Quantum Atomic Theory
Schrödinger (1887-1961)
Heisenberg (1901-1976)
Plank (1858-1947)
Einstein (1879-1955)
Pauli (1900-1958)
Bohr-Rutherford vs Quantum
Similarities
• Electrons have discrete amounts of energy
• Positive nucleus in the centre
Differences:
Bohr-Rutherford
Quantum
• Atom in 2D
• The location of the electrons
can be predicted
• Electrons travel on circular
paths around the nucleus
•Atom is 3D
•The location of the electrons
cannot be predicted
•Electrons move randomly in
‘clouds of probability’
Heisenburg Uncertainty Principle– the position and magnitude of
an electron cannot both be known. If one is measure, the
other is altered.
Orbitals - Redefined
• The Quantum model of the atom describes electrons in
different orbitals (or energy levels) around the nucleus. The
traditional orbits of the Bohr model are subdivided.
• Orbital/Sub-orbital: region around the nucleus where there is a
high probability of finding an electron
• The Period on the periodic table tells you the energy level, the
blocks contained within that period lets you know the kinds of
orbitals the atom has or where the last electrons of that
element are likely to be at any given moment.
The maximum number of electrons in each type of suborbital:
s = 2 electrons maximum
p = 6 electrons maximum
d = 10 electrons maximum
f = 14 electrons maximum
Electron Configuration &
The Periodic Table
• Lists e- location from low to high energy
in the following format
iron atom: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
So what does carbon look like?
• Carbon has 6 e-, so…
• Therefore, the electron configuration of carbon is:
1s2 2s2 2p2
The electron configuration of potassium is:
1s2 2s2 2p6 3s2 3p6 4s1
Cl: 1s2 2s2 2p6 3s2 3p5
becomes
Cl: [Ne] 3s2 3p5
• e- config. Is written only for the outer shell
electrons
• The noble gas indicates all inner shells are full
Shorthand:
[noble gas]
Energy Level Diagrams
6s
5p
5s
E
4s
3s
Hund’s rule – e- half-fill
Pictorial
each orbital in4d
a sublevel
representation of
before–pairing
up
- occupy
Aufbau principle
e
4p
electron distribution
the lowest energyinorbital
orbitals
3d
available
3p exclusion
Pauli
principle
– max 2
p. 188 in
text
e- per orbital (spin up and spin
down)
2p
2s
1s
n = 1 l = 0 m l = 0 ms = ½
• Pauli exclusion principle –
o no two electrons in an atom may have the same four
quantum numbers
o no two electrons in the same orbital may have the
same spin
o only two electrons with opposite spins may occupy an
orbital
• aufbau principle – (German for “building up’)
o each electron is added to the lowest available energy
orbital
• Hund’s rule –
o one electron is placed in each orbital at the same
energy level before the second electron is placed
O
(z = 8)
1s
2s
2p
3s
3p
P
(z = 15)
1s
2s
2p
3s
3p
1s
2s
2p
3s
3p
Ar
(z = 18)
Energy Level Diagrams
6s
5p
4d
5s
E
4p
3d
4s
3p
3s
2p
2s
1s
Anions - Add e- to lowest energy
sublevel available.
Energy Level Diagrams
6s
5p
4d
5s
E
4p
3d
4s
3p
3s
2p
2s
1s
Cations - Remove efrom sublevel with
highest value of n.
Energy Level Diagrams
6s
5p
4d
5s
E
4p
3d
4s
3p
3s
2p
2s
1s
Cations - Remove efrom sublevel with
highest value of n.
Shape of orbitals
• The diagram we used to represent oxygen is;
-
-
8
Protons
-
-
16
8
O
Shape of orbitals

The diagram we might currently use to
represent oxygen is;
Confidence building questions
1.Write out the shorthand notation for the electron configuration of B.
2.Write out the shorthand notation for the electron configuration of Cl.
3.Write out the shorthand notation for the electron configuration of F.
4.Write out the shorthand notation for the electron configuration of Ca.
5.Write out the shorthand notation for the electron configuration of Kr.
6.Write out the shorthand notation for the electron configuration of O2-.
Notice that this is an anion!
7.Write out the shorthand notation for the electron configuration of Na+.
Notice that this is a cation!
8.Why are Groups 1 and 2 referred to as the s-block of the periodic table?
9.Why are Groups 3 through 12 referred to as the d-block of the periodic
table?
10.Using what you now know about electron configurations explain the
notion that elements in the same column in the periodic table have similar
chemical and physical properties.