What you'll learn
This topic forms the foundation of CXC CSEC Chemistry, covering the structure of atoms and the three subatomic particles that compose them. You'll learn how to calculate the number of protons, neutrons and electrons in any atom or ion, interpret atomic notation, and understand isotopes—concepts that appear in multiple-choice, structured and extended response questions every year.
Key terms and definitions
Atom — the smallest particle of an element that can exist while retaining the chemical properties of that element.
Proton — a positively charged subatomic particle found in the nucleus of an atom, with a relative mass of 1 and a relative charge of +1.
Neutron — a subatomic particle with no charge (neutral) found in the nucleus, with a relative mass of 1.
Electron — a negatively charged subatomic particle that orbits the nucleus in energy levels (shells), with negligible mass (1/1840 of a proton) and a relative charge of -1.
Atomic number (Z) — the number of protons in the nucleus of an atom; this defines the element and determines its position in the periodic table.
Mass number (A) — the total number of protons and neutrons in the nucleus of an atom.
Isotopes — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers.
Ion — an atom or group of atoms that has lost or gained electrons, resulting in a net positive (cation) or negative (anion) charge.
Core concepts
Structure of the atom
The atom consists of a central nucleus containing protons and neutrons, surrounded by electrons arranged in energy levels or electron shells. The nucleus accounts for nearly all the atom's mass but occupies only a tiny fraction of its volume—if an atom were the size of the National Stadium in Trinidad, the nucleus would be smaller than a marble at the centre.
Key characteristics of subatomic particles:
| Particle | Relative Mass | Relative Charge | Location |
|---|---|---|---|
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | 1/1840 (≈0) | -1 | Shells |
All atoms of a given element contain the same number of protons. For example, every carbon atom contains exactly 6 protons, every aluminium atom (widely used in the Caribbean beverage industry for canning) contains 13 protons, and every calcium atom (essential for Jamaica's limestone deposits) contains 20 protons.
Atomic number and mass number
The atomic number (Z) identifies the element. Sodium has atomic number 11, meaning every sodium atom contains 11 protons. Since atoms are electrically neutral, a neutral sodium atom also contains 11 electrons.
The mass number (A) is the sum of protons and neutrons:
Mass number (A) = number of protons + number of neutrons
Rearranging this equation:
Number of neutrons = mass number - atomic number
or N = A - Z
Standard atomic notation
Atoms are represented using standard notation that shows both the mass number and atomic number:
$$^{A}_{Z}\text{X}$$
Where:
- X is the chemical symbol
- A is the mass number (top)
- Z is the atomic number (bottom)
For example, $^{23}_{11}\text{Na}$ represents a sodium atom with:
- 11 protons (atomic number)
- 23 - 11 = 12 neutrons
- 11 electrons (in a neutral atom)
Sometimes you'll see simplified notation showing only the mass number: sodium-23 or Na-23.
Isotopes
Isotopes are atoms of the same element with different numbers of neutrons. They have identical chemical properties because they have the same number of electrons, but different physical properties due to their different masses.
Carbon isotopes:
- $^{12}_{6}\text{C}$ — 6 protons, 6 neutrons, 6 electrons
- $^{13}_{6}\text{C}$ — 6 protons, 7 neutrons, 6 electrons
- $^{14}_{6}\text{C}$ — 6 protons, 8 neutrons, 6 electrons (radioactive; used in carbon dating)
Chlorine isotopes:
- $^{35}_{17}\text{Cl}$ — 17 protons, 18 neutrons, 17 electrons (75% abundance)
- $^{37}_{17}\text{Cl}$ — 17 protons, 20 neutrons, 17 electrons (25% abundance)
The existence of isotopes explains why relative atomic masses on the periodic table are rarely whole numbers. Chlorine's relative atomic mass of 35.5 reflects the weighted average of its two main isotopes.
Caribbean relevance: Isotopes of uranium are processed at conversion facilities, and understanding isotopic composition is essential for quality control in Trinidad and Tobago's petrochemical industry when analysing hydrocarbon feedstocks.
Ions and electron configuration
Cations (positive ions) form when atoms lose electrons:
- Sodium atom: 11 protons, 11 electrons → Sodium ion (Na⁺): 11 protons, 10 electrons
- Calcium atom: 20 protons, 20 electrons → Calcium ion (Ca²⁺): 20 protons, 18 electrons
- Aluminium atom: 13 protons, 13 electrons → Aluminium ion (Al³⁺): 13 protons, 10 electrons
Anions (negative ions) form when atoms gain electrons:
- Chlorine atom: 17 protons, 17 electrons → Chloride ion (Cl⁻): 17 protons, 18 electrons
- Oxygen atom: 8 protons, 8 electrons → Oxide ion (O²⁻): 8 protons, 10 electrons
Critical point: The number of protons never changes during ion formation. If the number of protons changed, the element itself would change. Only electrons are gained or lost.
For ions, use this relationship:
Number of electrons = number of protons - charge
For Mg²⁺: electrons = 12 - (+2) = 10 For S²⁻: electrons = 16 - (-2) = 18
Calculating subatomic particles
Follow this systematic approach for any atom or ion:
- Identify the element from the symbol to find the atomic number (number of protons)
- Read the mass number from the notation
- Calculate neutrons using: neutrons = mass number - atomic number
- Determine electrons:
- For neutral atoms: electrons = protons
- For ions: electrons = protons - charge
Example with $^{40}_{18}\text{Ar}$:
- Protons = 18 (atomic number)
- Neutrons = 40 - 18 = 22
- Electrons = 18 (neutral atom)
Example with $^{27}_{13}\text{Al}^{3+}$:
- Protons = 13
- Neutrons = 27 - 13 = 14
- Electrons = 13 - 3 = 10
Worked examples
Example 1: Determining subatomic particles
Question: An atom of potassium is represented as $^{39}_{19}\text{K}$.
