Atomic structure and the periodic table: atoms, elements and compounds

What you'll learn

This revision guide covers the foundational concepts of matter that underpin all of GCSE Chemistry. You'll understand how atoms form the building blocks of everything around you, how elements are organised, and how compounds are created through chemical bonding. These concepts are essential for understanding reactivity, chemical equations, and the periodic table—topics that appear throughout your AQA GCSE Chemistry papers.

Key terms and definitions

Atom — the smallest particle of an element that can exist, consisting of a nucleus containing protons and neutrons, surrounded by electrons in shells

Element — a substance made of only one type of atom, which cannot be broken down into simpler substances by chemical means

Compound — a substance formed from two or more elements chemically bonded together in fixed proportions

Molecule — two or more atoms chemically bonded together, which may be of the same element (e.g. O₂) or different elements (e.g. H₂O)

Chemical formula — a representation using symbols and numbers to show the elements present in a substance and their ratios

Mixture — two or more substances (elements or compounds) that are not chemically bonded together and can be separated by physical methods

Ion — an atom or group of atoms that has gained or lost electrons, resulting in a positive or negative electrical charge

Isotope — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers

Core concepts

The structure of atoms

All matter is made of atoms. Each atom contains three types of subatomic particles:

Protons — positively charged particles found in the nucleus

• Relative mass: 1
• Relative charge: +1

Neutrons — particles with no charge found in the nucleus

• Relative mass: 1
• Relative charge: 0

Electrons — negatively charged particles that orbit the nucleus in shells

• Relative mass: 0.0005 (effectively zero compared to protons and neutrons)
• Relative charge: -1

The nucleus is extremely small compared to the overall size of the atom. If an atom were the size of a football stadium, the nucleus would be smaller than a pea at the centre. Despite this tiny size, the nucleus contains almost all the atom's mass because protons and neutrons are much heavier than electrons.

Atoms are electrically neutral because they contain equal numbers of protons (positive charges) and electrons (negative charges). The number of protons determines which element an atom is—this is called the atomic number or proton number.

Atomic number and mass number

Every element is defined by two key numbers:

Atomic number (proton number) — the number of protons in the nucleus of an atom

• This number is unique to each element
• It also equals the number of electrons in a neutral atom
• Elements are arranged in order of atomic number in the periodic table

Mass number (nucleon number) — the total number of protons plus neutrons in the nucleus

• Mass number = number of protons + number of neutrons
• To find the number of neutrons: neutrons = mass number - atomic number

Standard notation shows these numbers with the chemical symbol:

• Mass number is written as a superscript (top left)
• Atomic number is written as a subscript (bottom left)
• Example: ²³₁₁Na represents sodium with 11 protons and 12 neutrons

Elements and the periodic table

An element contains only one type of atom. There are approximately 100 different elements, each with unique properties determined by their atomic structure.

Elements are represented by chemical symbols:

• One or two letters (the first always capitalised)
• Often based on Latin names (e.g. Fe for iron comes from ferrum)
• Examples: C (carbon), O (oxygen), Na (sodium), Cl (chlorine)

Common elements you must know include:

• Hydrogen (H), carbon (C), nitrogen (N), oxygen (O)
• Sodium (Na), magnesium (Mg), aluminium (Al)
• Sulfur (S), chlorine (Cl), potassium (K), calcium (Ca)
• Iron (Fe), copper (Cu), zinc (Zn)

Elements can exist as single atoms (e.g. noble gases like helium) or as molecules. Some elements exist naturally as diatomic molecules (containing two atoms):

• H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂

Compounds and chemical formulae

Compounds form when atoms of two or more different elements chemically bond together in fixed proportions. The properties of a compound are completely different from the properties of the elements it contains.

