Topic 1 of 10

Atomic Structure & the Periodic Table

Explore the building blocks of matter — from atoms and elements to isotopes, electron shells, and the development of the modern periodic table.

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Atoms, Elements & Compounds

What is an Atom?

An atom is the smallest part of an element that can exist. Think of it as the ultimate LEGO brick. Everything in the universe — the air you breathe, the water you drink, and even you — is made of atoms.

Atoms are incredibly empty! If an atom were the size of a football stadium, its nucleus (the central part) would only be the size of a marble.

Elements

An element is a substance made of only one type of atom. There are about 118 known elements, and each has a unique atomic number (the number of protons in its nucleus). All known elements are catalogued in the periodic table. Examples include Hydrogen (H), Oxygen (O), Carbon (C), and Iron (Fe).

Compounds

A compound is a substance formed when two or more elements are chemically bonded together. The atoms in a compound are held together by chemical bonds (ionic or covalent), and the compound has different properties from the individual elements. A compound can only be separated into its elements by chemical reactions, not by physical methods.

H₂ + Cl₂ → 2HCl

In this example, hydrogen and chlorine (both elements) react to form hydrogen chloride (a compound).

When describing the difference between a mixture and a compound, always state that the elements in a compound are chemically bonded, while in mixtures they are not. A compound has properties different from its constituent elements.

Mixtures

A mixture consists of two or more substances that are not chemically bonded together. The substances retain their individual properties and can be separated by physical methods.

Key Properties of Mixtures

  • The components are not chemically combined.
  • They retain their original properties.
  • They can be separated by physical processes (filtration, distillation, etc.).
  • No chemical reaction is needed to separate them.

Examples include air (a mixture of gases), salt water, and alloys like steel.

Separation Techniques

Filtration

This technique is used to separate an insoluble solid from a liquid. The mixture is poured through filter paper. The tiny pores in the paper allow the liquid particles to pass through, but they are too small for larger solid particles, which get trapped.

Crystallisation

This is used to obtain a dissolved solid (solute) from a solution. The solution is heated so the solvent evaporates. As the solution becomes more concentrated, crystals begin to form. The solution is then left to cool, allowing the crystals to grow.

Simple Distillation

This separates a solvent from a solution. The solution is heated until the solvent boils and turns into a gas. The gas is then cooled and condensed back into a pure liquid in a separate container using a condenser.

Fractional Distillation

This separates a mixture of liquids with different boiling points. The mixture is heated, and the liquid with the lowest boiling point evaporates first. The vapour passes up a fractionating column where it cools and condenses separately. This is repeated for each liquid in order of their boiling points.

Chromatography

This is used to separate mixtures of dissolved substances, typically dyes or inks. A spot of the mixture is placed on chromatography paper and placed in a solvent. As the solvent travels up the paper, it carries the different substances at different rates, separating them into distinct spots.

You should be able to describe when each technique should be used. For example, filtration for an insoluble solid from a liquid; distillation for a solvent from a solution; chromatography for dissolved substances like dyes.

The Changing Model of the Atom

Our understanding of atomic structure has evolved dramatically over centuries. Each new model built upon or replaced the one before it, driven by experimental evidence.

Early Ideas — Democritus (400 BC)

The Greek philosopher Democritus proposed that if you kept cutting matter into smaller and smaller pieces, you would eventually reach a tiny, indivisible particle. He called this an "atomos" — meaning "uncuttable".

John Dalton's Solid Sphere Model (1803)

Dalton described atoms as tiny, solid, indivisible spheres. He proposed that each element was made of a different type of atom, and that atoms could not be broken down further. This was the first scientifically-based atomic theory.

J.J. Thomson's "Plum Pudding" Model (1897)

Thomson discovered the electron using cathode rays. He proposed the "plum pudding" model: a sphere of positive charge with tiny negative electrons embedded within it, like plums in a pudding.

Rutherford's Nuclear Model (1909)

Rutherford's famous alpha particle scattering experiment (firing alpha particles at gold foil) showed that most of an atom is empty space, with a tiny, dense, positively charged nucleus at its centre. This disproved the plum pudding model and led to the nuclear model of the atom.

Niels Bohr's Shell Model (1913)

Bohr refined Rutherford's model by proposing that electrons orbit the nucleus in specific energy levels (shells) at fixed distances from the nucleus, rather than randomly. This model successfully explained the emission spectra of hydrogen.

The accepted model of the atom today is the nuclear model: a small, dense nucleus containing protons and neutrons, surrounded by electrons orbiting in energy levels (shells).

Subatomic Particles

Atoms are made up of three types of subatomic particles, each with specific properties:

  • Proton: Found in the nucleus. Relative mass = 1, Relative charge = +1.
  • Neutron: Found in the nucleus. Relative mass = 1, Relative charge = 0.
  • Electron: Found in shells orbiting the nucleus. Relative mass ≈ 1/2000 (negligible), Relative charge = −1.
The number of protons defines which element an atom is. This is called the atomic number (Z). The mass number (A) is the total number of protons + neutrons. In a neutral atom, the number of electrons equals the number of protons.
To find the number of neutrons: Neutrons = Mass Number − Atomic Number. Remember that in a neutral atom, protons = electrons.