(a) State the number of protons, neutrons and electrons in this atom. [3 marks]
(b) Potassium exists as three isotopes: K-39, K-40 and K-41. Explain what is meant by the term isotopes. [2 marks]
Solution:
(a)
- Protons = 19 ✓
- Neutrons = 39 - 19 = 20 ✓
- Electrons = 19 ✓
(b) Isotopes are atoms of the same element ✓ with the same number of protons but different numbers of neutrons/different mass numbers. ✓
Examiner note: For part (a), show your working for neutrons. For part (b), you must mention both the similarity (same element/protons) and the difference (different neutrons/mass numbers) to earn both marks.
Example 2: Ions and electron configuration
Question: The bauxite mined in Jamaica contains aluminium compounds. An aluminium ion has the symbol Al³⁺ and mass number 27.
(a) State the atomic number of aluminium. [1 mark]
(b) Determine the number of protons, neutrons and electrons in the Al³⁺ ion. [3 marks]
(c) Explain why the aluminium ion has a 3+ charge. [2 marks]
Solution:
(a) Atomic number = 13 ✓ (from periodic table)
(b)
- Protons = 13 ✓
- Neutrons = 27 - 13 = 14 ✓
- Electrons = 13 - 3 = 10 ✓
(c) The aluminium atom has lost 3 electrons ✓, leaving more protons than electrons, resulting in a net positive charge of 3+. ✓
Examiner note: Always check the periodic table for atomic numbers when not given. Remember that the number of protons never changes; only electrons are lost or gained during ion formation.
Example 3: Isotopes and average atomic mass
Question: Chlorine has two isotopes: $^{35}{17}\text{Cl}$ (75% abundance) and $^{37}{17}\text{Cl}$ (25% abundance).
(a) State one similarity and one difference between these two isotopes. [2 marks]
(b) The relative atomic mass of chlorine is 35.5. Explain why this value is not a whole number. [2 marks]
Solution:
(a)
- Similarity: Both have 17 protons/same atomic number/same number of electrons ✓
- Difference: Different numbers of neutrons/different mass numbers ✓
(b) The relative atomic mass is a weighted average ✓ of the two isotopes based on their natural abundance. ✓
Examiner note: Questions about why relative atomic masses aren't whole numbers regularly appear on CXC papers. Always mention "weighted average" and "abundance" or "proportion" of isotopes.
Common mistakes and how to avoid them
• Mistake: Confusing atomic number with mass number or using them interchangeably. Correction: Atomic number (Z) is always the number of protons and never changes for a given element. Mass number (A) is protons plus neutrons and can vary for isotopes of the same element.
• Mistake: Changing the number of protons when forming ions, such as stating that Na⁺ has 10 protons. Correction: Only electrons change during ion formation. Na⁺ still has 11 protons but only 10 electrons. If protons changed, it would become a different element.
• Mistake: Calculating electrons in an ion by forgetting to account for the charge, giving the neutral atom's electron count. Correction: For Ca²⁺, electrons = 20 - 2 = 18, not 20. For Cl⁻, electrons = 17 - (-1) = 18, not 17. Always apply: electrons = protons - charge.
• Mistake: Stating that isotopes have different numbers of electrons or different chemical properties. Correction: Isotopes have identical numbers of protons AND electrons in their neutral atoms, giving them the same chemical properties. Only the number of neutrons differs.
• Mistake: Writing the mass number at the bottom and atomic number at the top in standard notation. Correction: Always write mass number at the top: $^{A}_{Z}\text{X}$. Remember "Mass is at the top" or that mass is "heavier, so sits on top."
• Mistake: Forgetting that electrons have negligible mass when calculating mass number. Correction: Mass number = protons + neutrons only. Electrons contribute virtually nothing to atomic mass (1/1840 of a proton).
Exam technique for Atomic Structure: Protons, Neutrons and Electrons
• Command words matter: "State" requires a direct answer with no explanation (1 mark). "Explain" requires reasoning with linking words like "because," "therefore," or "so" (typically 2 marks). "Calculate" requires showing your working clearly, especially for neutron calculations.
• Standard notation questions: When asked to interpret or complete notation like $^{A}_{Z}\text{X}$, always identify which number is which first. Write down protons = Z, then calculate neutrons = A - Z before considering electrons. This systematic approach prevents errors under exam pressure.
• Isotope questions: These frequently ask you to "explain what is meant by isotopes." Your answer must include two parts: (1) atoms of the same element/same number of protons, and (2) different numbers of neutrons/different mass numbers. Missing either part loses marks.
• Show all working: For calculations involving subatomic particles, always show your method even for simple subtractions. Write "neutrons = 40 - 18 = 22" not just "22." Examiners award method marks even if your final answer is incorrect due to an earlier error.
Quick revision summary
Atoms contain three subatomic particles: protons (+1 charge, mass 1) and neutrons (0 charge, mass 1) in the nucleus, with electrons (-1 charge, negligible mass) in surrounding shells. Atomic number (Z) equals the number of protons and defines the element. Mass number (A) equals protons plus neutrons. Isotopes are atoms of the same element with different numbers of neutrons. Ions form when atoms gain or lose electrons, not protons. For any particle, use: neutrons = A - Z, and for ions, electrons = protons - charge.