Atomic Structure and Electron Configuration: Study Guide (Chemistry 1.5)

What is inside an atom

An atom has two parts:

• A tiny nucleus at the centre, holding protons (charge +1 each) and neutrons (no charge). Almost all of the mass is here.
• Electrons (charge −1 each) around the nucleus, filling nearly all of the atom’s volume.

The number of protons is the atomic number, Z. It decides which element the atom is. A neutral atom has as many electrons as protons. An ion has a different number of electrons: a cation (positive ion) has lost electrons, and an anion (negative ion) has gained them. The number of protons never changes when an ion forms.

Opposite charges attract. That attraction holds the atom together, and it explains almost everything in this topic.

Coulomb’s law: how strongly charges attract

Coulomb’s law describes the force between two charged particles:

F ∝ q₁q₂ / r²

• q₁ and q₂ are the two charges. For an electron and a nucleus, the product is negative, which means attraction. When you compare strengths, use the size of the product.
• r is the distance between the two charges.

You will use this law as a proportion, not to calculate a force in newtons. Two patterns matter:

• Charge. Double one of the charges and the force doubles.
• Distance. The force depends on 1/r². Double the distance and the force falls to ¼. Triple it and the force falls to 1/9.

For an electron in an atom, this means: the more positive charge it feels from the nucleus, and the closer it is, the more strongly it is held.

Shells and subshells

Electrons in an atom do not sit at random distances. They occupy shells (energy levels), numbered n = 1, 2, 3 and so on. Higher n means, on average, farther from the nucleus and higher in energy.

Each shell is split into subshells (sublevels), labelled s, p, d and f. Each type of subshell holds a fixed maximum number of electrons:

Subshell	Maximum electrons	First appears in shell
s	2	n = 1
p	6	n = 2
d	10	n = 3
f	14	n = 4

So shell 1 has only 1s; shell 2 has 2s and 2p; shell 3 has 3s, 3p and 3d.

A ground-state electron configuration lists the occupied subshells in order, with a superscript for the number of electrons in each. Nitrogen (7 electrons) is 1s²2s²2p³. The superscripts always add up to the number of electrons.

The filling order: the Aufbau principle

In the ground state, electrons fill the lowest-energy subshells first. This is the Aufbau principle (“Aufbau” is German for “building up”). The order is:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p

Note that 4s fills before 3d: when potassium and calcium are reached, 4s is the lower-energy subshell, so it takes electrons first.

You do not need to memorise this order: the periodic table is built from it.

To write a configuration from the table, start at hydrogen and “walk” along each period until you reach your element, writing down each block you cross. For example, iron is in period 4, sixth column of the d block. The walk gives 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶.

Box diagrams (useful background)

A box diagram shows each orbital as a box. An s subshell has one box, p has three, d has five. Two rules decide how arrows go in:

• Each box holds at most two electrons, drawn with opposite arrows (opposite spins).
• In a subshell with several boxes, electrons go one per box first, all with the same spin, before any box gets a second electron.

The course exam does not ask you to assign quantum numbers to electrons, so you can treat these boxes as a picture, not a set of numbers.

Noble-gas shorthand

Replace the part that matches the previous noble gas with its symbol in square brackets:

• Na: 1s²2s²2p⁶3s¹ = [Ne]3s¹
• Fe: 1s²2s²2p⁶3s²3p⁶4s²3d⁶ = [Ar]4s²3d⁶

Some books write [Ar]3d⁶4s², ordering by shell number. Both are accepted; they describe the same electrons.

Core and valence electrons

• Core electrons are the inner electrons: the ones in the noble-gas core.
• Valence electrons are the outer electrons, in the highest occupied shell. For main-group elements these are the s and p electrons of the largest n.

Chlorine, [Ne]3s²3p⁵, has 10 core electrons and 7 valence electrons (3s²3p⁵). Valence electrons take part in bonding.

Configurations of ions

Anions gain electrons. Add them to the next available subshell. Chlorine gains one electron: Cl⁻ is 1s²2s²2p⁶3s²3p⁶ = [Ar]. Sulfur gains two: S²⁻ is also [Ar]. These ions have the same configuration as argon (they are isoelectronic with it).

Cations lose electrons. Remove them from the highest shell (largest n) first, and from p before s within that shell. Calcium, [Ar]4s², loses both 4s electrons: Ca²⁺ is [Ar].

For transition metals this rule has a twist. The 4s subshell fills before 3d, but once 3d holds electrons, the 4s electrons are the outermost ones (n = 4 is larger than n = 3). So 4s electrons are removed first. Iron, [Ar]4s²3d⁶, becomes Fe²⁺ = [Ar]3d⁶, not [Ar]4s²3d⁴.

A note on exceptions: a few elements, such as chromium ([Ar]4s¹3d⁵) and copper ([Ar]4s¹3d¹⁰), do not follow the simple filling order. Writing configurations for these exceptions is not assessed on the exam.

How hard is an electron to remove?

