Chapter 7 - Multi-electron Species and Periodic Properties! Flashcards
what are we ignoring when making approx and assuming wavefunctions in multi-electron species are same as those in one-electron species
ignoring that in multi-electron species theres attraction between electrons and nucleus and repulsion among electrons (based on coulomb’s law)
in multielectron species, what does E of orbitals depend on?
n and fancy cursive l, l because orbitals overlap and in overlapping regeions ^ energy
how does an energy level diagram for multi-electron species look different to single electron species???
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)
Effective nuclear charge (Z eff)
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
what is the spin quantum number and what does it mean????
+ 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
how to draw an orbital (state) diagram
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
how to write electronic configurations
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
Aufbau principle
electrons occupy orbitals with lowest available energy levels before higher energy levels
Pauli exclusion principle
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
Hund’s rule
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
name for orbitals with the same energy
degenerate
s, p, d, f blocks of the periodic table
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
how do n values line up with s,p,d,f blocks on the periodic table
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
procedure for writing electron configurations multielectron species
- locate element in the period table, not period (n value, remember period - 1 for d block) and block (s,p,d,f)
- id noble gas in period above, write it’s symbol in brackets
- starting from left side of the table, write subshells and occupancies until element is reached
- for ionic species, add or remove electrons from neutral configuration
ex. Iron: period 4, d block, [Ar], [Ar]4s^2 3d^6 (no space)
which description of atomic structure aligns with the organization of the periodic table?
quantum mechanical
exceptions to regular electron filling: chromium and copper
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
exception to regular electron filling: transition metal ions
when transition metals ionized, electrons are removed from s subshells first
ex. manganese (II) is [Ar]3d5 NOT [Ar]4s2 3d3
order of orbital filling
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
ground state vs excited electron configuration
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
why are electrons paired spin up and spin down
magnetic fields cancel one another
paramagnetic
species with at least one unpaired electron
diamagnetic
species with all electrons paired
covalent radius
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
ionic radius
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
relationship between increasing nuclear charge & energy/increasing electron repulsion & energy
^ nuclear charge, decreases total energy
^ electron repulsion, increases total energy
van der waal radius
uses the distance between un-bonded atoms to estimate an atom’s boundary
how to approach questions about atomic radii
draw it out!!
trends in atomic radii across the periodic table
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)
relationship atomic radii cations, neutral, anions
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
isoelectric ions
atoms with identical electron configuration
how does atomic radii differ amongst isoelectric ions?
atomic radii decreases with increasing Z
difference between ionic and atomic radius
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
ionization energy
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
how would you represent ionization in an equation
M(g) —–> M+(g) + e-
where M is a general atom, e- is an electron
why are ionization energies measured in the gas phase?
so energy due to interactions between other species doesn’t affect the measurement
how much energy do atoms want to have???
as little as possible! that means lower energy means more stable and most likely higher ionization energy
trends in ionization energy across the periodic table
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
why would the trend in ionization energies reverse for some atoms? how would you approach answering something like that? ex. nitrogen and oxygen
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
how would you explain why second ionization energy is greater?
___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
electron affinity
energy change resulting from addition of a single electron to an atom or ion in gaseous state
trends in electron affinity across the periodic table
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
equation for number of electrons in n shell
2n^2
definition of ionization energy - imp things to include
energy required to remove one electron from a neutral atom in the gas phase
