3.6: ATOMIC STRUCTURE AND THE PERIODIC TABLE
CREATING ENERGY-LEVEL DIAGRAMS FOR ATOMS
See Fig 2, p. 187.(Draw the sample diagram into your notebooks)
Diagram:
DISTRIBUTION OF ELECTRONS IN ORBITALS
Begin with Hydrogen → one proton and one electron in the Is Orbital.
Then each time we add a proton to the nucleus, we also add an electron to an available
orbital.
Rules: 1. Electrons occupy the lowest energy orbital of the lowest energy level first. (The
Aufbau Principle)
2. No electron pairing takes place in the p,d or f orbital until each orbital of a given
set contains one electron (Hund’s Rule)
3. No orbital can contain more than two electrons.
(Pauli Exclusion principle) states that no two electrons in the same atom may
have identical values for all four quantum numbers.
creating energy-level diagrams for anions (negative charge ions)
The energy-level diagrams for anions, are done using the same method as for atoms. But
you need to add the extra electrons corresponding to the ionic charge to the total number
of electron before distribution of the electrons on the orbitals (See p190)
creating energy level diagram for cations (positive charged ions)
 This is different from anions.
 You must draw the energy level diagram for the corresponding neutral atom first,
and then remove the number of electrons (corresponding to the charge) from the
orbitals with the highest principal quantum number, n.
Note: This may not be the highest –energy electron (See p.190)
ELECTRON CONFIGURATION
The order of increasing energy, the orbital sequence.
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p……
↑
(4s has less energy than 3d.)
Valence electrons e.g. Group 1 – valence of one
(Oxidation Number Group 2 – valence of two
PERIODICITY OF OUTER ENERGY-LEVEL ELECTRON CONFIGURATION

Atoms of all elements in a given group or family of the Periodic table have the
same electron population and orbital configuration in their outer energy levels.

Each group is composed of elements having similar properties but each group has
properties, which differ from those of the other groups.
SUMMARY
GROUP A
ALKAL1 METALS
2s1
Lithium
Li [He]
Sodium
Na [Ne] 3s1
Potassium
k [Ar]
Rubidium
Rb [Kr] 5s1
Cesium
Cs [Xe] 6s1
GROUP II A
4s1
VERY REACTIVE METALS
ALKALINE – EARTH METALS
Beryllium
Be [He] 2s2
Magnesium
Mg [Ne] 3s2
Calcium
Ca [Ar] 4s2
Strontium
Sr [Kr] 5s2
GROUP
III B
IV B
VB
VI B
VII B
VIII B
Transition Elements filling the d orbital block
IB
IIB
GROUP II A
Boron
B [He] 2s2 2p1
Aluminum
Al [Ne] 3s2 3p1
GROUP IV A
Carbon
C[He] 2s2 2p2
Silicon
Si [Ne] 3s2 3p2
GROUP V A
Nitrogen
N[He] 2s2 2p3
Phosphorus
p [Ne] 3s2 3p3
GROUP VI A
P-Orbital block
Oxygen
[He] 2s2 2p4
Sulphur
[Ne] 3s2 2p4
GROUP V11 A
Fluorine
[He] 2s2 2p5
Chlorine
[Ne] 3s2 3p5
Bromine
[Ar] 4s2 4p5
Iodine
[Kr] 5s2 5p5
GROUP O
NOBLE GAS (Highly stable)
Helium
1s2
Neon
2s3 2p6
Argon
3s2 3p6
Krypton
4s2 4p6
Xenon
5s2 5p6
5s2 6p6
Radon
f-Orbital block
n
=4
n
=5
n
=6
n=4
4s
3d
n=5
4p
5s
4d
n=6
5p
6s
4f
n =7
5d
6p
7s
5f
Lanthanide series
Actinide series
Inner Transition Elements
They are in the f-orbital block
IONIC CHARGE FORMATION
The electrons in the highest or outermost energy level are the ones generally involved in
chemical reactions.
Valence electrons e.g. Group 1 – valence of one (Oxidation Number)
Group 2 – valence of two
The formation of ionic charges can be explained with the following:

Zinc ion: The short hand form of the electron configuration for the atom is
[Ar] 4s2 3d10. This shows 12 outer electrons.
 In forming the ion of Zn2+, the two electrons are removed from the 4s rather than
the 3d10 which is more stable because of the electron clouds.
 Then, a relatively stable state, like atom with filled sub-shell is: Zn2+: [Ar] 3d10
ANOMALOUS ELECTRON CONFIGURATIONS
The rules for electron Configuration as described above may not work for all of the
elements. This is to ensure stability.
Following the rules, we would expect the following configurations:
Cr: [Ar] 4s2 3d4
Cu: [Ar] 4s2 3d9
However, the actual electron configurations, determined experimentally, are:
Cr:
[Ar] 4s1 3d5
Cu: [Ar] 4s1 3d10
Corresponding diagrams would be:
Cr: [Ar]
Predicted
(↑↓) (↑) (↓) (↑) (↓) ( )
Cu: [Ar} (↑↓) (↑↓) (↑↓) (↑↓) (↑↓) (↑)
4s
3d
Actual
Cr: [Ar] (↑)
Cu: [Ar]
(↑)
4s
(↑) (↑) (↑) (↑) (↑)
(↑↓) (↑↓) (↑↓) (↑↓) (↑↓)
3d
Note:
 For chromium, an electron is “borrowed” from the 4s sub-shell to give to 3d subshell that is half-filled
 For Copper, same arrangement.
 A similar thing happens with silver and gold
Ag: [Kr] 5s1 4d10
Au: [Xe] 6s1 5d10 4f14
 The borrowing of electrons also occurs in some of the ions formed by the
transition elements.
3.7: Wave Mechanics and orbitals:

The German physicist, Werner Heisenberg described the limitations of our ability
to measure both a particle’s velocity and its position at the same instant.
 This is called Heisenberg’ uncertainty principle. This is particularly true for small
particles such as electrons.
This led to two ideas:
Electron Cloud:
 Electron behaves as if it were spread out around the nucleus. This relates
probability to amplitude or intensity

Electron Probability Density:
 This relates to how much of electron’s charge is packed into a given volume.
Because of the wave nature, the electron (and its charge) is spread out around the nucleus.
 The electron probability density for a 1s orbital of hydrogen atom shows a
spherical shape. All of the s-orbitals are spherical.
 Their sizes increase with increasing size of n.
 A p-orbitals consists of three orbitals whose directions lie at 900 to each other
along the axes of an imaginary xyz coordinate system.
 The p-orbital concentrated along the x-axis is called px , and so forth. Their
size also increases with increasing size of n.