Atomic Structure

⚛️ Atomic Structure Revision Notes: Definitions, Particles and Worked Examples

Atomic structure: the definition

An atom is the smallest particle of an element that still has the chemical properties of that element. Atomic structure is the way an atom is built: a tiny, dense, positively charged nucleus of protons and neutrons at the centre, surrounded by negatively charged electrons arranged in shells (energy levels). Atoms are around 0.1 nanometres across — a ten-millionth of a millimetre — and the nucleus is roughly ten thousand times smaller than the atom itself, so almost all of an atom is empty space.

The three sub-atomic particles

Because protons and neutrons each have a relative mass of 1 and electrons weigh almost nothing, nearly all of an atom's mass is in the nucleus. Because the proton's charge (+1) and the electron's charge (−1) cancel, an atom with equal numbers of the two has no overall charge. Neutrons carry no charge; their job is to add mass and hold the nucleus together.

Atomic number and mass number

Every element has two numbers on the periodic table. The atomic number (symbol Z) is the number of protons; it is what makes an element that element, so every carbon atom has 6 protons and every oxygen atom has 8. The mass number (symbol A) is the number of protons plus neutrons. From those two you can work out everything:

• Protons = atomic number
• Electrons = atomic number (in a neutral atom, protons and electrons are equal)
• Neutrons = mass number − atomic number

Worked example 1 — sodium, ²³Na, atomic number 11. Protons = 11. Electrons = 11. Neutrons = 23 − 11 = 12.

Worked example 2 — chlorine, ³⁵Cl, atomic number 17. Protons = 17. Electrons = 17. Neutrons = 35 − 17 = 18.

Worked example 3 — an ion. A magnesium ion, Mg²⁺, has atomic number 12, so it still has 12 protons; the 2+ charge tells you it has lost two electrons, so it has 12 − 2 = 10 electrons. Losing or gaining electrons makes an ion; the number of protons never changes.

Electron shells and electron configuration

Electrons occupy shells at increasing distances from the nucleus, and each shell fills before the next one starts. For the first twenty elements the rule is: the first shell holds up to 2 electrons, the second up to 8, the third up to 8, and any remaining electrons start the fourth shell. Writing the number in each shell, separated by commas, gives the electron configuration.

• Oxygen, 8 electrons: 2, 6
• Sodium, 11 electrons: 2, 8, 1
• Argon, 18 electrons: 2, 8, 8 (a full outer shell, so it is unreactive)
• Calcium, 20 electrons: 2, 8, 8, 2

Two shortcuts save time in exams. The number of shells equals the period (row) the element is in, and the number of electrons in the outer shell equals its group number for Groups 1 to 7. Sodium (2, 8, 1) is in Period 3, Group 1; chlorine (2, 8, 7) is in Period 3, Group 7. It is the outer electrons that take part in reactions, which is why elements in the same group behave alike.

Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They therefore share an atomic number but have different mass numbers. Chlorine-35 and chlorine-37 both have 17 protons; one has 18 neutrons and the other 20. Because their electron configurations are identical, isotopes react in the same way — only their mass differs. Carbon has three natural isotopes: carbon-12, carbon-13 and the radioactive carbon-14 used in dating ancient objects.

Relative atomic mass

Because most elements are a mixture of isotopes, the mass shown on the periodic table is an average. The relative atomic mass (A_r) is the weighted average mass of an element's atoms compared with one twelfth of the mass of a carbon-12 atom. To calculate it, multiply each isotope's mass number by its percentage abundance, add the results, and divide by 100.

Worked example — chlorine. Natural chlorine is about 75% chlorine-35 and 25% chlorine-37. A_r = (35 × 75 + 37 × 25) ÷ 100 = (2625 + 925) ÷ 100 = 35.5. That is why chlorine's mass on the periodic table is not a whole number even though every individual atom has a whole-number mass.

Quick recap for the activities

• Protons: mass 1, charge +1, nucleus. Neutrons: mass 1, charge 0, nucleus. Electrons: negligible mass, charge −1, shells.
• Atomic number = protons. Mass number = protons + neutrons. Neutral atom: electrons = protons.
• Shells fill 2, 8, 8, then 2 for the first twenty elements.
• Isotopes: same protons, different neutrons, same chemistry.
• Relative atomic mass is the abundance-weighted average of the isotope masses.

Key terms

Atom

The smallest particle of an element that keeps the properties of that element.

Nucleus

The tiny, dense centre of an atom, containing the protons and neutrons.

Atomic number

The number of protons in an atom; it identifies the element.

Mass number

The total number of protons and neutrons in an atom.

Electron configuration

The arrangement of electrons in shells, written as numbers such as 2, 8, 1.

Ion

An atom that has gained or lost electrons and so carries a charge.

Isotope

An atom of an element with the usual number of protons but a different number of neutrons.

Relative atomic mass

The average mass of an element's atoms, weighted by isotope abundance, on a scale where carbon-12 = 12.