Chapter 7 - Multi-electron Species and Periodic Properties!

Question 1

what are we ignoring when making approx and assuming wavefunctions in multi-electron species are same as those in one-electron species

Answer 1

ignoring that in multi-electron species theres attraction between electrons and nucleus and repulsion among electrons (based on coulomb’s law)

Question 2

in multielectron species, what does E of orbitals depend on?

Answer 2

n and fancy cursive l, l because orbitals overlap and in overlapping regeions ^ energy

Question 3

how does an energy level diagram for multi-electron species look different to single electron species???

Answer 3

for single electron species theres no overlap so all orbitals with a specific n value (1,2,3 etc) will have the same level of energy

for multielectron species there is overlap so energy depends on n value and l value 9 with corresponding orbital letter (s, p, d, etc)

Question 4

Effective nuclear charge (Z eff)

Answer 4

an approximation for the nuclear charge felt by valence electrons (based on how core electrons somewhat shield the valence electron from the charge of the nucleus)

calculated with Z - S where Z is atomic number/identity/number of protons and S is number of shielding (non-valence electrons)

NOW you can use the equation for wavefunction of single electron species for multielectron species, subbing Z - S for Z

Question 5

what is the spin quantum number and what does it mean????

Answer 5

+ or - 1/2, represents how electrons interact with an exernal magnetic field, like imagine for + 1/2 electron is spinning to the right so the magnetic field would go from the electrons north pole around to the south pole. for - 1/2, electron spinning to the left so magnetic field extending from the north pole around to the south pole

Question 6

how to draw an orbital (state) diagram

Answer 6

the one where each orbital is a horizontal line with the orbital name written underneath (ex 1s, 2s), and 2 electrons for each orbital rep by the half headed arrows pointing up and down

Question 7

how to write electronic configurations

Answer 7

list ex 1s^2 2s^2 etc (no space between numbers) where first number is n, then which orbital, then how many electrons in that orbital. For s it would be up to 2, for p up to 6, for d up to 10, for f up to 14

Question 8

Aufbau principle

Answer 8

electrons occupy orbitals with lowest available energy levels before higher energy levels

Question 9

Pauli exclusion principle

Answer 9

no two electrons in an atom or ion can have the same four quantum numbers, so two electrons in an orbital must have opposite spin

Question 10

Hund’s rule

Answer 10

with orbitals that have the same energy (ex multiple 2p orbitals) electrons occupy them each single with the same spin before being paired within an orbital

Question 11

name for orbitals with the same energy

Answer 11

degenerate

Question 12

s, p, d, f blocks of the periodic table

Answer 12

s block includes groups 1, 2

p block includes groups 13, 14, 15, 16, 17, 18

d block includes 3-12

f block is two periods at the bottom

Question 13

how do n values line up with s,p,d,f blocks on the periodic table

Answer 13

for s and p, n value = period (1,2,3 etc)

for d block, n = period -1 (first row of d block elements is n = 3)

f block, first row is 4, second is 5

Question 14

procedure for writing electron configurations multielectron species

Answer 14

1. locate element in the period table, not period (n value, remember period - 1 for d block) and block (s,p,d,f)
2. id noble gas in period above, write it’s symbol in brackets
3. starting from left side of the table, write subshells and occupancies until element is reached
4. for ionic species, add or remove electrons from neutral configuration

ex. Iron: period 4, d block, [Ar], [Ar]4s^2 3d^6 (no space)

Question 15

which description of atomic structure aligns with the organization of the periodic table?

Answer 15

quantum mechanical

Question 16

exceptions to regular electron filling: chromium and copper

Answer 16

since 4s and 3d orbitals are close in energy, it’s more favourable to have one electron in 4s, and 1 in each 3d rather than filling 4s first, so for Cr the configuration would be 4s1 3d5

for Cu, configuration 4s1, 3d10

Question 17

exception to regular electron filling: transition metal ions

Answer 17

when transition metals ionized, electrons are removed from s subshells first

ex. manganese (II) is [Ar]3d5 NOT [Ar]4s2 3d3

Question 18

order of orbital filling

Answer 18

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s

also just look at arrangement on the periodic table

Question 19

ground state vs excited electron configuration

Answer 19

if following aufbau, pauli, hunds, exceptions - configuration will be ground state, when configuration violates aufbau or hunds (never pauli, only ever two electrons max in an orbital) the electron is in an excited state

