Atomic Structure and the Periodic Table

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:

mass number A (top left), atomic number Z (bottom left), symbol X — as in ²³₁₁Na

Where:

• X is the chemical symbol
• A is the mass number (top)
• Z is the atomic number (bottom)

For example, ²³₁₁Na 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:

• ¹²₆C — 6 protons, 6 neutrons, 6 electrons
• ¹³₆C — 6 protons, 7 neutrons, 6 electrons
• ¹⁴₆C — 6 protons, 8 neutrons, 6 electrons (radioactive; used in carbon dating)

Chlorine isotopes:

• ³⁵₁₇Cl — 17 protons, 18 neutrons, 17 electrons (75% abundance)
• ³⁷₁₇Cl — 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:

1. Identify the element from the symbol to find the atomic number (number of protons)
2. Read the mass number from the notation
3. Calculate neutrons using: neutrons = mass number - atomic number
4. Determine electrons:
   • For neutral atoms: electrons = protons
   • For ions: electrons = protons - charge

Example with ⁴⁰₁₈Ar:

• Protons = 18 (atomic number)
• Neutrons = 40 - 18 = 22
• Electrons = 18 (neutral atom)

Example with ²⁷₁₃Al³⁺:

• 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 ³⁹₁₉K.

(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: ³⁵₁₇Cl (75% abundance) and ³⁷₁₇Cl (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.

The Periodic Table

The Periodic Table is the chemist's most powerful organising tool. It arranges all the known elements so that their properties follow a repeating ("periodic") pattern, allowing you to predict how an element will behave from its position alone. For CSEC Chemistry you must understand how the table is arranged and be able to describe the trends in Group I, Group VII and Group 0, plus the general features of the transition metals.

How the table is arranged

• Elements are arranged in order of increasing atomic (proton) number.
• A horizontal row is a period. The period number equals the number of electron shells the atoms have.
• A vertical column is a group. The group number equals the number of electrons in the outer shell (for the main groups).

This last point is the key to everything: elements in the same group have the same number of outer electrons, so they react in similar ways. Chemical properties depend on the outer electrons.

• Metals are found on the left and centre; non-metals on the right. A "staircase" line separates them.

Group I — the alkali metals (Li, Na, K …)

These are soft, reactive metals with one outer electron, which they lose easily to form +1 ions.

• They react with water to give a metal hydroxide (an alkali) and hydrogen gas: e.g. sodium + water → sodium hydroxide + hydrogen.
• Reactivity increases down the group (lithium → sodium → potassium become more vigorous).

Why? Going down the group the atoms get larger, the outer electron is further from the nucleus and is shielded by more inner shells, so it is held less strongly and lost more easily — making the metal more reactive.

Group VII — the halogens (F, Cl, Br, I)

These are reactive non-metals with seven outer electrons. They gain one electron to form −1 ions (halides), or share electrons to form diatomic molecules (Cl₂, Br₂…).

• Reactivity decreases down the group (chlorine is more reactive than iodine) — the opposite trend to Group I, because a larger atom gains an electron less easily.
• A more reactive halogen displaces a less reactive one from a solution of its salt. For example, chlorine added to potassium bromide displaces bromine, turning the solution orange:

  chlorine + potassium bromide → potassium chloride + bromine

Group 0 — the noble gases (He, Ne, Ar …)

These are unreactive (inert) gases. Their atoms have a full outer shell of electrons (2 for helium, 8 for the others), so they have no tendency to gain, lose or share electrons. Their stability is the reason other atoms react — atoms react in order to achieve a full outer shell like a noble gas.

Uses follow from their inertness: helium in balloons (low density, non-flammable), argon in light bulbs and welding (provides an unreactive atmosphere), neon in lights.

The transition metals

Found in the central block, these are typical metals with extra characteristic properties:

• high melting points and high density;
• often form coloured compounds (e.g. copper salts are blue, iron(III) salts are orange-brown);
• can show more than one oxidation state / valency (e.g. iron forms Fe²⁺ and Fe³⁺);
• many act as catalysts (e.g. iron in the Haber process, nickel in hydrogenation).

These contrast with Group I metals, which are soft, have low densities and melting points, are very reactive and form only white/colourless +1 compounds.

Using the table to predict

Because position reveals outer-shell electrons, you can predict:

• the charge of an ion (Group I → +1, Group II → +2, Group VI → −2, Group VII → −1);
• the formula of a compound (e.g. magnesium, Group II, with chlorine, Group VII → MgCl₂);
• whether an element is a metal or non-metal, and roughly how reactive it is.

A short history — why the table looks as it does

The modern table grew out of the work of Dmitri Mendeleev in 1869. He arranged the known elements in order of atomic mass and lined up those with similar properties into groups. His insight was twofold: he left gaps for elements not yet discovered, and he even predicted their properties from their position. When elements such as gallium and germanium were later found and matched his predictions, the periodic table was accepted. The one flaw — a few pairs that seemed out of order by mass — was resolved when elements were re-ordered by atomic (proton) number instead of mass, which is how the table is arranged today. Knowing this story helps you explain why position predicts properties: the table is built so that repeating patterns in electron arrangement line up.

Period 3 across the table

Looking across a single period also shows clear trends. Moving across Period 3 from sodium to argon:

• the elements change from metals (Na, Mg, Al) through a metalloid (Si) to non-metals (P, S, Cl) and finally a noble gas (Ar);
• the number of outer electrons increases from 1 to 8;
• oxides change from basic (metal oxides, e.g. Na₂O) to acidic (non-metal oxides, e.g. SO₂).

So both the group (down) and the period (across) give predictable patterns — the essence of "periodicity."

Worked example — predicting a formula

Suppose you are asked for the formula of the compound formed between calcium and bromine. Calcium is in Group II, so it forms Ca²⁺. Bromine is in Group VII, so it forms Br⁻. To balance the charges you need two bromide ions for each calcium ion, giving CaBr₂. You reached the answer using nothing but the elements' positions — which is exactly what the periodic table is for.

Common exam mistakes

• Saying elements are arranged by mass — modern tables use atomic (proton) number.
• Giving the wrong reactivity trend: Group I gets more reactive down the group, Group VII gets less reactive down the group.
• Explaining a trend by "more shells" alone — for full marks state the chain: larger atom → outer electron further from nucleus → more shielding → held less strongly.
• Calling the noble gases reactive; they are inert because their outer shell is full.

Key terms to remember

• Group — a vertical column; elements share the same number of outer electrons and similar properties.
• Period — a horizontal row; the period number equals the number of electron shells.
• Atomic (proton) number — the number of protons; the basis for ordering the table.
• Alkali metals (Group I) — soft, reactive metals forming +1 ions; reactivity increases down.
• Halogens (Group VII) — reactive non-metals forming −1 ions; reactivity decreases down.
• Noble gases (Group 0) — inert gases with a full outer shell.
• Transition metals — central block; high melting points, coloured compounds, variable valency, good catalysts.
• Displacement — a more reactive halogen pushing a less reactive one out of its salt solution.

Quick recap

• The table is ordered by atomic number; period = number of shells, group = number of outer electrons.
• Same group = similar chemistry, because the outer electrons match.
• Group I: soft, +1 ions, reactivity increases down (lost electron more easily).
• Group VII: −1 ions, reactivity decreases down; more reactive halogens displace less reactive ones.
• Group 0: inert, full outer shells.
• Transition metals: high m.p., coloured compounds, variable valency, good catalysts.

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 the mass number at the top left and the atomic number at the bottom left, as in ²³₁₁Na. 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 and the Periodic Table

• 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 of this kind, 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.