Key characteristics of compounds:

• Elements are chemically bonded (not just mixed)
• Fixed composition and ratio of elements
• Can only be separated into elements by chemical reactions
• New properties different from constituent elements

Examples:

• Water (H₂O): hydrogen and oxygen combine to form a liquid at room temperature
• Sodium chloride (NaCl): reactive sodium metal and poisonous chlorine gas form edible table salt
• Carbon dioxide (CO₂): carbon and oxygen form a gas used by plants in photosynthesis

Chemical formulae show the elements in a compound and their ratios:

• Subscript numbers indicate how many atoms of each element
• No subscript means one atom
• H₂O contains 2 hydrogen atoms and 1 oxygen atom
• H₂SO₄ contains 2 hydrogen atoms, 1 sulfur atom, and 4 oxygen atoms

Prefixes in compound names indicate numbers of atoms:

• Mono- = 1 (often omitted)
• Di- = 2
• Tri- = 3
• Tetra- = 4
• Examples: carbon dioxide (CO₂), sulfur trioxide (SO₃), dinitrogen tetroxide (N₂O₄)

Naming compounds

Simple compounds follow clear naming patterns:

Metal + non-metal compounds:

• Name the metal first, then the non-metal with an '-ide' ending
• Examples: sodium chloride (NaCl), magnesium oxide (MgO), calcium sulfide (CaS)

Compounds containing oxygen and another non-metal:

• Name ends in '-ate' or '-ite'
• Examples: copper sulfate (CuSO₄), calcium carbonate (CaCO₃), sodium nitrate (NaNO₃)

Compounds containing three elements (one being oxygen):

• Usually end in '-ate'
• Sulfates contain sulfur and oxygen (SO₄ group)
• Carbonates contain carbon and oxygen (CO₃ group)
• Nitrates contain nitrogen and oxygen (NO₃ group)

Mixtures versus compounds

Understanding the difference between mixtures and compounds is crucial:

Mixtures:

• Components not chemically bonded
• Can be separated by physical methods (filtration, evaporation, distillation, chromatography)
• Components retain their individual properties
• No fixed composition—proportions can vary
• Examples: air (mixture of gases), seawater (salt dissolved in water), crude oil

Compounds:

• Elements chemically bonded
• Can only be separated by chemical reactions
• New properties completely different from elements
• Fixed composition—always same ratio of elements
• Examples: water (H₂O), methane (CH₄), sodium chloride (NaCl)

Air is a good example of a mixture commonly found in the Caribbean and UK contexts:

• Approximately 78% nitrogen
• Approximately 21% oxygen
• Small amounts of argon, carbon dioxide, water vapour, and other gases
• Proportions vary slightly by location (e.g. more water vapour in humid Caribbean climates)

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons. Because they have the same number of protons (and electrons), isotopes have the same chemical properties but different physical properties due to different masses.

Key points about isotopes:

• Same atomic number (same number of protons)
• Different mass numbers (different numbers of neutrons)
• Same chemical behaviour
• Different physical properties (e.g. density, rate of diffusion)

Common examples:

• Carbon-12 (¹²₆C): 6 protons, 6 neutrons

• Carbon-13 (¹³₆C): 6 protons, 7 neutrons

• Carbon-14 (¹⁴₆C): 6 protons, 8 neutrons (radioactive, used in carbon dating)

• Chlorine-35 (³⁵₁₇Cl): 17 protons, 18 neutrons (75% abundance)

• Chlorine-37 (³⁷₁₇Cl): 17 protons, 20 neutrons (25% abundance)

The existence of isotopes explains why relative atomic masses shown on the periodic table are not whole numbers—they represent weighted averages based on the abundance of each isotope.

Worked examples

Example 1: Determining particle numbers

Question: An atom of aluminium is represented as ²⁷₁₃Al. How many protons, neutrons, and electrons does this atom contain? (3 marks)

Answer:

• Number of protons = 13 (the atomic number/bottom number) ✓
• Number of electrons = 13 (same as number of protons in a neutral atom) ✓
• Number of neutrons = 27 - 13 = 14 (mass number - atomic number) ✓

Mark scheme notes: Each particle type is worth 1 mark. Students must show working for neutrons calculation or state the correct number.