The energy needed to remove an electron from an atom or ion is its ionization energy. You can estimate which electrons are harder to remove with Coulomb’s law, using two factors:

1. Distance (r). Electrons in a higher shell are farther from the nucleus, so they are held less strongly.

2. Effective nuclear charge. An outer electron does not feel the full charge of the nucleus, because the inner (core) electrons between them shield it. A simple estimate is the core charge:

   core charge ≈ protons − core electrons

   For sodium’s 3s electron: 11 − 10 = +1. For magnesium’s 3s electrons: 12 − 10 = +2.

The core charge is a rough model: shielding is not perfect. But it gets the direction of comparisons right, which is what this topic needs.

Use the two factors for two kinds of comparison:

• Different subshells in the same atom. Inner electrons are closer and feel more of the nucleus. Removing a 2p electron from sodium takes far more energy than removing its 3s electron.
• The same subshell in different atoms. Compare the charge felt and the distance. Mg’s 3s electrons feel about +2, Na’s about +1, at a similar distance, so Mg’s 3s electrons are harder to remove.

Worked example 1: cobalt, its ions and chloride

Question. Write the full and noble-gas ground-state configurations for a cobalt atom (Z = 27), then for Co²⁺ and Co³⁺. Then write the configuration of Cl⁻ (Z = 17).

1. Locate cobalt: period 4, seventh column of the d block. Walking the table gives 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁷.
2. Check the count: 2 + 2 + 6 + 2 + 6 + 2 + 7 = 27. Correct.
3. Shorthand: the first 18 electrons match argon, so Co is [Ar]4s²3d⁷.
4. Co²⁺ has 25 electrons. Remove two electrons from the highest shell, n = 4. Both 4s electrons go: Co²⁺ = [Ar]3d⁷.
5. Co³⁺ has 24 electrons. The 4s subshell is already empty, so the third electron comes from 3d: Co³⁺ = [Ar]3d⁶.
6. Cl has 17 electrons, [Ne]3s²3p⁵. Cl⁻ gains one electron into 3p: 1s²2s²2p⁶3s²3p⁶ = [Ar].

Check. The superscripts add to 27, 25, 24 and 18. The proton numbers (27 and 17) never change.

Common slip. [Ar]4s²3d⁵ for Co²⁺ removes the “last electrons in”, not the outermost ones. Its electron count is right, so counting alone will not catch it.

Worked example 2: ranking electrons with Coulomb’s law

Question. Rank these four electrons from hardest to easiest to remove, and justify the ranking:

• (i) a 3s electron in Na
• (ii) a 3s electron in Mg
• (iii) a 2p electron in Na
• (iv) the 4s electron in K

Step 1: estimate the core charge each electron feels.

Electron	Protons	Electrons in inner shells	Core charge
(i) Na 3s	11	10	+1
(ii) Mg 3s	12	10	+2
(iii) Na 2p	11	2	+9
(iv) K 4s	19	18	+1

Step 2: compare distances. The 2p electron is in shell 2, much closer than shell 3. The K 4s electron is in shell 4, farther than shell 3.

Step 3: apply F ∝ q₁q₂ / r².

• (ii) against (i): same shell, similar distance, but twice the core charge. At the same r the force is 2 times larger. (Mg’s 3s electrons are also a little closer, which makes the difference bigger still.)
• (iii) against (i): larger charge and smaller distance. In a rough model where the 2p electron is half as far away, the force is 9 × 2² = 36 times larger. The real number differs, but the direction is clear.
• (iv) against (i): same core charge, greater distance, so a weaker force.

Answer. Hardest to easiest: (iii) Na 2p > (ii) Mg 3s > (i) Na 3s > (iv) K 4s.

Tip. In a written answer, name both factors, charge and distance, and say which way each one pushes.

Common misconceptions

• “The 3d electrons are lost first because they went in last.” Cations lose electrons from the highest shell first. For first-row transition metals that is 4s, even though 4s filled earlier.
• “An outer electron feels the whole nuclear charge.” Core electrons shield it. Sodium’s 3s electron feels roughly +1, not +11.
• “More protons always means electrons are held more tightly.” Distance matters too. Potassium has more protons than sodium, but its outer electron is in a farther shell with the same core charge, so it is easier to remove.
• “Doubling the distance halves the force.” The force depends on 1/r², so doubling the distance gives one quarter of the force.
• “Ions change the number of protons.” Only electrons are gained or lost. Fe³⁺ still has 26 protons.
• “Isoelectronic species are the same element.” Cl⁻, Ar and K⁺ share a configuration but have different numbers of protons.

Where this leads

Next, Topic 1.6, Photoelectron Spectroscopy, gives experimental evidence for shells and subshells: each peak matches a subshell, and its energy follows the Coulomb’s law reasoning of Worked example 2. Later topics use the same ideas for periodic trends and bonding. If you need a refresher on the previous topic, see Topic 1.4, Composition of Mixtures.

Try the practice questions now, then use the revision notes and the checklist to consolidate.