The electronic configuration based on orbitals is one concept which is deceptively easy to understand yet tricky enough to force students to make mistake. It is without a doubt, one of the most important concepts that you should build your chemistry foundation upon. So fret not, because Berry Berry Easy is back with Part 2 and final part of this two-part series on “The Electronic Structure of Atoms – STPM Chemistry Form 6“. All essential concepts regarding the electronic configuration, the structure will all be summarised here. However, you should draw all the structures instead of just memorising the concepts. The ability to draw the configurations is a hallmark of true understand of this topic. Hence, the Berry Berry Teacher will recommend all students to do more exercise on this and draw as much as possible.
Other parts covered in this topic includes also the atomic orbital, Heisen uncertainty principle, the types of orbital, core shell, valence shell, effective nuclear charge, ionisation energy, arrangment of electrons in an atom and also classification into the Periodic table. Combine what you have learnt in Part 1 with this part and you’ll fully understand the electronic structure of atoms.
STPM Chemistry Form 6 – Terminology and Concepts: The Electronic Structure of Atoms (Part 2 – Final)
Atomic orbital – the region, or volume, of space in an atom within the high probability (95% chance) of finding an electron in an atom.
Neil Bohr developed the model of atomic structure assuming that the electrons in an atom are in constant motion around the nucleus in circular orbits.
Heisenberg uncertainty principle states the position and momentum (mass x velocity) of an electron cannot be known with great precision. Charged particles in motion create magnetic fields, therefore it is possible to learn about the pathway and position of the moving electron.
Types of orbital: s, p, d and f orbital
Core shell – first shell that holds two electrons
Valence shell – the outermost shell
Effective nuclear charge (Nuclear attraction) accounts for (increases from left to right of the periodic table):
- i) attraction to the nucleus
- ii) repulsion from core electrons
- iii) minimal repulsion by other valence electrons
Ionisation energy is influenced by:
- i) nuclear charge (nuclear charge increases, the force of attraction on the electrons becomes stronger and the ionisation energy increases.)
- ii) distance of the electrons from the nucleus (further the outer electrons are from the nucleus, ionisation energies will be lower.)
- iii) screening effect (outermost electrons in an atom are shielded from the attraction of the nucleus by the repelling effect of the inner effect. The higher the screening effect, lower the ionization energies.)
Number of electrons in shell = 2(n)2
- Example: Lithium atom.
Nucleus = made up of both neutrons and protons
Core shell = 1st energy level (electron occupancy of 2)
Valence shell = 2nd energy level (electron occupancy up to 8 )
Arrangement of electrons in an atom
- Aufbau principle – Electrons occupy orbitals with the lowest energy level first
- Pauli exclusion principle – Each orbital can hold maximum of two electrons with opposite spin
- Hund’s rule – Orbital with the same energy level (degenerate orbitals), electron will occupy different orbital singly/one electron first the parallel spin, before pairing
The electrons are filled according the orbitals (Aufbau principle).
- Fills the 1s orbital to: 1s2
- Fills the 2s orbital to: 2s2
- Fills the 2p then 3s orbitals to: 2p63s2
- Fills the 3p then 4s orbitals to: 3p64s2
- Fills the 3d, 4p then 5s orbitals to: 3d104p65s2
(Extended knowledge: the process is repeated until all of the electrons have been accounted for. g-, h- and j-orbital exist in theory but the periodic table contains no elements that have electrons in either g-, h- and j-orbitals.)
The first break from numerical sequencing comes when the 4s level is filled before the 3d level, despite the fact that the perimeter of the 3d level is closer to the nucleus than the 4s orbital. The reason is that the energy of the level is based on an average position of the electron, not the extreme position.
Ionising electrons are not removed from the atom in reverse order! However, the outer shell electrons are always removed first when forming cations.
Example 1: Electronic configuration for manganese.
-> Solution 1:
- Neutral manganese (Mn) atom must contain 25 electrons.
- Electronic configuration of Mn: 1s22s22p63s23p64s23d5
Example 2: Which column of the periodic table is diamagnetic?
-> Solution 2:
- Column 2 (alkaline earth metals) and Column 8 (noble gas). A diamagnetic compound has its entire electron spin-paired. There must be an even number of electrons in the element. Valence electronic configuration for alkaline earth metals is ns2. Valence electronic configuration for noble gas is ns2np6.
- Column 1 (alkali metals) and Column 7 (halogen) are not diamagnetic.
- Column 6 (chalcogen) are paramagnetic. Valence electronic configuration for chalcogen is ns2np4.
Example 3: Electronic configuration for chromium
-> Solution 3:
- Half-filled d-shell stability in chromium: 1s22s22p63s23p64s13d5 rather than 1s22s22p63s23p64s23d4. (Others half-filled d-shell element are molybdenum and tungsten)
Example 4: Electronic configuration for copper.
-> Solution 4:
- Fully-filled d-shell stability in copper: 1s22s22p63s23p64s13d10 rather than 1s22s22p63s23p64s23d9.
Example 5: Which of the following electronic configuration represents an exited state?
A. He: 1s2
B. Li: 1s22p1
C. N: 1s22s22p3
D. F: 1s22s22p6
-> Solution 5:
- B (Li should have 1s22s1 as a ground state and the electronic configuration has the last electron in a 2p-orbital that is higher energy than the ground state 2s.)
- An excited state electronic configuration does not follow energetic sequence. An excited state has at least one electron in an energy level higher than the ground state.
Important: Not to confuse an ion (either cation or anion) with an excited state. A cation is an atom that has a deficit of at least one electron and thus carries a positive charge. An anion is an atom that has an excess of at least one electron and thus carries a negative charge.
Periodic Table can be classified into 4 main groups.
1) The s-block elements:
- Group 1 – general electronic configuration ns1.
- Group 2 – ns2.
2) The p-block elements
- Group 13 – ns2np1.
- Group 14 – ns2 np2.
- Group 15 – ns2 np3.
- Group 16 – ns2 np4.
- Group 17 – ns2 np5.
- Group 18 – ns2 np6.
3) The d-block elements
- Between Group 2 and Group 13 that the d orbitals are partially occupied.
4) The f-block elements
- Lanthanides (15 elements) – 4f orbitals are partially filled and must have a 6s2.
- Actinides (15 elements) – 5f orbitals are partially filled and must have a 7s2.