Atomic Size & Mass Distribution

The Scale of an Atom

Atoms are unimaginably small, and the nucleus is even smaller in comparison.

  • The atomic radius is approximately 1 × 10⁻¹⁰ metres.
  • The radius of the nucleus is about 1 × 10⁻¹⁴ metres.

This means the nucleus is roughly 10,000 times smaller than the entire atom. If an atom were the size of a football stadium, the nucleus would be the size of a pea placed in the centre.

Distribution of Mass

Despite its tiny size, the nucleus contains nearly all of the atom's mass. This is because protons and neutrons (found in the nucleus) each have a relative mass of 1. Electrons have a negligible mass (approximately 1/2000th of a proton or neutron). Therefore, when calculating the mass of an atom, the mass of the electrons is considered to be zero.

Be prepared to use standard form to compare the sizes of the atom and its nucleus. For example, an atom (radius 1 × 10⁻¹⁰ m) is 10,000 (or 10⁴) times larger than its nucleus (radius 1 × 10⁻¹⁴ m).

Isotopes & Relative Atomic Mass

Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons. Because isotopes have the same number of protons, they also have the same number of electrons — so they have identical chemical properties.

Understanding Isotopes

The key difference between isotopes is their mass number. Let's look at carbon as an example — all carbon atoms have 6 protons:

  • Carbon-12: 6 protons + 6 neutrons (Mass number = 12)
  • Carbon-13: 6 protons + 7 neutrons (Mass number = 13)
  • Carbon-14: 6 protons + 8 neutrons (Mass number = 14)

Relative Atomic Mass (Ar)

The relative atomic mass is the weighted mean mass of an atom of an element, taking into account its naturally occurring isotopes. This is why some elements on the periodic table have mass numbers that are not whole numbers.

Calculating the Ar of Chlorine

In nature, chlorine exists as two isotopes: 75% is Chlorine-35 and 25% is Chlorine-37.

Step 1: Multiply abundance by mass: (75 × 35) = 2625; (25 × 37) = 925

Step 2: Add: 2625 + 925 = 3550

Step 3: Divide by 100: 3550 ÷ 100 = 35.5

Electron Shells & Configuration

The way electrons are arranged in an atom is fundamental to chemistry. It dictates how an element behaves, why it reacts, and where it sits in the periodic table.

Rules for Filling Shells

  • Electrons always occupy the lowest available energy level first (the innermost shell).
  • Shell 1: holds a maximum of 2 electrons.
  • Shell 2: holds a maximum of 8 electrons.
  • Shell 3: holds a maximum of 8 electrons.

How to Work Out Electron Configurations

Sodium (Na) — Atomic number 11

First shell fills with 2 electrons.

Second shell fills with 8 electrons. (2 + 8 = 10 used).

Remaining 1 electron goes into the third shell.

Electron configuration: 2.8.1

Chlorine (Cl) — Atomic number 17

First shell: 2 electrons. Second shell: 8 electrons. Third shell: 7 electrons.

Electron configuration: 2.8.7

Why Electron Arrangement Matters

The electrons in the outermost shell are called valence electrons. The number of valence electrons determines the chemical properties of an element. Elements in the same group have the same number of valence electrons — which is why they have similar chemical properties.

An atom is most stable when it has a full outer shell. The noble gases (Group 0) have full outer shells, which is why they are so unreactive. Other atoms react by gaining, losing, or sharing electrons to achieve a stable, full outer shell — this is the basis for all chemical bonding.

The Periodic Table

The periodic table is a masterfully organised chart of all the known chemical elements. It's arranged to reveal patterns in the properties of elements — a concept known as periodicity.

How the Table is Arranged

The modern periodic table arranges elements in order of increasing atomic number (Z). The atomic number is the number of protons — the element's unique ID number.

A common mistake is to say the table is ordered by relative atomic mass. Always state that the modern periodic table is ordered by atomic number.

Periods (Horizontal Rows)

The rows across the periodic table are called periods. The period number tells you how many occupied electron shells an element has.

  • Period 1 elements (H, He) have 1 electron shell.
  • Period 2 elements (Li to Ne) have 2 electron shells.
  • Period 3 elements (Na to Ar) have 3 electron shells.

Groups (Vertical Columns)

The columns down the periodic table are called groups. For the main groups, the group number tells you the number of electrons in the element's outermost shell (valence electrons).

  • Group 1 (Alkali Metals): 1 outer electron — very reactive.
  • Group 7 (Halogens): 7 outer electrons — very reactive (need one more for a full shell).
  • Group 0 (Noble Gases): Full outer shell — stable and very unreactive.