Question 20

why are electrons paired spin up and spin down

Answer 20

magnetic fields cancel one another

Question 21

paramagnetic

Answer 21

species with at least one unpaired electron

Question 22

diamagnetic

Answer 22

species with all electrons paired

Question 23

covalent radius

Answer 23

an estimate for radius determined experimentally for single bonded homonuclear species

ex two hydrogen atoms bonded, find distance between their nuclei, half of this is covalent radius

Question 24

ionic radius

Answer 24

estimate for atomic radii made from a crystal lattice structure and using crystallography measurements of the distance between nuclei in a crystal lattice

kind of same as covalent radius, half of the distance between bonded nuclei

Question 25

relationship between increasing nuclear charge & energy/increasing electron repulsion & energy

Answer 25

^ nuclear charge, decreases total energy ^ electron repulsion, increases total energy

Question 26

van der waal radius

Answer 26

uses the distance between un-bonded atoms to estimate an atom's boundary

Question 27

how to approach questions about atomic radii

Answer 27

draw it out!!

Question 28

trends in atomic radii across the periodic table

Answer 28

atomic radii decrease from left to right along the periodic table because as Z and Z eff increase, probability distributions shift closer to the nucleus (electrons held more tightly) atomic radii increase moving down the periodic table as n increases (THIS AFFECTS RADIUS MORE THAN Z)

Question 29

relationship atomic radii cations, neutral, anions

Answer 29

cations have ^ Z eff and less electron-electron repulsion than neutral, smaller than neutral because pull of protons in the nucleus is greater, especially if all valence electrons are removed anions greater electron repulsion, lower Z eff so larger than neutral

Question 30

isoelectric ions

Answer 30

atoms with identical electron configuration

Question 31

how does atomic radii differ amongst isoelectric ions?

Answer 31

atomic radii decreases with increasing Z

Question 32

difference between ionic and atomic radius

Answer 32

atomic is 1/2 the diameter of a NEUTRAL atom, measuring across the outer stable electrons ionic is half distance between atoms that are just touching, could be the same as atomic, or larger for anions, smaller for cations

Question 33

ionization energy

Answer 33

minimum energy required to remove a single electron from an atom, molecule, or ion in gaseous state related to magnitude of electrostatic repulsion between electrons removed and nucleus and repulsion from other electrons

Question 34

how would you represent ionization in an equation

Answer 34

M(g) -----> M+(g) + e- where M is a general atom, e- is an electron

Question 35

why are ionization energies measured in the gas phase?

Answer 35

so energy due to interactions between other species doesn't affect the measurement

Question 36

how much energy do atoms want to have???

Answer 36

as little as possible! that means lower energy means more stable and most likely higher ionization energy

Question 37

trends in ionization energy across the periodic table

Answer 37

ionization energy increases from left to right, as Z is growing larger and electrons feel greater pull from nucleus. these atoms have lower energy therefore are more stable ionization energy decreases from top to bottom as n is increasing, electrons further from the nucleus so experience less pull, easier to remove

Question 38

why would the trend in ionization energies reverse for some atoms? how would you approach answering something like that? ex. nitrogen and oxygen

Answer 38

first draw the electron configurations for the two atoms to compare in the case of N and O, both have 3 2p orbitals while N has only one electron in each, O has 1 in 2, and 2 in 1. this means theres more electron repulsion and it's a bit easier to remove an electron compared to N

Question 39

how would you explain why second ionization energy is greater?

Answer 39

___cation is smaller due to less electron-electron repulsion and therefore it's valence electron(s) experience a greater Z eff than those in the neutral atom

Question 40

electron affinity

Answer 40

energy change resulting from addition of a single electron to an atom or ion in gaseous state

Question 41

trends in electron affinity across the periodic table

Answer 41

affinity increases across the periodic table because those atoms are smaller, so electron will feel greater attraction to the nucleus UNTIL halogens because they have very low electron affinities electron affinity decreases down a the table because entering shell greater n value, higher in energy, and less attraction to the nucleus

Question 42

equation for number of electrons in n shell

Answer 42

2n^2

Question 43

definition of ionization energy - imp things to include

Answer 43

energy required to remove one electron from a neutral atom in the gas phase