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Chemistry

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Atomic structure

Learning objectives · Know what is meant by the terms atom and molecule; Know the structure of an atom in terms of the positions, relative masses and relative charges of sub-atomic particles; Know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Aᵣ); Be able to calculate the relative atomic mass of an element (Aᵣ) from isotopic abundances

Atomic structure — the chapter map

SB p.24

Take a piece of copper and imagine cutting it in half, then in half again, on and on. The thought experiment ends at a definite stopping point: eventually you reach the smallest piece of copper that can exist — and at that stage you are holding a single copper atom. You could split that atom further, into protons, neutrons and electrons, but the moment you do, you no longer have copper: the pieces are not copper, they are sub-atomic particles. That is the definition of the atom — the smallest piece of an element that can still be recognised as that element.

The 118 elements are the building blocks from which everything is made, from a simple substance such as carbon to something as complex as DNA, and every one of those elements is built from atoms. What makes one element different from another is what this chapter is about: the numbers of protons, neutrons and electrons inside the atom. And the story reaches further than Earth: new atoms are produced in stars, which is why the story of the atom is also, quietly, the story of the universe.

By the end of this lesson you will be able to:

say what atom and molecule mean, in mark-scheme language;
describe the structure of the atom — where the protons, neutrons and electrons are, and their relative masses and charges;
define atomic number, mass number, isotopes and relative atomic mass (Aᵣ);
calculate a relative atomic mass from isotopic abundances, exactly as the worked example does for boron (§s08).

Key point

An atom is the smallest piece of an element that can still be recognised as that element. Split it into protons, neutrons and electrons and you no longer have the element — that splitting happens only in nuclear processes (nuclear bombs, nuclear reactors, radioactive decay), never in chemistry.

Prior knowledge

Open this chapter in the Rail All chapters

Atoms and molecules

SB p.24

Atoms rarely stay single. They can be joined together to make molecules: a molecule consists of two or more atoms chemically bonded (by covalent-bonds). One point repays care, and examiners love it: the atoms in a molecule can be the same element or different elements. A hydrogen molecule, H₂, is two hydrogen atoms bonded together — one element only, and still a molecule. A water molecule, H₂O, is two hydrogen atoms and an oxygen atom bonded together. In diagrams like these the lines between atoms represent the chemical bonds.

Splitting an atom into smaller pieces is not a chemical process at all. As §s01 established, that only happens in nuclear processes — nuclear bombs, nuclear reactors, radioactive decay. Chemistry, the whole subject you are studying, happens above that floor: it rearranges how atoms are joined, it never splits them.

One caution before we go inside the atom. You may have seen diagrams in which electrons are drawn orbiting the nucleus rather like planets around the sun. That picture is misleading. Electrons are constantly moving, and it is impossible to know exactly where an electron is at any moment; what you can say is that it has a particular energy and is likely to be found in a certain region of space at some particular distance from the nucleus. Electrons with different energies are found at different distances from the nucleus — that energy-region picture is the one to carry through the course.

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Cross-curricular links

Physics: everything in §s02's nuclear list — nuclear bombs, nuclear reactors, radioactive decay — is physics territory: chemistry cannot split the atom, and the energy side of that line belongs to your physics lessons. Stars: new atoms are produced in stars — stellar nuclear fusion is where the elements' story begins. :::key-point A molecule is two or more atoms chemically bonded (covalently). The atoms may be the same element (H₂) or different elements (H₂O). And the planet-orbit picture of electrons is misleading: electrons have particular energies and are likely to be found in certain regions of space, not on fixed tracks.

Common misconception

"A molecule must contain at least two different elements." No — H₂ is a molecule of one element only. What makes it a molecule is two or more atoms chemically bonded, not the variety of elements. (One atom alone is an atom; two bonded atoms are a molecule, same element or not.)

Exam hint

Write molecular formulae with subscripts: H₂ and H₂O, never "H2" / "H2O" — in "explain/draw a molecule" answers the subscript carries the meaning (two hydrogens in H₂O). If asked to draw the molecules, use the standard convention: lines between atoms represent the bonds.

Quick check

Question-bank join: spec 1.14 × 3 items — define molecule; is H₂ a molecule of a compound or of an element; why the planet-orbit diagram is only a model.

