Atomic Structure & the Periodic Table

Inside the Atom

Everything around you, from the air you breathe to the screen you are reading, is made of atoms. Atoms are incredibly small, but they are not the end of the story: each atom is built from three even smaller subatomic particles — protons, neutrons and electrons.

The centre of an atom is a tiny, dense nucleus containing the protons and neutrons. Around the nucleus, the electrons move in regions called shells (or energy levels). Almost all of the atom is empty space.

Key terms

Nucleus — the central core of an atom, containing protons and neutrons.

Proton — a positively charged particle found in the nucleus.

Neutron — an uncharged (neutral) particle found in the nucleus.

Electron — a negatively charged particle found in shells around the nucleus.

Relative Mass and Charge

We measure the masses and charges of these particles relative to each other, rather than in tiny real numbers. Learn this table — it is asked about constantly.

An electron is so light its mass is treated as negligible. Because an atom has equal numbers of protons and electrons, the positive and negative charges cancel out, so a complete atom is electrically neutral.

Atomic Number and Mass Number

Two numbers describe every atom, and they are written next to the element symbol like this: $^{A}_{Z}\text{X}$.

Atomic number (proton number), $Z$ — the number of protons in the nucleus. This defines the element. Every carbon atom has 6 protons; if it had 7, it would be nitrogen.
Mass number, $A$ — the total number of protons and neutrons in the nucleus.

So: number of neutrons = mass number − atomic number.

For sodium, $^{23}_{11}\text{Na}$: 11 protons, 11 electrons, and $23 - 11 = 12$ neutrons.

Exam tip

In a neutral atom, electrons = protons = atomic number. Find neutrons by subtracting the atomic number from the mass number. Mark schemes love this calculation, so always show your subtraction.

Isotopes

Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons. Because the proton number is unchanged, isotopes have the same atomic number but different mass numbers.

For example, chlorine has two common isotopes:

$^{35}_{17}\text{Cl}$ — 17 protons, 18 neutrons
$^{37}_{17}\text{Cl}$ — 17 protons, 20 neutrons

Isotopes of an element have identical chemical properties. This is because chemical reactions involve electrons, and isotopes have the same number of electrons arranged in the same way. They differ only in physical properties such as density, because the extra neutrons make some isotopes slightly heavier.

Calculating Relative Atomic Mass

Because most elements are mixtures of isotopes, we use a weighted average called the relative atomic mass ($A_r$) — the average mass of an atom compared with $\tfrac{1}{12}$ the mass of a carbon-12 atom. We weight the average by how common each isotope is (its relative abundance, given as a percentage).

$$A_r = \frac{\sum(\text{isotope mass} \times \text{\% abundance})}{100}$$

Worked example

A sample of chlorine contains 75% of the $^{35}\text{Cl}$ isotope and 25% of the $^{37}\text{Cl}$ isotope. Calculate the relative atomic mass of chlorine.

Multiply each mass by its abundance and add:

$(35 \times 75) + (37 \times 25) = 2625 + 925 = 3550$

Electronic Configuration

Electrons occupy shells around the nucleus, filling the inner shells first. At IGCSE you use the 2, 8, 8 rule:

1. The first shell holds a maximum of 2 electrons.
2. The second shell holds a maximum of 8 electrons.
3. The third shell holds up to 8 electrons (for the first 20 elements).

We write the configuration as numbers separated by commas, starting from the inner shell. For example:

Carbon (6 electrons): 2, 4
Sodium (11 electrons): 2, 8, 1
Argon (18 electrons): 2, 8, 8
Calcium (20 electrons): 2, 8, 8, 2

To draw a shell diagram, put the symbol in the centre, draw a circle for each shell, and add the correct number of crosses or dots, filling inner shells first.

From Plum Pudding to the Nuclear Atom

Our picture of the atom developed over time. Around 1900, J. J. Thomson proposed the plum pudding model: a ball of positive charge with electrons dotted through it like fruit in a pudding.

In 1909, Rutherford's team fired positively charged alpha particles at thin gold foil. Most passed straight through, but a few bounced back sharply. This could only happen if the positive charge and mass were concentrated in a tiny central nucleus, with electrons orbiting in mostly empty space. This nuclear model replaced the plum pudding model and is the basis of what we use today.

Real world

Rutherford's gold foil experiment is a classic example of how scientists revise models when new evidence appears. A surprising result — alpha particles bouncing back — overturned the accepted theory. Good science follows the evidence.

The Modern Periodic Table

The periodic table arranges elements in order of increasing atomic number. Its layout is not random — it groups elements with similar behaviour together.

Periods are the horizontal rows. The period number tells you the number of occupied electron shells.
Groups are the vertical columns. The group number tells you the number of electrons in the outer shell (for the main groups, 1 to 7, and 0 for the noble gases).

A zig-zag line (a "staircase") separates metals on the left from non-metals on the upper right. Most elements are metals.

Linking Configuration to the Table

Electronic configuration and table position are directly connected:

Number of outer-shell electrons = group number. Sodium is 2, 8, 1 → one outer electron → Group 1.
Number of occupied shells = period number. Sodium has 3 shells → Period 3.

So just from "2, 8, 1" you can place sodium exactly: Group 1, Period 3.

Exam tip

If asked which group and period an element is in, write its electronic configuration first. The last number is the group; the count of numbers is the period. Noble gases (full outer shells) are placed in Group 0.

Why a Group Shares Properties

Elements in the same group have the same number of outer-shell electrons. Since chemical reactions depend on these outer electrons, elements in a group react in similar ways and form similar compounds.

For example, the Group 1 metals (lithium, sodium, potassium) each have one outer electron. They all react with water in the same way to form a hydroxide and hydrogen gas. The noble gases (Group 0) have full outer shells, which makes them stable and very unreactive.

Watch out

"Same group = same number of outer electrons" is the reason for similar chemistry, not "same number of total electrons." Don't confuse the two — each element in a group still has a different total number of electrons.