1.2.2 Atomic Structure — OCR Gateway GCSE Combined Science A J250 - Chemistry Revision Notes & Topical Questions

Part 1 of 7Key ideas

1

Atoms are neutral because the number of protons equals the number of electrons.

2

The nucleus is tiny compared with the whole atom but contains almost all the mass.

3Electron shells fill in order

2, 8, 8 for the first 20 elements (e.g. Ca 2.8.8.2).

4

Isotopes have identical chemical properties because they have the same electron configuration.

5

A_r = Σ(isotope mass × % abundance) ÷ 100.

6Chlorine

75% ³⁵Cl and 25% ³⁷Cl gives A_r = 35.5.

7

Relative atomic mass is measured relative to carbon-12, which is given a mass of exactly 12.

8For a neutral atom

For a neutral atom electrons = protons, and neutrons = nucleon number − proton number.

9

Shells fill from the one nearest the nucleus outwards: the first holds 2, the second and third 8 each, and the fourth only starts filling once the third has 8.

10The configuration gives the position directly

The configuration gives the position directly the number of shells is the period, and the number of outer electrons is the group (for the main groups).

11Same

Isotopes of an element have the same number of protons and electrons and a different number of neutrons, so their chemistry is identical and only their mass-related properties differ.

12No units

Relative atomic mass is an average over the isotopes present, compared with one twelfth of the mass of a carbon-12 atom — and it has no units.

Part 2 of 7Learning objectives

1

Describe the structure of an atom and the properties of its particles.

2

Use atomic and mass numbers to find numbers of protons, neutrons and electrons.

3

Define isotopes and explain their chemical similarity.

4

Calculate relative atomic mass from isotopic abundances.

Part 3 of 7Essential knowledge

Essential knowledge

Atoms and molecules

Atoms and molecules an atom is the smallest particle of an element; a molecule is a group of atoms bonded together, which may be of one element (O₂, Cl₂) or of a compound (H₂O, CH₄).

Structure

Structure at the centre is a very small, dense, positively charged nucleus containing protons and neutrons. Electrons occupy shells (energy levels) around it. Most of an atom's volume is empty space.

Atomic number

The atomic number identifies the element: every carbon atom has 6 protons. The mass number is protons + neutrons. Notation: mass number top-left, atomic number bottom-left of the symbol.

Electronic configuration

Electronic configuration first shell holds up to 2 electrons, second up to 8, third up to 8 (for the first 20 elements). The number of outer electrons equals the group number; the number of occupied shells equals the period.

Isotopes

Isotopes of an element have the same atomic number but different mass numbers, e.g. carbon-12 and carbon-14. Because chemistry depends on electrons, isotopes react the same way but have slightly different physical properties such as density.

Relative atomic mass

Relative atomic mass is not a whole number for many elements because it is the average of the isotopes weighted by their abundance. It is compared to 1/12 the mass of a ¹²C atom, so it has no units.

When atoms gain or lose electrons they become ions, but the number of protons (and so the element) does not change.

Part 4 of 7Process & sequence

1

Write down each isotope's mass number and its percentage abundance.

2

Multiply each mass number by its percentage abundance.

3

Add the products together.

4

Divide the total by 100 (the sum of the percentages).

5

Round sensibly (usually to 1 decimal place) and give no units.

Calculating A_rwrite these steps in this order; each step is a marking point.

1

List isotope masses

2

Note % abundances

3

Multiply mass by abundance

4

Add the products

5

Divide by 100

Part 5 of 7Worked examples

163.6

Copper is 69.2% ⁶³Cu and 30.8% ⁶⁵Cu. Calculate its relative atomic mass.

A_r = (63 × 69.2 + 65 × 30.8) ÷ 100 = (4359.6 + 2002.0) ÷ 100 = 6361.6 ÷ 100 = 63.6.

2How many protons, neutrons and electrons are in 3216S2−?