Key terms in this section

molecule: — Two or more atoms chemically bonded (by covalent bonds); the atoms may be of the same element (H2) or different elements (H2O).
covalent bond: — The chemical bond that joins atoms into molecules, shown as a line between atoms in the book's diagrams; detailed treatment comes in later chapters.
electron: — Sub-atomic particle outside the nucleus with relative mass 1/1836 and relative charge -1; its number equals the proton number in a neutral atom.

Keep goingAtoms: Definitions & Structure 1-3-atomic-structure

Inside the atom — protons, neutrons and electrons

SB p.25

Atoms are made of protons , neutrons and electrons — particles smaller than an atom — hence the name sub-atomic particles. The protons and neutrons sit together in the nucleus at the centre of the atom; the electrons occupy the space around it.

How small is the nucleus really? Scale a helium atom up to the size of a sports hall and the nucleus would be no more than the size of a grain of sand. Almost all of the atom is empty space — yet almost all of its mass is in that grain of sand, because electrons have a far smaller mass than protons and neutrons.

The numbers you must know are relative values — measured against each other, because the real values are incredibly small (it would take about 600 000 000 000 000 000 000 000, that is 6 × 10²³, protons to weigh 1 g). One small table carries the whole story:

Particle	Relative mass	Relative charge
proton	1	+1
neutron	1	0
electron	1/1836	−1

Protons and neutrons have (relative) mass 1 — the electron's mass, 1/1836, is so small it is usually treated as negligible when counting an atom's mass. Charges: the proton is +1, the neutron is 0 (neutral), the electron is −1.

Key point

Relative masses and charges: proton mass 1, charge +1; neutron mass 1, charge 0; electron mass 1/1836, charge −1. Protons and neutrons are in the nucleus; electrons are outside it. Virtually all the atom's mass is in the nucleus.

Common misconception

"The electron's mass of 1/1836 means a negative mass." No — 1/1836 is the relative mass (about 0.0005 of a proton's); the −1 is its charge, a separate column. Mass tiny, charge negative: keep the two properties apart and the table reads cleanly.

Exam hint

"State the relative masses and relative charges of the proton, neutron and electron" is a stock 3-mark question (spec 1.15 verbatim). Learn the table as a 2 × 3 grid and write the charge with its sign: +1, 0, −1. A missing − sign loses the mark even when the number is right.

Quick check

Question-bank join: spec 1.15 × 3 items — where each particle sits; the relative masses; the relative charges (signs included).

Key terms in this section

nucleus: — The central part of the atom containing the protons and neutrons; tiny compared with the atom but holding virtually all its mass.
proton: — Sub-atomic particle in the nucleus with relative mass 1 and relative charge +1; its number is the atomic number.
neutron: — Sub-atomic particle in the nucleus with relative mass 1 and relative charge 0; its count varies between isotopes.
electron: — Sub-atomic particle outside the nucleus with relative mass 1/1836 and relative charge -1; its number equals the proton number in a neutral atom.
sub-atomic particle: — A particle smaller than an atom; the three of record are the proton, neutron and electron.

Fill the tableblock 2

Complete the subatomic particle table (relative values).

Particle	Relative mass	Relative charge
Proton
Neutron
Electron

Keep goingAtoms: Definitions & Structure 1-3-atomic-structure 1-4-the-periodic-table

Atomic number and mass number

SB p.25

Two numbers label every atom. The number of protons in an atom's nucleus is its atomic-number (sometimes called the proton number):

atomic number = number of protons

Each of the 118 elements has its own atomic number — no two elements share it. If an atom has 8 protons, it must be an oxygen atom; the atomic number doesn't just describe the element, it defines it. A playful way to see it: we could talk about "a wristwatch made from the element with atomic number 79" instead of "a gold wristwatch" — true, but nobody talks that way, because names are easier.

The second number counts the whole nucleus. The mass-number (sometimes known as the nucleon number) is the total number of protons and neutrons:

mass number = number of protons + number of neutrons

For any particular atom the two numbers can be written as a label on the symbol — for example cobalt, written in nuclide-style notation as ⁵⁹₂₇Co. This particular cobalt atom contains 27 protons (the atomic number, bottom-left); to bring the total of protons and neutrons up to 59 (the mass number, top-left) it must also contain 32 neutrons. Rearranged, that is the counting formula you will use constantly:

number of neutrons = mass number − atomic number

One more caution, and it matters for the Periodic Table (§s10): the number above each symbol in the International GCSE Periodic Table papers is a relative atomic mass, not a mass number. In most cases it happens to equal the mass number of the most common isotope — the the only exceptions are chlorine (35.5) and copper (63.5), and §s06 explains exactly why those two break the pattern.

