Atomic models, subatomic particles, isotopes and electronic structure

Chemistry · Atomic structure and the periodic table. The idea in plain words, the rules to remember, the mistake to avoid and a worked example, written and checked by Mega Lecture.

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The idea

Ideas about the atom changed as new evidence was found. The plum pudding model was a ball of positive charge with electrons spread through it. The alpha particle scattering experiment showed that the mass and positive charge are in a tiny nucleus and that most of the atom is empty space. Bohr then placed the electrons in shells at fixed distances, and Chadwick gave evidence for neutrons.

The atomic number is the number of protons. The mass number is the number of protons plus neutrons. Isotopes are atoms of the same element with different numbers of neutrons. They react in the same way because they have the same electronic structure.

Electrons fill the shells in order: up to 2 in the first shell, then 8, then 8.

Rules to remember

Relative charge: proton +1, neutron 0, electron −1.
Relative mass: proton 1, neutron 1, electron very small.
number of neutrons = mass number − atomic number
In an atom, electrons = protons. A positive ion has lost electrons; a negative ion has gained electrons.
Radius of an atom is about 0.1 nm (1 × 10⁻¹⁰ m); the radius of the nucleus is less than 1/10 000 of this.
relative atomic mass = sum of (mass number × percentage of each isotope) ÷ 100
Electronic structure example: sodium (11 electrons) is 2,8,1.

Common mistake

For an ion, students change the number of protons. Only the number of electrons changes: Mg²⁺ still has 12 protons but has 10 electrons.

Worked example

Chlorine is 75% chlorine-35 and 25% chlorine-37. Calculate its relative atomic mass.

Multiply each mass number by its percentage: 35 × 75 = 2625 and 37 × 25 = 925.
Add them: 2625 + 925 = 3550.
Divide by 100: 3550 ÷ 100 = 35.5.

Answer: Relative atomic mass of chlorine = 35.5

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