Week 2:

Question 1

Periodic Table:

Answer 1

quantization of energy level in atoms into
orbitals and shells (n) and subshells (s,p,d,f)

ordering of the elements by the number
of electrons in their valence shell.

empirically realized by Dmitri Ivanovich
Mendeleev (1834-1907) long before quantum
mechanics were discovered/developed

Question 2

The Aufbau Principle:

Answer 2

Aufbau: building up

• Atomic orbitals are filled starting from the lowest
  to higher energies with quantum number n
  defining a shell and l defining a subshell

Question 3

Madelung Rule:

Answer 3

Orbital energies increase as (n+l) increases. For two
orbitals with the same (n+l), the one with the smaller n lies lower in
energy

Question 4

drop in energy is not linear with increasing Z (diagram with wavy lines)

Answer 4

e curves cross each other, meaning the order of filling
the orbitals is not only dependent on the quantum numbers.
* Suttle differences are due to shielding and penetration, which we
will discuss in just a moment.

Question 5

Three rules give us the proper order for interactions between electrons:

Answer 5

1. Electrons are placed in orbitals to give the lowest total electronic
   energy. Quantum numbers n and l are filled first.
2. Hund’s rule of maximum multiplicity requires that electrons be
   placed in orbitals to give the maximum total spin possible
   (maximum numbers of parallel spins). Spin multiplicity is number
   of unpaired electrons plus 1 (see text table 2.6).
3. The Pauli exclusion principle requires that each electron in an
   atom have a unique set of quantum numbers. Two electrons in the
   same orbital must be spin correlated, i.e. have opposite spins.

Question 6

Orbital Stabilization:

As the nuclear charge increases
and the electron repulsions, the
3d and 4s levels both become…

Answer 6

…more stable, but 3d moves down
in energy faster than 4s.

-The levels cross each other at
some point. 3d starts higher but
then moves lower than 4s and the
complex way in which the orbital
energies change with increasing
nuclear charge accounts for the
orbital filling.

Question 7

Effective nuclear charge:

Answer 7

net positive charge experienced by an electron
in a polyelectronic atom.

Z* = Z - S

Z* = effective nuclear charge, Z = nuclear charge, S = shielding parameter

Question 8

There are two effects that result in an effective nuclear charge:

Answer 8

1) The direct shielding of the outer electrons from the nuclear charge by the
“inner” electrons.
2) The penetration of “inner” electron density by “outer” electron density.

Question 9

calculation of the energy:

Answer 9

-Pwermitted by Slaters rules:

E = -13.6 (Z/n)2

Question 10

Ionization Energy:

Answer 10

Energy required to remove an electron from a gaseous atom
or ion

• Affected by shielding effects.
• Generally, observe an increase in ionization energy as the nuclear charge
  increases

Question 11

Electron Affinity:

Answer 11

Energy required to remove an electron from a negative
ion (endothermic process)

Question 12

Atomic Radii:

Answer 12

As nuclear charge increases, the electrons are pulled in towards the
center of the atom, and the size of any particular orbital decreases. The opposite is
true for an increase in nuclear charge.

Question 13

Ionic radius:

Answer 13

Radius of a monoatomic ion in an ionic crystal structure.

Question 14

Covalent radius:

Answer 14

Measure of the size of an atom that forms part of one covalent
bond.

Question 15

Van der Waals radius:

Answer 15

radius of an imaginary hard sphere representing the
distance of closest approach for another atom.

The covalent radius is half of the inter-nuclear distance between two identical (or
almost) atoms bonding by a single covalent bond.

• The Van der Waal’s radius is half of the inter-nuclear distance between the nuclei
  of two non-bonding adjacent atoms belong to different molecules.

Question 16

Cations are always smaller than the…

Answer 16

neutral atom and anions are always larger.

• Cations are always smaller than the neutral atom and anions are always larger.

Question 17

Electronegativity: (Linus Pauling)

Answer 17

(“The Nature of the Chemical Bond, 1948):
“the power of an atom in a molecule to attract electrons to itself

• Pauling noticed that the bond energies of polar bonds (A-B) were greater than
  the average of the bond energies of the two homonuclear species (avg. A-A &
  B-B).
• Derives values for electronegativity P by comparing bond energies:

Question 18

Electronegativity: (Mulliken)

Answer 18

• Derives values for an Absolute Electronegativity M from ionization energy and
  electron affinity of atoms or ions

The absolute electronegativity (x) is
defined as the change in energy of the
system as a function of the number of
electrons N present. E.g., for a neutral
species, N → N+1 = EA and N → N-1
= IP, (x) is defined as the negative
tangent at N to the curve shown above

Question 19

Absolute Hardness:

Answer 19

Resistance of the chemical potential to change the number of
electrons. The harder a chemical species, the more difficult it will be to change its
oxidation state.

n = 0.5(IE - EA)

Question 20

Polarizability:

Answer 20

An atom’s ability to be distorted by an electric field.

• Atoms are highly polarizable if their electron distribution can be distorted readily.
• High polarizability is most likely to occur when unfilled atomic orbitals lie close in
  energy to the highest occupied atomic orbital(s), i.e. when there is a small energy
  difference between HOMO and LUMO.
• Closely separated frontier orbitals are typically found for large, heavy atoms and
  ions, whereas small, light atoms typically have widely spaced energy levels

Question 21

Fajan’s Rules: can be used to
predict whether…

Answer 21

a chemical bond
is expected to be predominantly
ionic or valent, and depend on
the relative charges and sizes of
the cation and anion.