Key point

atomic number = number of protons (defines the element). mass number = protons + neutrons. neutrons = mass number − atomic number. In nuclide notation (⁵⁹₂₇Co) the top-left number is the mass number, the bottom-left the atomic number.

Common misconception

"The bigger number printed on the Periodic Table is the mass number." Not always — on exam-paper Periodic Tables that number is the relative atomic mass. It matches the most common isotope's mass number for most elements, but chlorine (35.5) and copper (63.5) prove it is not a mass number: no atom contains half a neutron.

Exam hint

Symbols with two letters: first capital, second lower-case. The classic trap is CO versus Co — "CO" is carbon monoxide (a compound), "Co" is cobalt (an element). Case is not decoration in chemistry; it is the difference between an element and a compound.

Quick check

Question-bank join: spec 1.16 × 3 items — define atomic number; define mass number; neutrons from mass number − atomic number (cobalt 59/27). spec 1.15 × 1 item — which particles the two numbers count.

Key terms in this section

atomic number: — The number of protons in the nucleus of an atom; it defines the element and every atom of the same element shares it.
mass number: — The total number of protons and neutrons in the nucleus of an atom.
relative atomic mass: — The average mass of an atom of an element, taking into account the amount of each isotope in a naturally occurring sample; a weighted average, usually not a whole number.

Keep goingAtoms: Definitions & Structure 1-3-atomic-structure 1-4-the-periodic-table 1-8-metallic-bonding

Isotopes — same element, different mass

SB p.26

The number of neutrons in an atom can vary slightly. Carbon is the classic example: there are three kinds of carbon atom, called carbon-12, carbon-13 and carbon-14. All of them have 6 protons — every carbon atom does, because 6 is carbon's atomic number — but their neutron counts differ, so their mass numbers differ too. These different atoms of the same element are called isotopes:

Isotopes are atoms (of the same element) which have the same atomic number but different mass numbers — the same number of protons but different numbers of neutrons.

Does the isotope change how the element behaves? No — and the reason is worth memorising because it is the bridge to the rest of chemistry. The fact that isotopes have varying numbers of neutrons makes no difference to their chemical reactions: chemical properties (how something reacts) are controlled by the number and arrangement of the electrons, and that is identical for all three carbon isotopes. Carbon-12 and carbon-14 are chemically the same element; where they differ is mass (and, for some isotopes, nuclear stability — a nuclear property, not a chemical one).

Key point

Isotopes = same atomic number (same protons), different mass numbers (different neutrons). Chemistry is unchanged between isotopes because chemical reactions depend on the electrons, which are identical.

Common misconception

"Isotopes are the radioactive versions of an element." No — isotope just means same protons, different neutrons. Carbon-12 and carbon-13 are stable isotopes; radioactivity is a separate nuclear property that only some isotopes (like carbon-14) have. Don't fold nuclear behaviour into a definition that is only about particle counts.

Exam hint

"Explain what isotopes are" scores with the paired wording: same number of protons (same atomic number), different numbers of neutrons (different mass numbers). Both halves must appear — protons alone, or neutrons alone, is half an answer and half a mark. And "explain the difference between two isotopes" wants the neutrons/mass number, never "one is more reactive".

Quick check

Question-bank join: spec 1.16 × 2 items — define isotopes; why ¹²C, ¹³C and ¹⁴C react identically.

Key terms in this section

isotope: — Atoms of the same element with the same atomic number but different mass numbers - same protons, different numbers of neutrons; chemically identical.

Keep goingAtoms: Definitions & Structure 1-3-atomic-structure 1-8-metallic-bonding

Relative atomic mass — the 35.5 puzzle

SB p.26

Look up chlorine on the periodic-table and you meet a paradox. The number beside chlorine appears to be a mass number of 35.5. Try the §s04 formula on it:

number of neutrons = 35.5 − 17 = 18.5

It is not possible to have half a neutron — so something must be wrong with this. And something is: 35.5 is not a mass number at all. It is the relative-atomic-mass (Aᵣ) of chlorine. Chlorine consists of two isotopes, ³⁵Cl and ³⁷Cl, and a naturally occurring sample contains a mixture of these — so the 35.5 is an average across the whole sample, not the mass of any single atom.

