Year 3 chemistry: the atom and chemical bonds

Mid-September 2026: in the Year 3 class of an Italian high school specialising in the human sciences (the liceo delle scienze umane), natural sciences has only just started, and the teacher is introducing (or about to introduce) the structure of the atom and the periodic table. The first test arrives in five to eight weeks, so between the end of October and mid-November. This is not yet the time for last-minute revision. It is the time to understand the topic properly, because from here on all of chemistry — bonds, naming, reactions — is built on this foundation.

What the test really asks at this point

At the human-sciences high school, natural sciences gets 2 hours a week in all five years. It is squeezed in between human sciences at 5 hours, philosophy (new, at 3 hours), Latin, which drops to 2 hours but stays, and history of art, which makes its debut in Year 3. With this workload, natural sciences is almost always the subject that gets the least attention, from revision at home to private tutoring, which almost nobody here offers specifically. Yet the grade counts like all the others, and it goes into the average and into the school credit (credito scolastico, the points that feed the final school-leaving mark), which starts precisely in the third year.

The end-of-October test almost always asks for three things together:

1. being able to write the electron configuration of an atom and place it in the periodic table (period and group);
2. being able to explain how and why some properties change along a period and down a group (not just list them);
3. being able to predict the type of bond between two elements and give a name to the resulting compound.

Why students who got through the first two years “by heart” get stuck here

In Years 1 and 2, natural sciences is taught through content that can be learned as a list of facts: states of matter and changes of state, what a mixture is and what a pure substance is, the laws of Lavoisier and Proust, balancing an equation. You can reach a 6 even by repeating definitions without really understanding them, because the test often asks you to recognise or apply a rule, not to connect causes and effects.

In Year 3 the register changes. Atomic models, electron configuration and periodic trends are not facts to list: they are reasoning you have to rebuild each time. “Why does sodium lose an electron more easily than chlorine?” is not answered by reciting a definition. You answer it by linking the structure of the atom (how many electrons it has in the outer level, how far it is from the nucleus) to the chemical property that follows. It is the first topic of the year where memory alone is no longer enough, and that is exactly why so many students who had a steady 6-7 in the first two years find themselves facing a failing grade in November without understanding what has changed.

The mistakes that cost marks

• Confusing period and group. The period (horizontal row) tells you how many energy levels the atom has occupied; the group (column) tells you how many electrons it has in the outer level (for the main-group elements). Swapping them leads you to get both the configuration and the prediction of properties wrong.
• Writing the electron configuration in the wrong order. The filling order is 1s, 2s, 2p, 3s, 3p, 4s… A frequent mistake is to fill the whole third level (3s and 3p) and then the fourth, skipping the real order of increasing energy, or to forget that each s orbital holds at most 2 electrons and each p sublevel holds at most 6.
• Reversing the periodic trends. Atomic radius increases going down a group (more levels occupied) and decreases going right along a period (more protons pull the electrons closer). Ionisation energy and electronegativity go in the opposite direction to the radius: they grow towards the right and towards the top. Mixing up “increases” and “decreases” on one of these three properties is by far the most common mistake in end-of-October tests.
• Deciding the type of bond without looking at the electronegativity difference. “Metal + non-metal, so ionic” is almost always true but it is not the criterion: the criterion is the electronegativity difference between the two atoms. Anyone who skips this step struggles when faced with a less familiar compound.
• Getting the suffixes of traditional nomenclature wrong. The suffix -oso indicates the lower valence of an element with variable valence, -ico the higher one. Swapping them turns one compound into another (ferrous oxide, FeO, and ferric oxide, Fe₂O₃, are two different substances, not the same one written in two ways).

The method, step by step

For electron configuration and position in the periodic table:

1. Count the electrons of the neutral atom: they equal the atomic number Z.
2. Fill the levels in the order 1s, 2s, 2p, 3s, 3p, 4s, distributing the electrons while respecting the maximum capacity of each sublevel (2 for s, 6 for p).
3. The period is the number of the outermost occupied energy level.
4. The group (for the main-group elements) corresponds to the number of electrons in the outermost level.

To predict the type of bond:

1. Find the position of the two elements in the periodic table (metal, non-metal).
2. If possible, compare the electronegativity difference: above about 1.9 the bond is ionic; between 0.4 and 1.9 it is polar covalent; below 0.4, or between two atoms of the same element, it is pure (non-polar) covalent.
3. Write the formula respecting the overall electrical neutrality of the compound.