History of the Periodic Table

Dmitri Mendeleev's Revolutionary Idea

Mendeleev arranged the known elements in order of increasing relative atomic mass and noticed a periodic pattern in their properties. His genius was demonstrated in two key ways:

  • He left gaps for elements he believed had not yet been discovered.
  • He made predictions about the properties of these missing elements (e.g., "eka-aluminium" and "eka-silicon").

When Gallium (1875) and Germanium (1886) were discovered, their properties almost perfectly matched Mendeleev's predictions — powerful evidence his arrangement was correct.

The Shift to Atomic Number

Despite its success, some pairs of elements (like Argon and Potassium) appeared in the wrong order when ordered by mass. The discovery of isotopes showed this was because an element's identity comes from its proton number, not its mass.

When asked why the modern table is ordered by atomic number instead of atomic mass, mention that the existence of isotopes meant that ordering by mass could incorrectly place some elements. Ordering by atomic number, which is unique to each element, fixes this.

Metals & Non-Metals

The periodic table has a fundamental dividing line that separates elements into metals (left and centre) and non-metals (right).

Reactions of Oxides

Metal Oxides are Basic

Metal oxides react with acids in neutralisation reactions to produce a salt and water.

MgO + 2HCl → MgCl₂ + H₂O

Non-Metal Oxides are Acidic

Non-metal oxides typically react with water to form acidic solutions.

SO₂ + H₂O → H₂SO₃
When asked to classify an element, first state its position on the periodic table. Then support this by citing two typical physical or chemical properties.

Group 0: The Noble Gases

The elements in Group 0 (Helium, Neon, Argon, etc.) are known as the noble gases. They are chemically inert — extremely unreactive.

Why are Noble Gases So Unreactive?

They all have a full outer shell of electrons:

  • Helium: 2
  • Neon: 2.8
  • Argon: 2.8.8

A full outer shell is the most stable arrangement. They have no tendency to lose, gain, or share electrons, so they exist as individual monatomic atoms.

Trends in Group 0

Boiling points increase as you go down the group. As atoms become larger with more electrons, the weak intermolecular forces of attraction between them become stronger. More energy is needed to overcome these forces.

Whenever you're asked why noble gases are unreactive, you must mention that they have a stable electron arrangement because they have a full outer shell.

Group 1: The Alkali Metals

The alkali metals (Lithium, Sodium, Potassium, etc.) all have one electron in their outermost shell.

Physical Properties

  • Very soft — can be cut with a knife.
  • Low density — Li, Na, and K float on water.
  • Low melting/boiling points — decrease as you go down the group.

Reactivity Trend

Reactivity increases down the group. As you go down, the outer electron is further from the nucleus with more inner shells providing shielding. This weaker attraction means the outer electron is lost more easily.

Reactions with Water

All alkali metals react vigorously with water to produce a metal hydroxide and hydrogen gas:

2Na + 2H₂O → 2NaOH + H₂
  • Lithium: Floats and fizzes steadily.
  • Sodium: More vigorous — melts into a silvery ball darting across the surface.
  • Potassium: Very rapid — ignites the hydrogen with a lilac flame.
When explaining the reactivity trend in Group 1, always refer to the single outer electron and how easily it is lost due to increasing atomic size and shielding.

Group 7: The Halogens

The halogens (Fluorine, Chlorine, Bromine, Iodine) all have seven electrons in their outermost shell. They exist as diatomic molecules (X₂), joined by a single covalent bond.

Physical Property Trends

Melting and boiling points increase down the group. As molecules become larger with more electrons, intermolecular forces strengthen, requiring more energy to change state.

Reactivity Trend

Reactivity decreases down the group. As the outer electron shell gets further from the nucleus with more shielding, it becomes harder for the nucleus to attract an extra electron to complete the outer shell.

Displacement Reactions

A more reactive halogen will displace a less reactive halide from an aqueous solution of its salt. For example, chlorine displaces bromide ions from potassium bromide solution.

Note the opposite reactivity trend compared to Group 1: Group 1 metals get MORE reactive down the group (easier to lose electrons), while Group 7 halogens get LESS reactive (harder to gain electrons).

Transition Metals

Transition metals occupy the central block of the periodic table. They have very different properties from the alkali metals.

High Melting Points & Densities

Transition metals have strong metallic bonding with powerful forces holding the atoms tightly together, resulting in high melting points, boiling points, and densities.

Coloured Compounds

Their compounds are often coloured (unlike Group 1 compounds which are white):

  • Copper(II) ions (Cu²⁺): characteristic blue colour.
  • Iron(III) ions (Fe³⁺): pale yellow-brown colour.

Variable Oxidation States

Most transition metals can form positive ions with different charges:

  • Iron forms Fe²⁺ (Iron(II)) and Fe³⁺ (Iron(III)).
  • Copper forms Cu⁺ (Copper(I)) and Cu²⁺ (Copper(II)).

Catalytic Activity

Many transition metals and their compounds are excellent catalysts:

  • Iron is used in the Haber process for making ammonia.
  • Copper(II) ions can catalyse the decomposition of hydrogen peroxide.
N₂ + 3H₂ ⇌ 2NH₃