A plain average would not do, because the isotopes are not present in equal amounts. The tool for this is the same one you use in maths when some values count for more than others: the weighted-average (or weighted mean). The relative atomic mass is a weighted average of the isotope masses, weighted by how much of each isotope is present in a naturally occurring sample. That single sentence is the definition behind spec 1.16's last term — and §s07 and §s08 turn it into a calculation you can perform.

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Cross-curricular links

Maths: the weighted mean is a taught GCSE maths skill — values scaled by weights before averaging. Aᵣ is that exact calculation applied to isotope masses and abundances; the maths department's "weighted average" and chemistry's "relative atomic mass" are one and the same operation. :::key-point 35.5 for chlorine is a relative atomic mass (Aᵣ), not a mass number — no atom has half a neutron. Aᵣ is a weighted average of the isotope masses across a naturally occurring sample; chlorine's sample is a mixture of ³⁵Cl and ³⁷Cl.

Common misconception

"Chlorine's mass number is 35.5, so a chlorine atom has 18.5 neutrons." Twice wrong — no chlorine atom has 18.5 neutrons (impossible), and 35.5 is not any atom's mass number. Real chlorine atoms are ³⁵Cl (18 neutrons) or ³⁷Cl (20 neutrons); the 35.5 is the sample's weighted average.

Exam hint

If an exam gives a non-integer "mass number" (35.5, 63.5), it is telling you the quantity is a relative atomic mass — say so explicitly. Naming the quantity correctly ("Aᵣ, not mass number") is usually a mark in itself on questions of this type.

Quick check

Question-bank join: spec 1.16 × 3 items — why 35.5 cannot be a mass number; what Aᵣ is an average of; why a plain average would not do.

Key terms in this section

relative atomic mass: — The average mass of an atom of an element, taking into account the amount of each isotope in a naturally occurring sample; a weighted average, usually not a whole number.
weighted average: — An average in which each value is multiplied by its weight before averaging; Ar weights isotope masses by their abundances.
periodic table: — The chart of all the elements, arranged in order of increasing atomic number; each cell carries the atomic number plus a mass number or relative atomic mass.

Keep goingAtoms: Definitions & Structure 1-3-atomic-structure 1-8-metallic-bonding

Counting the number of electrons in an atom

SB p.27

Here is the question you can now almost answer for yourself: given an element, how do you count the electrons in its atom? The full charge-balance argument arrives in §s09 — the anchor that everything else hangs from comes here: the number of protons is the same as the atomic number of the element. Count protons by looking up the atomic number; electrons then follow (§s09); neutrons come from mass number − atomic number (§s04). One look-up, three particle counts.

The hint box also closes the chlorine story from §s06, and it is worth reading slowly because it is examiner-true reasoning. Relative atomic mass is defined as the average mass of an atom, taking into account the amount of each isotope present in a naturally occurring sample of an element (Chapter 5 builds on this further). Apply that to chlorine: the Aᵣ of 35.5 is closer to 35 than to 37, so a naturally occurring sample of chlorine must contain more ³⁵Cl than ³⁷Cl — the average sits nearer the isotope that contributes more.

That "amount of each isotope" is the isotopic-abundance — the percentage of each isotope in the sample. Knowing the abundances, you can stop reasoning about the average and calculate it, in exactly the way you would calculate a weighted average in maths — which is exactly what the next section's worked example does.

Key point

Protons = atomic number (one Periodic-Table look-up powers all three particle counts). Aᵣ = the average mass of an atom taking the isotopic abundances into account; the average lies closest to the most abundant isotope.

Common misconception

"Chlorine's Aᵣ of 35.5 means the sample is mostly ³⁷Cl, since 37 is the bigger isotope." No — the weighted average is closer to 35, which is only possible if ³⁵Cl is the more abundant isotope. Distance from the average tells you which isotope dominates; bigger isotope ≠ bigger share.