For basic nomenclature:

1. Recognise the class of the compound: oxide/hydroxide (with a metal), acid (with hydrogen and a non-metal or a polyatomic group), salt.
2. If the element has a fixed valence, the name is unambiguous: “sodium oxide”, “calcium hydroxide”.
3. If the element has a variable valence, use -oso for the lower valence and -ico for the higher one (traditional nomenclature), or the Roman numeral in brackets (IUPAC nomenclature): iron(II) and iron(III).

Three worked examples

1. Electron configuration and position in the periodic table

A neutral atom has atomic number $Z=15$ (phosphorus). Write its electron configuration and give its period and group.

I fill the levels in order, for a total of 15 electrons:

Check: $2+2+6+2+3=15$, correct. The outermost occupied level is the third ($n=3$): phosphorus is in period 3. The electrons in the outermost level are $2+3=5$: phosphorus is in group 15 (5 valence electrons), the same group as nitrogen.

2. Comparing periodic properties

Why does sodium (Z = 11) easily form the ion $\mathrm{N}{\mathrm{a}}^{+}$, while chlorine (Z = 17) easily forms the ion $\mathrm{C}{\mathrm{l}}^{-}$?

Sodium has the configuration $1s^{2}2s^{2}2p^{6}3s^1$: a single electron in the outermost level, far from the nucleus and held loosely. Its ionisation energy is low (about 496 kJ/mol): losing that electron and reaching the stable configuration of neon costs little energy.

Chlorine has the configuration $1s^{2}2s^{2}2p^{6}3s^{2}3p^5$: seven electrons in the outermost level, just one electron short of the stable configuration of argon. Chlorine is much further to the right in the period than sodium (a nucleus with more protons, a smaller radius), so it holds on to its electrons strongly and is happy to take one more: its electron affinity is high.

Put together, sodium “wants” to lose an electron and chlorine “wants” to gain one. This is exactly the pair of tendencies that leads to the ionic bond of $\mathrm{NaCl}$, the next exercise.

3. Type of bond in three compounds

Classify the bond in $\mathrm{NaCl}$, ${\mathrm{H}}_2\mathrm{O}$ and $\mathrm{C}{\mathrm{l}}_2$, knowing that the electronegativity (Pauling scale, approximate values) is: $\mathrm{Na}=0.9$; $\mathrm{Cl}=3.0$; $\mathrm{H}=2.1$; $\mathrm{O}=3.5$.

For $\mathrm{NaCl}$:

Above 1.9: ionic bond. This is consistent with the previous exercise: sodium gives its electron to chlorine, it does not share it.

For ${\mathrm{H}}_2\mathrm{O}$ (the O–H bond):

Between 0.4 and 1.9: polar covalent bond. Oxygen attracts the bonding electrons more strongly, but it does not pull them away completely as in the ionic case.

For $\mathrm{C}{\mathrm{l}}_2$ (two atoms of the same element):

Pure (non-polar) covalent bond: neither atom attracts the electrons more than the other, by definition, since they are the same element.

How I work on this topic in my lessons

With Year 3 human-sciences students I almost always start from the periodic table as an object to “read”, not to learn by heart. First the structure (periods, groups, blocks), then electron configuration as a consequence of that structure, and only at the end bonds and nomenclature, which by then become logical applications and not new rules to remember separately. The material on my science pages follows the same order. Electron configuration, which sits in the middle, has a post of its own: Configurazione elettronica e numeri quantici: il metodo.

To practise outside the lesson too, the site has the high-school test practice papers: for now they mainly cover maths and physics, and there is no chemistry practice test yet, but you will still find them useful for the method of working against the clock. If you are coming from the first two years and balancing the reactions of Year 2 is not yet solid, I have written a dedicated guide: balancing chemical equations in Year 2. Later in the programme you will get to the mole and stoichiometric calculations: when you do, the typical mistakes (and how to avoid them) are collected here: /blog/stechiometria-errori-tipici-liceo-medicina/.

If you want to set up the topic properly before the end-of-October test, or you already have a marked test to go through together, write to me: we can start with a single lesson or the 10-hour package (the costs are on the pricing page) if you want to follow the whole module continuously, plus a free 30-minute introductory call if you first just want to understand how I work. Payment is always offline — bank transfer, Satispay, PayPal, card, Apple Pay or Revolut.