Atomic structure

Atoms and Molecules

An atom is the smallest particle of an element that can take part in a chemical reaction. Atoms cannot be broken down into simpler substances by chemical means. Every element on the periodic table is made up of one type of atom, and the type of atom defines the element.

A molecule is a group of two or more atoms chemically bonded together. These atoms can be of the same element or of different elements.

For example, O₂ is a molecule containing two oxygen atoms bonded together, while H₂O is a molecule containing two hydrogen atoms and one oxygen atom bonded together. When a molecule contains atoms of different elements, the substance is called a compound.

MisconceptionStudents often confuse atoms and molecules. A single oxygen atom (O) is an atom; two oxygen atoms bonded together (O₂) form a molecule. An atom is always a single particle, never a bonded group.

Exam TipIf asked to distinguish between an atom and a molecule, stress that a molecule contains two or more atoms bonded together.

Structure of the Atom

The structure of an atom consists of a central nucleus surrounded by electrons arranged in energy levels (shells). The nucleus is extremely small compared to the overall size of the atom, yet it contains almost all of the atom’s mass.

The nucleus contains two types of sub-atomic particle: protons and neutrons. Electrons orbit the nucleus in shells at different distances from the centre.

Protons carry a relative charge of +1, electrons carry a relative charge of −1, and neutrons carry no charge (0). Because an atom has equal numbers of protons and electrons, the positive and negative charges cancel out, so atoms are electrically neutral overall.

The relative masses and charges of the three sub-atomic particles are summarised below.

Reading the table: Each row describes one sub-atomic particle. “Relative mass” and “relative charge” compare particles to each other rather than giving absolute values. A relative mass of 1 is the baseline set by the proton.

MisconceptionStudents sometimes state that the electron has a relative mass of zero. The correct exam phrasing is “negligible” or “1/1836”. It has mass, but it is so tiny compared to a proton or neutron that it is effectively ignored in mass calculations.

Exam TipWrite “negligible” rather than “zero” for the relative mass of an electron.

Atomic Number, Mass Number and Isotopes

The atomic number (also called proton number) is the number of protons in the nucleus of an atom. This number defines which element an atom belongs to, so all atoms of the same element have the same atomic number. In a neutral atom, the number of electrons equals the number of protons, because the overall charge must be zero.

The mass number is the total number of protons and neutrons in the nucleus of an atom. Electrons are not included because their mass is negligible. The number of neutrons in an atom can therefore be calculated:

neutrons = mass number − atomic number

Notation convention: When an element’s symbol is written with two numbers, the mass number appears as a superscript (top) and the atomic number as a subscript (bottom) to the left of the symbol. For example, ²³₁₁Na means sodium with mass number 23 and atomic number 11.

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Because isotopes have the same atomic number, they are the same element and have identical chemical properties. They differ only in mass number, because the number of neutrons varies.

For example, carbon-12 (¹²₆C) and carbon-14 (¹⁴₆C) are isotopes of carbon — both have 6 protons, but carbon-12 has 6 neutrons while carbon-14 has 8 neutrons.

Examiner InsightWhen asked to define isotopes, examiners require three components: same element (or same number of protons), different numbers of neutrons, and different mass numbers. Missing any one of these typically costs a mark.

Exam TipAlways state “same number of protons” rather than just “same element” to secure both marks.

MisconceptionIsotopes do not have different numbers of electrons. A common error is to confuse neutrons with electrons. Isotopes differ only in their neutron count, which changes the mass number but not the charge or chemical behaviour.

Exam TipChemical properties depend on electron arrangement, so isotopes of the same element react identically.

Relative Atomic Mass from Isotopic Abundances

Relative atomic mass (Aᵣ) is the weighted average mass of the atoms of an element, measured relative to one-twelfth the mass of a carbon-12 atom. Most elements exist as a mixture of isotopes in nature, so the Aᵣ accounts for the mass of each isotope and how common (abundant) it is in a natural sample.

Key Equations

Relative atomic mass:

$${A}_{r}=\frac{\text{(mass of isotope 1}\times \text{abundance)}+\text{(mass of isotope 2}\times \text{abundance)}+…}{100}$$

Variables:

• mass of isotope = mass number of each isotope
• abundance = percentage of each isotope in a natural sample
• 100 = total percentage

SI unit: Aᵣ has no unit (it is a ratio)

The calculation multiplies each isotope’s mass number by its percentage abundance, sums these products, and divides by 100. This gives a weighted average that reflects the contribution of each isotope to the overall mass.