Exam hint

"Explain why a sample of chlorine contains more ³⁵Cl than ³⁷Cl" scores by one line: the relative atomic mass (35.5) is closer to 35 than to 37. Quote the closeness argument — it is the whole mark.

Quick check

Question-bank join: spec 1.17 × 2 items — define isotopic abundance; which chlorine isotope dominates and why. spec 1.16 × 1 item — the Aᵣ definition (average, abundances taken into account).

Key terms in this section

isotopic abundance: — The amount (percentage) of each isotope present in a naturally occurring sample of an element.

Step throughblock 3

Chlorine: 75% is chlorine-35, 25% is chlorine-37. Find Ar of chlorine.

1•Contribution of Cl-35 (abundance x mass number)
2•Contribution of Cl-37
3•Divide by the total abundance (100)

Examiner note - mass language

'Relative atomic mass' and 'mass number' are NOT interchangeable. Mass number belongs to ONE isotope (protons + neutrons, a whole number); Ar is the weighted mean for the ELEMENT. In 'explain why Ar of chlorine is 35.5', the marking points are: chlorine exists as isotopes AND abundance-weighted mean - '35.5 neutrons' scores nothing.

Watchblock 8

GCSE Chemistry - Elements, Isotopes & Relative Atomic Mass (2027/28 exams)

Isotopes and relative atomic mass — abundance-weighted Ar calculations, worked step by step.

Watch on YouTube

Video © Cognito (cognito.org) — embedded via YouTube

Keep goingAtoms: Definitions & Structure Calculate Relative Atomic Mass 1-3-atomic-structure 1-8-metallic-bonding 2-2-group-7-halogens

Worked example — the relative atomic mass of boron

SB p.27

This worked example turns §s07's definition into a routine. A naturally occurring sample of boron contains 20% ¹⁰B and 80% ¹¹B. Calculate the relative atomic mass.

The method — imagine there are exactly 100 atoms — makes the percentages do the work:

20% of 100 atoms = 20 atoms, each of mass 10 → total mass 20 × 10
80% of 100 atoms = 80 atoms, each of mass 11 → total mass 80 × 11

Add the masses, divide by the number of atoms, and the average falls out:

relative atomic mass = (20 × 10 + 80 × 11) ÷ 100 = (200 + 880) ÷ 100 = 10.8

Therefore the relative atomic mass of boron is 10.8. Notice what the number tells you before you check it: 10.8 sits much closer to 11 than to 10, exactly as §s07 predicts for a sample that is 80% ¹¹B.

The same method works even if there are three or four different isotopes: multiply each isotope's mass by its abundance, add all the products, divide by 100 (when abundances are percentages). The imagine-100-atoms framing survives every exam version of this question.

Worked example

Boron: 20% ¹⁰B, 80% ¹¹B. Imagine 100 atoms → 20 atoms of mass 10, 80 of mass 11. Total mass = 20 × 10 + 80 × 11 = 1080. Divide by 100 atoms: Aᵣ = 10.8. Mark lines: (1) multiply each mass by its abundance; (2) add the products; (3) divide by 100; (4) quote 10.8 (no units — Aᵣ is a relative quantity).

Maths skills

The skill is the weighted mean, in the exact form maths uses: multiply each value by its weight, add, divide by the total weight. Here the values are isotope masses (10, 11) and the weights are the abundances (20, 80). The imagine-100 trick converts percentages into counts so you never have to manipulate percentage fractions under exam pressure.

Common misconception

"The relative atomic mass of boron should be a whole number, like a mass number." No — Aᵣ is an average over a mixture, so it is usually not a whole number (boron 10.8, chlorine 35.5, copper 63.5). Whole numbers belong to individual isotopes (mass numbers 10 and 11); the element's Aᵣ averages across them.

Exam hint

Show the full weighted-mean line — (20 × 10 + 80 × 11) ÷ 100 — not just the answer. The three method steps carry the marks; a bare "10.8" with no working risks losing method marks if any slip occurs. Quote the final value without units.

Quick check

Question-bank join: spec 1.17 × 3 items — reproduce the boron calculation; a two-isotope Aᵣ with different abundances; the three-isotope generalisation.