Protons = 16; neutrons = 32 − 16 = 16; electrons = 16 + 2 = 18 (the 2− charge means two extra electrons).
Counting the particles
- An atom of copper has proton number 29 and nucleon number 63.
- Protons = 29; electrons = 29 (neutral atom); neutrons = 63 − 29 = 34.
From configuration to position in the table
- Chlorine is 2,8,7: three shells → Period 3; seven outer electrons → Group VII.
- Sodium is 2,8,1: three shells → Period 3; one outer electron → Group I.
- Argon is 2,8,8: a full outer shell, which is why it sits in Group VIII and does not react.
Relative atomic mass of rubidium
- Rubidium is 72 % mass-85 and 28 % mass-87.
- A_r = (72 × 85 + 28 × 87) ÷ 100 = 85.6.
- The answer is written without units, and it lies between the two isotope masses, closer to the more abundant one — a useful sanity check.

Part 6 of 7Summary

1

Atoms have a nucleus of protons and neutrons surrounded by electrons.

2

Atomic number = protons; mass number = protons + neutrons.

3

Isotopes differ only in neutron number.

4

A_r is a weighted mean of isotope masses.

5

Neutrons = nucleon number − proton number; electrons = protons in a neutral atom.

6

Shells fill 2, 8, 8 then the fourth shell begins.

7

From a configuration: shells → period, outer electrons → group.

8

Isotopes differ in neutrons only, so their chemical properties are the same.

Part 7 of 7Compare: Compare

Sub-atomic particles
Particle	Relative mass	Relative charge	Location
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus
Electron	1/1836 (negligible)	−1	Shells around nucleus

Working out particle numbers
Atom / ion	Protons	Neutrons	Electrons
²³₁₁Na	11	12	11
³⁵₁₇Cl	17	18	17
³⁷₁₇Cl	17	20	17
²⁴₁₂Mg²⁺	12	12	10

Neutrons = mass number − atomic number. Ions: subtract electrons for positive charge, add for negative.

Key terms

Command terms

Give / State / Name

Recall a fact; one idea per mark.

Define

Give the exact meaning using the specification keywords.

Describe

Give an account of what happens / the pattern in data, quoting values; no reasons needed.

Explain

Give reasons with linked statements (cause → effect). Each linked step is usually a separate marking point.

Suggest

Use knowledge in a new context; more than one answer may be accepted.

Compare

Similarities and differences; each point must refer to both items.

Calculate

Show working; include units; check significant figures. Method marks are given even if the final answer is wrong.

Discuss

Explore the issue from different angles and link the points.

Evaluate / Assess

Weigh strengths and weaknesses and reach a supported judgement.

Deduce

Draw a conclusion from the information given.

Justify

Support a statement/choice with evidence.

Exam technique

Keywords the mark scheme looks for in this lesson

Atom · Molecule · Proton · Neutron · Electron · Atomic number (Z) · Mass number (A) · Isotopes · Relative atomic mass (A_r)

Learn each definition in the Key words tab exactly; a missing keyword in a definition usually loses the mark.

How to get full marks (examiner guidance)

1

Use specification keywords exactly (e.g. electronegativity, nucleophile, dynamic equilibrium, activation energy).

2Link ideas

Link ideas for “explain” each marking point is usually a cause → consequence step; unlinked facts often score nothing.

3Use the information in the question

“discuss/evaluate” answers must use the data provided and match the number of marks.

4

Convert units carefully (mg → g, cm³ → dm³) and show working for calculations.

5Avoid vague words

“it”, “stuff”, “things”, “goes up”, “energy is made”. Name the substance and the process.

6

Do not extrapolate a graph line beyond the data or to zero unless asked.

7

Six-mark extended responses are marked on content AND a logical, linked line of reasoning.

8Definitions that are marked word for word

Isotopes"atoms of the same element with the same number of protons and different numbers of neutrons" — "different mass" alone is not enough.
Relative atomic massinclude the comparison with one twelfth of carbon-12 if the question asks for the definition.
Never give a relative atomic mass with a unit such as g or kg.