chem exam 2

hunds rule

fill in each orbital with 1 electron first then go back and fill the rest

what does electron energy actually mean

high energy means its easier to remove electrons and less stable. the higher the n value, the higher the energy. energy increases with more spdf because of increased shielding. shielding makes them repel each other, which makes it easier to remove

pauli exclusion

an orbital can only hold 2 electrons. no 2 electrons in the same atom can have the exact same quantum numbers, differentiated by spin

zeff

determined by number of protons and shielding by electrons. higher zeff leads to lower energy. shielding reduces full nuclear charge to zeff

what happens when theres a greater nuclear charge

lowers sublevel energy because they are strongly attracted to each other, making it more stable

how to tell the max amount of electrons in a subshell

the amount of elements in that row

the more shielding,

the less attraction

where do you remove electrons from

remove from the highest n value first. then, go from f d p s

where do you add electrons to

add to lowest n value that still has room, then go to s p d f

inner/core electrons

lower energy electrons. electrons in noble gas configuration, take the configuration of the nearest noble gas and those are your core electrons. AND any completed transition series (d10 or f14).

ex: [Xe]6s^s5d¹⁰4f¹⁴6p²

the electrons in xe are core, 4f14 and d10 are also core.

outer electrons

electrons with the highest n value

ex: [Xe]6s²5d¹⁰4f¹⁴6p²

6s2 and 6p2 would be outer electrons

valence electrons

involved in compound formation. for main group (1A-8A), the valence are the same as the outer electrons, same # as group number

transition metals valence electrons

put the ng configuration in brackets. the remaining electrons are valence

paramagnetic

has unpaired e, attracted by a magnetic field.

when determining which one is more paramagnetic, its the one with more unpaired electrons

diamagnetic

all electrons are paired, slightly repulsed by magnets

exceptions to electron config

Cu, Ag, Au, Cr, Mo, Nb, Ru, Rh

double exception to electron configuration

Pd

how do the exceptions to electron configuration work

put into noble gas configuration and proceed. but, take 1 electron from the s and put it into d. for pd, take 2

isoelectric

same electron configuration

main group elements valence and outer electrons

same thing

shielding and zeff relationship

inversely proportional

zeff trend

as it moves to the right, there are more protons. this pulls the electrons closer together. when you go up a period, there are more shells meaning they get further away. it cant pull as well when its far

atomic size trend

transition elements (d block) dont follow this pattern. theyre all approx the same size because zeff is counteracted by electron shielding

size comparison for ions

whatever element it corresponds to with the same number of e, assume that size. if there are 2 isoelectric elements, the one with the most protons is smaller because it pulls the electrons closer together

penetrating

can go closer to the nucleus. 4s is more penetrating than 4d because it can go closer

ionization energy

the amount of energy needed to remove (endothermic rxn) a mol of electrons from a mol of neutral gaseous atoms

ionization energy trend

inverse to size. as it gets smaller, they are more attracted and harder to separate

ionization energy exceptions

Be>B

Mg>Al

Ca>Ga

N>O

P>S
As>Se

successive ionization energies

removing more than 1 electron, gets bigger because attractive forces increase. when you see a large jump, thats where valence electrons end

pseudo noble gas

when an metal ion has the same configuration as a noble gas and its core electrons are completely full

metallic and nonmetallic character trend

nonmetallic character

same as acidity and covalent oxides

metallic character

same as basicity and ionic oxides

amphoteric oxide

an oxide that can act as either an acid or base in water

electron affinity

attraction to electrons, an energy level. the energy released (exothermic) when 1 mol of electrons is added to 1 mol of gaseous atoms

magnitude of electron affinity trend

usually negative because its releasing energy. more negative from left to right but no clear trend from top to bottom

ea exceptions

2A and 8A ea is 0 because they dont want to gain electrons

element with positive ea

nitrogen

exception for ea

Cl>F

ea endo and exo

only the first electron added is exothermic, the rest after are endothermic

atoms with low ea

form cations

atoms with high ea

form anions

noble gases ea and ie

very high ie and slightly positive eas. dont want to lose nor gain electrons

aufbau principle

electrons are always placed in the lowest energy sublevel

elements in the same groups

similar electron configs correlate with similar chemical behavior

atomic size increases

when n increases and decreases with inc zeff because electrons are pulled closer

oxidation numbers

what somethings tendency to lose/gain electrons when forming compounds

on for 1A-3A

just 1, 2, 3 because they always lose electrons

on for 4A-8A

can either be positive or negative (highest and lowest oxidation number)