Keep goingCalculate Relative Atomic Mass 1-3-atomic-structure 2-2-group-7-halogens

The electrons

SB p.27

One charge argument finishes the particle census. The charge on a proton (+1) is equal but opposite to the charge on an electron (−1), and an atom is neutral — so the two charges must cancel exactly:

number of electrons = number of protons

Two checks: an oxygen atom (atomic number 8) has 8 protons, so it must also have 8 electrons; a chlorine atom (atomic number 17) has 17 protons and therefore 17 electrons. The key feature throughout is knowing the atomic number — one look-up on the Periodic Table (§s10) and all three counts follow: protons from the atomic number, electrons from the protons, neutrons from mass number − atomic number.

Why, then, is the atomic number defined in terms of protons rather than electrons, when either count starts from the other in a neutral atom? The answer matters for the whole course: the number of electrons can change in chemical reactions, for example when atoms form ions (Chapter 7) — but the number of protons in the nucleus never changes chemically. Protons are the stable identity; electrons are the reactive surface. That is why "number of protons = atomic number" is the definition of record, and "electrons = protons" is merely true for neutral atoms.

Key point

Charge balance: proton +1 vs electron −1, equal and opposite, so in an atom electrons = protons. The atomic number is defined via protons because protons never change chemically, while electrons can (ions, Chapter 7).

Common misconception

"An oxygen atom has 8 protons and 10 electrons, like its mass number suggests." No — mass number counts protons + neutrons, never electrons. A neutral oxygen atom: 8 protons, 8 electrons, and (for ¹⁶O) 8 neutrons. Electrons follow the atomic number, not the mass number.

Exam hint

"Explain why the number of protons in an atom must always equal the number of electrons" is a stock end-of-chapter part (d) — and the marks are: charges equal and opposite (+1/−1) and atoms are neutral, so the charges must cancel. Both halves, in that order.

Quick check

Question-bank join: spec 1.15 × 2 items — the charge-cancellation argument; count all three particles for oxygen. spec 1.16 × 1 item — why the atomic number is defined via protons.

Key terms in this section

ion: — What an atom becomes when it gains or loses electrons, so its particle counts stop matching; the book develops ions fully in Chapter 7.

Keep goingAtoms: Definitions & Structure 1-3-atomic-structure 1-4-the-periodic-table 1-8-metallic-bonding

The Periodic Table — counting particles from the chart

SB p.28

Everything in this chapter runs on one reference chart. Atoms are arranged in the periodic-table in order of increasing atomic number — that ordering is the chapter's payoff: the chart is a particle-counting device. (Chapter 4 deals in detail with everything else the Periodic Table tells you; the Periodic Table in your exam paper carries the full version.)

Most Periodic Tables show two numbers against each symbol, and the warning is blunt: be careful to choose the right one. The atomic number will always be the smaller number. The other number is either the mass number of the most common isotope of the element or the relative atomic mass of the element (§s06 showed you why those are not the same thing — chlorine's 35.5). The Periodic Table you are using will clarify which; exam papers always state their convention.

The complete counting recipe for any atom:

the number of protons = the atomic number (the smaller number);
the number of electrons = the number of protons;
the number of neutrons = mass number − atomic number.

Three lines, one chart, every atom on the syllabus. §s11 works the recipe end-to-end on uranium.

Key point

Periodic Table recipe: protons = atomic number (always the smaller number); electrons = protons; neutrons = mass number − atomic number. Watch the second number in the cell: mass number (most common isotope) or relative atomic mass — they are not the same quantity.

Common misconception

"The two numbers in a Periodic Table cell are the mass number and the atomic number." Only sometimes — the larger number is the relative atomic mass on exam papers (chlorine 35.5, copper 63.5), which equals the most common isotope's mass number for most elements but not all. Read the paper's convention; do not assume.

Exam hint

When a question says "use the Periodic Table", it is licensing the three-line recipe. Write the atomic number down first (from the chart), then run the recipe in order: protons → electrons → neutrons. Stating the atomic number explicitly protects the whole chain of marks.

Quick check

Question-bank join: spec 1.16 × 3 items — which number is which in a cell; the three-line recipe; apply it to a named element.

Key terms in this section

periodic table: — The chart of all the elements, arranged in order of increasing atomic number; each cell carries the atomic number plus a mass number or relative atomic mass.

