The atom is the smallest unit of matter that retains the chemical properties of an element. By understanding atomic structure — protons, neutrons and electrons — you can explain the periodic table, chemical bonding and all of chemistry.
11.1 Atomic Structure
Definition
Atomic number (Z)
The atomic number of an element is the number of protons in the nucleus of one atom of that element. The atomic number defines the element — all atoms of the same element have the same atomic number.
Definition
Mass number (A)
The mass number of an atom is the total number of protons and neutrons (nucleons) in the nucleus of that atom. A = Z + N.
Definition
Protons
Protons are positively charged particles found in the atomic nucleus. Each proton carries a charge of +1,6 × 10⁻¹⁹ C and has a relative mass of 1 atomic mass unit (u).
Definition
Neutrons
Neutrons are electrically neutral particles found in the atomic nucleus. A neutron has approximately the same mass as a proton (1 u) and carries no electric charge.
Definition
Electrons
Electrons are negatively charged particles that occupy energy levels (shells) around the atomic nucleus. Each electron carries a charge of −1,6 × 10⁻¹⁹ C and has a mass approximately 1/1 836 of a proton.
Formula
Number of neutrons
N = number of neutrons, A = mass number, Z = atomic number (number of protons)
SI unit: no unit (count)
A neutral atom contains equal numbers of protons and electrons, so the positive and negative charges cancel. If an atom loses electrons, it becomes a positively charged ion (cation). If it gains electrons, it becomes a negatively charged ion (anion). The number of protons NEVER changes in ordinary chemical reactions — changing the number of protons creates a different element (nuclear reaction). Atomic notation is written as: ᴬ_Z E, where E is the element symbol, A is the mass number (top left) and Z is the atomic number (bottom left). For example, sodium is written ²³₁₁Na.
Note
It is the number of PROTONS that determines the element. All carbon atoms have exactly 6 protons — if you add or remove a proton you no longer have carbon. Neutrons contribute to mass but not to chemical identity.
Worked Example
An atom has atomic number Z = 11 and mass number A = 23. Find: (a) number of protons, (b) number of neutrons, (c) number of electrons in the neutral atom, (d) write the atomic notation.
Given
- Z = 11
- A = 23
Find
(a) protons (b) neutrons (c) electrons (d) atomic notation
Solution
- 1(a) Number of protons = Z = 11
- 2(b) N = A − Z = 23 − 11 = 12 neutrons
- 3(c) Neutral atom: electrons = protons = 11
- 4(d) ²³₁₁Na
Practice Question
For ⁴⁰₂₀Ca, find: (a) the atomic number Z, (b) the mass number A, (c) the number of neutrons, (d) the number of electrons in the neutral atom.
(4 marks)
11.2 Isotopes and Electron Configuration
Definition
Isotopes
Isotopes are atoms of the same element (same atomic number Z) that have different mass numbers (different numbers of neutrons). Isotopes of an element have identical chemical properties but different masses.
Definition
Relative atomic mass (Ar)
The relative atomic mass of an element is the weighted average of the masses of the naturally occurring isotopes of that element, relative to one-twelfth of the mass of a ¹²C atom.
Definition
Electron configuration
The electron configuration of an atom is the arrangement of electrons in the energy levels (shells) of that atom, starting from the innermost shell.
Definition
Valence electrons
Valence electrons are the electrons in the outermost occupied energy level (valence shell) of an atom. They determine the chemical reactivity and bonding behaviour of the atom.
Formula
Relative atomic mass (weighted average)
Ar = relative atomic mass (no unit), fractional abundance = fraction of naturally occurring atoms of that isotope, isotope mass = mass number of that isotope
SI unit: no unit (ratio)
Electron shell filling rules (Grade 10)
- First shell (n = 1): maximum 2 electrons
- Second shell (n = 2): maximum 8 electrons
- Third shell (n = 3): maximum 8 electrons (at Grade 10 level)
- Always fill inner shells first before adding electrons to the next shell (aufbau principle)
Write the electron configuration as numbers separated by commas, showing how many electrons are in each shell from the inside out. For example, sodium (Na, Z = 11) has 11 electrons: 2 in shell 1, 8 in shell 2, and 1 in shell 3 — written 2,8,1. The last number (1) is the number of valence electrons. The group number of a main-group element equals the number of valence electrons. The period number equals the number of occupied shells.
Worked Example
Chlorine has two naturally occurring isotopes: ³⁵Cl (relative abundance 75%) and ³⁷Cl (relative abundance 25%). Calculate the relative atomic mass of chlorine.
Given
- ³⁵Cl: abundance = 75% = 0,75; mass = 35
- ³⁷Cl: abundance = 25% = 0,25; mass = 37
Find
Ar of chlorine
Solution
- 1Ar = (fraction₁ × mass₁) + (fraction₂ × mass₂)
- 2Ar = (0,75 × 35) + (0,25 × 37)
- 3Ar = 26,25 + 9,25
- 4Ar = 35,5
Worked Example
Write the electron configuration for: (a) Oxygen (Z = 8), (b) Magnesium (Z = 12), (c) Chlorine (Z = 17).
Given
- Z(O) = 8
- Z(Mg) = 12
- Z(Cl) = 17
Find
Electron configurations
Solution
- 1(a) O: 8 electrons → fill shell 1 (2), then shell 2 (6) → 2,6
- 2(b) Mg: 12 electrons → fill shell 1 (2), shell 2 (8), shell 3 (2) → 2,8,2
- 3(c) Cl: 17 electrons → shell 1 (2), shell 2 (8), shell 3 (7) → 2,8,7
Exam Tip
The group number of a main-group element = number of valence electrons. The period number = number of occupied shells. So: Period 2, Group 6 → the element has 6 valence electrons in 2 shells (electron config 2,6) → that is oxygen.
Practice Question
Bromine has two naturally occurring isotopes: ⁷⁹Br (relative abundance 50,5%) and ⁸¹Br (relative abundance 49,5%). Calculate the relative atomic mass of bromine.
(4 marks)