either

+4/-4

+5/-3

+6/-2

+7/-1

inert pair configuration

metal only loses np electrons

ex [Kr] 5s²4d¹⁰5p²

[Kr] 5s²4d¹⁰

what determines if something will be a pseudo ng or inert pair

what it bonds with

cations size

smaller than parent atoms. size continues to decrease as charge increases

-3>-2>-1>1>2>3

anions size

larger than parent ions

3 types of chemical bonding

ionic, covalent, metallic

lewis dot structure for metals

the total number of dots is the number of e a metal loses to form a cation

lewis dot structure for nonmetals

number of unpaired electrons equals the number of electrons it shares to form a covalent mond and the number of electrons the atom gains to form an anion

octet rule

when atoms bond, they lose, gain, or share e to attain an outer filled level of 8 electrons. exception: Li and H only need 2`

ionic binding model

ionic bonds form when e are transferred to form ions, which attract to form solid compounds

born haber cycle

looks at energies involved in making a substance in a chemical reaction

combining ions to form a crystal

exothermic process

lattice energy

energy required to separate 1 mol of ionic solid into gaseous ions. always has a positive H, magnitude depends on size, charge, and arrangement

coulombs law

electrostatic energy (cation charge*anoion charge/cation radius+anion radius)

how is lattice energy affected by ionic charge

the greater the ionic charge, the greater the lattice energy because its harder to break the bonds.

properties of ionic compounds

hard, rigid, brittle, with high melting points. do not conduct energy in solid states because they are fixed in place in the lattice. can conduct when melted or dissolved and carry a current

how are melting points related to lattice energy

the higher the lattice energy, the higher the melting point

covalent when theyre far vs close

changing the distance affects potential energy. when theyre closer together, reactions become more energetically favorable because theyre attracted to each other. theres a middle point where it has the most energy but when they get too close, they begin to repel.

shared pairs/bonding pairs

atoms that share electrons

lone pair

outer level pair not involved in bonding

bond order

number of electron pairs being shared by a given pair of electrons.

ex single bond has bond order of 1

double bond has bond order of 2

bond energy

energy needed to overcome the attraction between nuclei and shared e (covalent). the stronger the bond, the higher the bond energy. you cant just double the energy of the previous bond, you need to find it on the table

bond length

distance between nuclei of bonded atoms

relationship between bond order and bond length and bond enerfy

higher bond order results in shorter bond length because density is increase. shorter bonds are also stronger

bond length trend

increases down a group and decreases across a period

polar covalent bond

shared electron pair isnt shared equally, remains closer to 1 atom than the other

electronegativity

ability of an e in a covalent bond to attract shared electron pair

electronegativity and size

inversely related, the larger a molecule gets, the less it can attract things

most electronegative element

fluorine. as you move away from f, it gets less electronegative

electronegativity oxidation numbers and formula ex HCl

the more en atom is assigned all the shared electrons while the less atom is assigned none.

number of valence electrons- (number of shared electrons + number of unshared electrons)

ex: HCl has 7 valence electrons

Cl= -1

shared e: 2

unshared e: 6

H= +1

shared e:0

unshared e: 0

valence e: 1

depicting polar bonds

shown by the arrow, points towards the more electronegative element

properties of metals

solids with middle to high melting points and very high boiling points. can be shaped without breaking because of sea of metals slide past each other. good conductors of energy and heat as solids and liquids

∆EN

ranges from 0 (for diatomic or elements with the same en) to 3.3 for CsF. the bigger the number, the more ionic the bond

% ionic character

increases with ∆EN, there are always none that are 100% ionic

electroon sea

all metal atoms contribute their valence electrons to form a delocalized sea. the metal ions lie in an orderly array. held together by attraction between metal atoms cations and valence electrons

alloys

2 or metals with variable composition