Example 2: Identifying elements and compounds

Question: The table shows four substances. State which are elements and which are compounds. Explain your reasoning. (4 marks)

Answer:

• Elements: oxygen (O₂) and nitrogen (N₂) ✓
• Explanation: They contain only one type of atom ✓
• Compounds: ammonia (NH₃) and magnesium oxide (MgO) ✓
• Explanation: They contain two or more different elements chemically bonded together ✓

Mark scheme notes: 1 mark for correctly identifying each category, 1 mark for each explanation. Must mention "one type of atom" for elements and "different elements bonded/combined" for compounds.

Example 3: Writing chemical formulae

Question: A compound contains 2 aluminium atoms and 3 oxygen atoms. (a) Write the chemical formula for this compound. (1 mark) (b) Calculate the total number of atoms in one formula unit of this compound. (1 mark)

Answer: (a) Al₂O₃ ✓ (b) 2 + 3 = 5 atoms in total ✓

Mark scheme notes: In part (a), must have correct symbols with correct subscripts (Al₂O₃). Accept ratio 2:3 shown clearly. In part (b), must show working or correct answer.

Common mistakes and how to avoid them

• Confusing atoms and molecules: Remember that atoms are single particles, while molecules contain two or more atoms bonded together. O represents one oxygen atom, but oxygen gas exists as O₂ (a molecule containing two oxygen atoms).

• Mixing up atomic number and mass number: The atomic number (smaller, bottom number) tells you the number of protons. The mass number (larger, top number) tells you protons + neutrons. Always subtract atomic number from mass number to find neutrons, not the other way around.

• Thinking mixtures and compounds are the same: Compounds have elements chemically bonded in fixed ratios and require chemical reactions to separate them. Mixtures have no chemical bonds between components and can be separated by physical methods like filtration or distillation.

• Incorrect subscript placement in formulae: In H₂O, the 2 applies only to hydrogen (2 hydrogen atoms, 1 oxygen atom). Writing HO₂ would represent a completely different substance (the hydroperoxyl radical). Always place subscripts immediately after the element they apply to.

• Forgetting that isotopes have the same chemical properties: Students often think isotopes behave differently in chemical reactions. Isotopes of an element react identically because they have the same number of electrons (which determine chemical behaviour). Only physical properties differ due to mass differences.

• Not recognising diatomic elements: Remember the seven elements that exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. When writing equations or discussing these elements in their natural state, you must write them as molecules, not single atoms.

Exam technique for "Atomic structure and the periodic table: atoms, elements and compounds"

• "State" and "Give" questions: Provide a direct answer without explanation. These are typically 1-mark questions. For example, "State the number of protons in a sodium atom" requires only the answer "11"—no working needed.

• "Explain" and "Describe" questions: You must provide reasoning or detail. "Explain why sodium chloride is a compound" requires you to state that it contains different elements (sodium and chlorine) chemically bonded together in fixed proportions—not just "it's made of two elements."

• Formula questions: Write clearly with subscripts in the correct position. If handwriting, write small numbers slightly below the line to show they're subscripts. Check you've used the correct chemical symbols with proper capitalisation (Ca not ca or CA).

• Calculation questions: Always show your working, even for simple arithmetic. If you make an arithmetic error but your method is correct, you can still gain method marks. For particle number calculations, write the formula (e.g. neutrons = mass number - atomic number) before substituting numbers.

Quick revision summary

Atoms are the smallest particles of elements, consisting of protons and neutrons in a central nucleus surrounded by electrons in shells. Elements contain only one type of atom and are represented by chemical symbols. Compounds form when different elements chemically bond in fixed ratios, creating substances with new properties. Mixtures contain substances not chemically bonded, retaining individual properties. Isotopes are atoms of the same element with different neutron numbers. Chemical formulae use symbols and subscripts to show element ratios in compounds.