3.5 - Atomic Structure & Periodic Table

Question 1

orbitals of multi-electron atoms

Answer 1

The size and energy of their orbitals are
different than for Hydrogen (single electron)
because of the interplay between nuclear
attraction and the repulsions from other
electrons.

Energy in the sublevels increases as follows:
s<p<d<f

Question 1

aufbau principle

Answer 1

Theory stating that an atom is “built up” by the
addition of electrons, which fill orbitals starting at
the lowest available energy orbital before filling
higher energy orbitals. (eg. Fill 1s first before 2s)

The location and number of electrons in the
energy levels of an atom or ion is called
ELECTRON CONFIGURATION.

Can also be represented by an ENERGY-LEVEL
DIAGRAM (ORBITAL DIAGRAM).

Question 2

hund’s rule

Answer 2

Rule stating that in a set of orbitals of the
same energy, the lowest energy configuration
for an atom is the one with the maximum #
number of unpaired electrons allowed by the
Pauli Exclusion Principle.

Ie. Do not pair up the electrons yet until all
other sublevels are full at that energy level.

Question 3

Transition metals

Answer 3

Elements whose highest-energy electrons are in d
orbitals.

Some exceptions to the Rule: Cr-

Supposed to be [Ar]4s^23d^4. However it is not. It is:

[Ar]4s^13d^5.

Question 4

half full effect

Answer 4

an unfilled subshell is less stable than half-filled or filled subshell.

It is more stable for the atom’s subshell to be half filled or filled.

Question 5

Ions

Answer 5

Lose electrons or gain electrons from the
highest energy level.

Eg. Pb: [Xe]6s24f145d106p2.

Pb2+: [Xe]6s24f145d10