Orbitals and electronic configurationEdexcel International A Level Chemistry: Revision notes
Section 1
Shells, sub-shells and orbitals
Electrons occupy shells (quantum shells, n = 1, 2, 3, 4), which are divided into sub-shells (s, p, d, f), which contain orbitals.
An orbital is a region within an atom that can hold up to two electrons, which must have opposite spins.
- s sub-shell: 1 orbital, 2 electrons
- p sub-shell: 3 orbitals, 6 electrons
- d sub-shell: 5 orbitals, 10 electrons
Shell 1 has only 1s; shell 2 has 2s and 2p; shell 3 has 3s, 3p and 3d; shell 4 has 4s, 4p, 4d and 4f.
An orbital is not an orbit or path. It is a region of space where there is a high probability of finding an electron.
Section 2
Shapes of s and p orbitals
An s orbital is spherical, centred on the nucleus. The 2s orbital is larger than the 1s.
A p orbital is dumbbell-shaped, with two lobes either side of the nucleus. Each p sub-shell has three p orbitals, labelled px, py and pz, at right angles to each other along the x, y and z axes.
Section 3
Filling the orbitals
Electrons fill the lowest-energy orbitals first. The order for the first four shells is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
The 4s sub-shell fills before the 3d because it is slightly lower in energy.
In a sub-shell with several orbitals, each orbital takes a single electron before any pair up (Hund's rule). The pairing of electrons in one orbital causes repulsion, so pairing only happens when every orbital already has one electron.
In electrons-in-boxes notation each orbital is a box and each electron an arrow; paired electrons have arrows in opposite directions.
Section 4
Writing configurations: hydrogen to krypton
Work through the order of filling, using the number of electrons:
- Oxygen (8): 1s²2s²2p⁴
- Chlorine (17): 1s²2s²2p⁶3s²3p⁵
- Calcium (20): 1s²2s²2p⁶3s²3p⁶4s²
- Iron (26): 1s²2s²2p⁶3s²3p⁶3d⁶4s²
Two exceptions in the first four periods are chromium (24) and copper (29), which are 3d⁵4s¹ and 3d¹⁰4s¹ rather than 3d⁴4s² and 3d⁹4s².
Check your answer by adding the superscripts: they must equal the atomic number for an atom.
Section 5
Configurations of ions
Positive ions form by removing the highest-energy electrons. For the d-block elements, the 4s electrons are removed first, even though 4s fills before 3d.
- Fe: 3d⁶4s². Fe²⁺: 1s²2s²2p⁶3s²3p⁶3d⁶. Fe³⁺: …3d⁵.
- Zn²⁺: …3d¹⁰.
Negative ions add electrons to the outer sub-shell: Cl⁻ is 1s²2s²2p⁶3s²3p⁶.
Do not take electrons from the 3d sub-shell of a transition metal before the 4s.
Section 6
Configuration and chemical properties
Chemical properties depend on the electronic configuration, especially the outer electrons.
- Elements in the same group have the same number of outer electrons, so they react in similar ways.
- A full outer shell, as in neon, makes an element unreactive.
- One outer electron, as in sodium, is easily lost, forming a 1+ ion.
- One electron short of a full shell, as in chlorine, gains an electron to form a 1− ion.
That's the notes covered.
Carry on to the next subtopic.
Exam questions on Orbitals and electronic configuration
- A student is learning to write the electronic configurations of atoms and ions of elements in the first four periods.State what is meant by the term orbital.2 marks
- Iron (atomic number 26) is the most widely used transition metal. Iron(II) and iron(III) ions are found in haemoglobin and in rust.Describe how the six 3d electrons of an Fe²⁺ ion are arranged in the 3d orbitals, and explain why they are arranged in this way.2 marks
- A transition metal course covers elements from scandium to zinc. The atomic numbers are: vanadium 23, manganese 25 and zinc 30.Write the full electronic configuration, using s, p, d notation, of (i) a vanadium atom, (ii) a V³⁺ ion and (iii) a Zn²⁺ ion.3 marks
Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).