Keep goingAtoms: Definitions & Structure 1-3-atomic-structure 1-8-metallic-bonding

Worked example — uranium, and the chapter close

SB p.28

This worked example runs the §s10 recipe on a heavy atom. The symbol for uranium is given in a Periodic Table as ²³⁸U (uranium, atomic number 92). Calculate the number of protons, neutrons and electrons in an atom of uranium.

The atomic number is the smaller number, so uranium's atomic number is 92 → a uranium atom contains 92 protons.
Protons = electrons → 92 electrons.
Neutrons = mass number − atomic number = 238 − 92 = 146.

Same three lines as always — the atom just gets heavier. (For this you will need the Periodic Table in your exam paper.)

The chapter-closing skill list is worth self-checking: particle locations and charges; counting the particles in a fluorine atom; reading nuclide-style symbols like ⁵⁶Fe and ⁹³₄₁Nb; chlorine's two isotopes; and a run of relative-atomic-mass calculations — lithium, magnesium with three isotopes, and lead with four. This lesson's practice is wired to the question bank below.

One last fact for the road: considering only the most common isotope of each element, there is only one element that has more protons than neutrons. Work it out with the recipe — the answer is hydrogen, whose most common atom (¹H) has 1 proton and 0 neutrons; every other element's most common isotope carries at least as many neutrons as protons.

Worked example

Uranium ²³⁸U, atomic number 92: atomic number = smaller number = 92 → 92 protons; electrons = protons → 92 electrons; neutrons = 238 − 92 = 146. Mark lines: (1) identify the atomic number as the smaller number; (2) protons from atomic number; (3) electrons = protons; (4) subtract for neutrons.

Key point

The counting recipe is atom-independent: ²³⁸U → 92 protons, 92 electrons, 146 neutrons. And a chapter-closing curiosity — hydrogen-1 is the only most-common-isotope atom with more protons than neutrons (1 proton, 0 neutrons).

Common misconception

"Heavy atoms like uranium have far more electrons than protons." No — charge balance holds for every atom: 92 protons, 92 electrons. What grows faster in heavy atoms is the neutron count (146 here), which is exactly why neutrons = mass number − atomic number is the line you keep computing.

Exam hint

For "state the numbers of protons, neutrons and electrons" answers, present all three counts in order (protons, electrons, neutrons) with the substitution shown for neutrons (238 − 92 = 146). The substitution line is the method mark; the three numbers are the answer marks.

Quick check

Question-bank join: spec 1.16 × 2 items — reproduce the uranium counts; the hydrogen-1 exception. spec 1.17 × 1 item — the lead four-isotope weighted mean.

Summary

An atom is the smallest piece of an element still recognisable as that element; a molecule is two or more atoms chemically bonded (H₂, H₂O). Only nuclear processes split atoms — chemistry never does.
Atoms contain protons (mass 1, charge +1) and neutrons (mass 1, charge 0) in the nucleus, with electrons (mass 1/1836, charge −1) outside; virtually all the mass is nuclear.
Atomic number = protons (defines the element); mass number = protons + neutrons; neutrons = mass number − atomic number; electrons = protons (neutral atoms).
Isotopes: same atomic number, different mass numbers — chemically identical, because chemistry belongs to the electrons.
Aᵣ is the weighted average of isotope masses across a natural sample (chlorine 35.5: mostly ³⁵Cl; boron 10.8 from 20%/80%).
Periodic Table: ordered by increasing atomic number; smaller number = atomic number; run the three-line recipe for any atom.

Can you…?

define atom and molecule, and say why H₂ is a molecule of an element?
state the relative masses and charges of the proton, neutron and electron (signs included)?
define atomic number and mass number, and find neutrons from mass number − atomic number (cobalt 59/27)?
define isotopes and explain why ¹²C, ¹³C and ¹⁴C react identically?
explain why chlorine's 35.5 is a relative atomic mass and not a mass number?
deduce which isotope dominates a sample from how close Aᵣ lies to each isotope's mass?
calculate a relative atomic mass from isotopic abundances (boron 20%/80% → 10.8), including three- and four-isotope cases?
take any element (say uranium ²³⁸U) and state its protons, electrons and neutrons from the Periodic Table?

Keep goingAtoms: Definitions & Structure Calculate Relative Atomic Mass 1-3-atomic-structure 1-8-metallic-bonding 2-2-group-7-halogens