other e

electronegativity

electron pulling power

ionic compound

diff in EN is greater than 2.0, e are NOT shared, (metal + nonmetal)

covalent compounds

diff in E is less than 2.0, e are shared (metal + metal)

anions

take e, negative (nonmetals)

cations

give e, positive (metals)

lattice energy

bond strength btw cations and anions

ionic radi __ as electrons inc

inc

ionic radi __ as protons inc

dec

lattice energy is bigger when ionic radi is __

smaller

Steps for Drawing Lewis Dot Structure

1) count total avail e

2) count total e needed

3) calcuate # of bonds ½ (e needed - e avail)

4) determine # of e remaining (e avail - 2 x bond number)

5) calculate formal charge (valence e - # of lone pair - # of bonds)

resonance structures

structures with same number of atoms but different arrangement of e

isomers

structures with diff arangements of atoms

elements with exapnded octect

elements below period 2

molecules with incomplete octet

B and Al

radicals

molecules with unpaired e, very reactive

internulcear PE - atoms are far apart

baseline, PE = 0

internulcear PE - moving closer

e are attracted by nucleus, attraction > repulsion, more stable, lower PE

internulcear PE - bond length

attraction = repulsion, PE is lowest

internulcear PE - too close

repulsion > attraction, less stable, higher PE

dissocation energy

energy needed to pull atoms apart/break bonds

bigger atomic radius mean __ bond length

longer

longer bond = length means ____ bond

weaker

all covalent bonds between different elements are

polar to some extent

only covalent bonds btw atoms of the same element are

nonpolar

VSPER Model

no mixing, explains the shape of molecules, doesn’t have orbitals 

Valence Bond (VB) Theory

same atoms form hybrid orbitals, explains locations of electrons in VSPER structure

steric #

number of bonds + number of lone pairs around the central atom 

lone pairs and double bonds

take up more space so angles are slightly smaller

the higher the electronegativity

the more polar the bond 

Radicals

molecules w/ unpaired electrons

bigger atomic radius means

longer bond length

dissociation energy

energy needed to pull atoms apart/break bonds

larger dissociation energy means

stronger bond

longer bond length means

weaker bond

nonpolar covalent bonds

covalent bonds btw atoms of the same element

hybrid orbitals

combination of s,p, anf d atomic orbitals

𝜎 bonds

formed by hybrid + hybrid or hybrid + H

π bonds

formed by p-orbital + p-orbital

Molecular Orbital Theory 

molecular orbitals are made by combining orbitals from different atoms, energy level diagram, (diagnetic vs. paramagnetic)

bond order

½ ( # electrons in bonding MO - # of electrons in antibonding MO)

diamagnetic 

does not have unpaired electrons

paramagnetic

has unpaired electrons

atomic orbitals with more electronegative elements have

lower energy

When at least one element is group 5 and below

σs < σs* < πp < σp < πp* σp*

when at least one element is group 6 and above 

σs < σs* < σp < πp < πp* σp*

molarity

mole/liter

mole 

concentration (M) x volume (L)

gram-to-gram conversion

1) find molar mass,

2) gram to gram conversion (molar mass, rxn)

find empirical formula

1) add all massessand divde by smallest one to get ratio

precent yield

actual yield/ theoretical yield

Magic # 

initial/coefficient, limiting reagent has smallest magic #)

Aufbau 

add electron to orbitals w/ lowest energy

Hund’s

if more than one orbitals have same energy, add electrons to empty orbital with paralel spin rather than pairing up

pauli’s exclusion

each orbital can only have 2 electrons and they must be opposite spin

Why is first ionization energy of B lower than Be or Al lower than Mg?

bc removing electrons from p orbital rather than s

why is first ionization energy of O lower than N and S lower than P? 

bc N and P have half-filled subshells 

there is a big jump in successsive ionization energy when removing

core e

more e → more e-e repulsion

bigger size

more p → more nuc-e attraction

smaller size

lattice energy is bigger when 

charges are higher, ionic radii is smaller 

lattice energy is smaller when

charge is lower, ionic radii is larger

which direction does the dipole moment point?

from lower to higher electronegativity

bond w/ greatest electronegativity

is most polar bc electrons are pulled more unevenly

bon length and bond strength are 

inversely proportional 

bond length and atomic radius are

directly proportional

mole

concentration x volume

what degress is equatorial?

120

what degrees is axial?

90

Sn 2

linear

sn 3

trigonal planar, lone pair = bent 

sn 4

tetrahedral, lone pair = trigonal pyramidal 

sn 5

trigonal pyrmidal, 

sn 6

octahedral