Atomic Structure

Basic Atomic Structure

Atoms are the smallest units of matter that retain the properties of an element. They are made up of three main particles:

• Protons 6 positively charged particles found in the nucleus.
• Neutrons 6 neutral particles (no charge) also located in the nucleus.
• Electrons 6 negatively charged particles orbiting the nucleus in shells.

The nucleus is the tiny, dense centre of the atom containing protons and neutrons. Electrons move around the nucleus in fixed energy levels called electron shells.

Two important numbers describe an atom:

• Atomic number (Z): the number of protons in the nucleus. This defines the element.
• Mass number (A): the total number of protons and neutrons in the nucleus.

For example, a carbon atom has 6 protons (atomic number 6) and usually 6 neutrons, so its mass number is 12.

The number of electrons in a neutral atom equals the number of protons. Electrons are arranged in shells around the nucleus, with the first shell holding up to 2 electrons, the second up to 8, and so on.

For instance, a sodium atom has 11 protons, so it has 11 electrons arranged as 2 in the first shell, 8 in the second, and 1 in the third shell.

Isotopes

Isotopes are atoms of the same element that have the same number of protons (same atomic number) but a different number of neutrons, so they have different mass numbers.

Because isotopes have the same number of protons, they have the same chemical properties but may have different physical properties, such as stability or mass.

For example:

• Carbon-12: 6 protons and 6 neutrons (mass number 12)
• Carbon-14: 6 protons and 8 neutrons (mass number 14)

Carbon-14 is radioactive and used in carbon dating, while Carbon-12 is stable and the most common isotope.

Isotopes are often written with the element name followed by the mass number, e.g. hydrogen-1, hydrogen-2 (deuterium), and hydrogen-3 (tritium).

Atomic Models

Our understanding of the atom has changed over time through several models:

Plum Pudding Model

Proposed by J.J. Thomson in 1904 after discovering the electron. It pictured the atom as a positively charged 'pudding' with negatively charged electrons (the 'plums') scattered inside.

This model explained the atom as a whole being neutral but could not explain later experimental results.

Rutherford Scattering Experiment

In 1909, Ernest Rutherford and his team fired alpha particles at a thin gold foil. Most particles passed straight through, but some were deflected at large angles.

This showed that:

• The atom has a tiny, dense, positively charged nucleus that repels alpha particles.
• Most of the atom is empty space, allowing most particles to pass through.

This disproved the Plum Pudding Model and introduced the nuclear model of the atom.

Bohr
92s Model

Niels Bohr improved the nuclear model by suggesting that electrons orbit the nucleus in fixed shells or energy levels, not randomly.

Electrons can jump between shells by absorbing or emitting energy, explaining atomic spectra.

Bohr
92s model matches experimental evidence and is still used to explain atomic structure at GCSE level.

Ions and Ion Formation

An ion is an atom or group of atoms that has an electric charge because it has lost or gained electrons.

Atoms become ions by:

• Losing electrons to form positive ions (cations).
• Gaining electrons to form negative ions (anions).

The number of protons stays the same, but the number of electrons changes, so the overall charge is no longer zero.

For example, a sodium atom (Na) has 11 protons and 11 electrons. If it loses 1 electron, it becomes a sodium ion (Na⁺) with 11 protons and 10 electrons, giving it a +1 charge.

Similarly, a chlorine atom (Cl) has 17 protons and 17 electrons. If it gains 1 electron, it becomes a chloride ion (Cl^2121) with 17 protons and 18 electrons, giving it a 2121 charge.

Examples of Atomic Structure Calculations