Pearson Edexcel International GCSE Chemistry 4CH1: Atomic structure and the Periodic Table – Study Guide

This study guide teaches sub-topics 1(c) Atomic structure and 1(d) The Periodic Table from Topic 1, Principles of chemistry, of the Pearson Edexcel International GCSE Chemistry (4CH1) specification, Issue 3 (September 2024). It covers specification points 1.14–1.24. The qualification is untiered, and none of these points carries a “C” reference, so all of them can be examined on Paper 1 and on Paper 2. Paper 1 lasts 2 hours (110 marks) and Paper 2 lasts 1 hour 15 minutes (70 marks); a calculator may be used on both.

Use it with the 4CH1 Chemistry course hub and the printable 4CH1 checklist. Then condense it with the atomic structure and Periodic Table revision notes and test yourself with the atomic structure and Periodic Table practice questions. To find your weak spots across the whole course, take the free 4CH1 all-topics diagnostic. For all of Topic 1, see the Principles of chemistry study guide.

What these sub-topics cover

4CH1 point	What you must be able to do
1.14	Say what is meant by an atom and a molecule
1.15	Give the positions, relative masses and relative charges of protons, neutrons and electrons
1.16	Say what is meant by atomic number, mass number, isotopes and relative atomic mass (Ar)
1.17	Calculate Ar from isotopic abundances
1.18	Explain how the Periodic Table is arranged: by atomic number, in groups and periods
1.19	Work out the electronic configuration of the first 20 elements from their positions
1.20	Use electrical conductivity and the acid–base character of oxides to classify elements as metals or non-metals
1.21	Identify an element as a metal or non-metal from its position
1.22	Link the electronic configuration of a main group element to its position
1.23	Explain why elements in the same group have similar chemical properties
1.24	Explain why the noble gases (Group 0) do not readily react

1(c) Atomic structure

Atoms and molecules (1.14)

An atom is the smallest particle of an element that keeps the chemical properties of that element. A molecule is two or more atoms chemically bonded together (by covalent bonds). The atoms can be the same element, as in O₂ or Cl₂, or different elements, as in H₂O or CO₂. So oxygen gas is made of molecules, but helium and neon exist as single atoms. A molecule containing two or more elements is a compound.

Sub-atomic particles (1.15)

An atom has a tiny central nucleus that contains protons and neutrons. Electrons occupy shells (energy levels) around the nucleus.

Particle	Position	Relative mass	Relative charge
Proton	Nucleus	1	+1
Neutron	Nucleus	1	0
Electron	Shells around the nucleus	about 1/1840 (negligible)	−1

Almost all of an atom’s mass is in the nucleus. An atom has no overall charge because it has equal numbers of protons and electrons.

Atomic number, mass number and isotopes (1.16)

• Atomic number (proton number): the number of protons in the nucleus of an atom. It identifies the element. In a neutral atom it also equals the number of electrons.
• Mass number: the total number of protons and neutrons in the nucleus of an atom.
• Isotopes: atoms of the same element that have the same atomic number but different mass numbers, because they have different numbers of neutrons.
• Relative atomic mass (Ar): the weighted mean mass of the atoms of an element, compared with one-twelfth of the mass of an atom of carbon-12.

An atom is written with its mass number at the top left and its atomic number at the bottom left of the symbol, for example ³⁹₁₉K. The number of neutrons is mass number minus atomic number.

Worked example 1 — counting particles. State the numbers of protons, neutrons and electrons in ³⁹₁₉K, ³¹₁₅P and the ion ²⁴₁₂Mg²⁺.

39K:    protons = 19   neutrons = 39 - 19 = 20   electrons = 19
31P:    protons = 15   neutrons = 31 - 15 = 16   electrons = 15
24Mg2+: protons = 12   neutrons = 24 - 12 = 12   electrons = 12 - 2 = 10

An ion has the same nucleus as its atom; a 2+ ion has lost two electrons.

Isotopes of an element have the same electron arrangement, so they react in the same way.

Calculating relative atomic mass (1.17)

Most elements are a mixture of isotopes. Ar is a weighted mean: each isotope counts in proportion to its abundance, so Ar is rarely a whole number.

Ar = sum of (mass number x percentage abundance) / 100

If abundances are given as a ratio or as relative amounts that do not add to 100, divide by their total instead of by 100.

Worked example 2 — two isotopes. A sample of gallium contains 60.1% gallium-69 and 39.9% gallium-71. Calculate the relative atomic mass of gallium. Give your answer to 1 decimal place.

Ar = (69 x 60.1 + 71 x 39.9) / 100
   = (4146.9 + 2832.9) / 100
   = 6979.8 / 100
   = 69.798
   = 69.8 (1 d.p.)

Check: the answer is closer to 69, the more common isotope.

Worked example 3 — three isotopes. A sample of magnesium contains 79% magnesium-24, 10% magnesium-25 and 11% magnesium-26. Calculate Ar to 1 decimal place.

Ar = (24 x 79 + 25 x 10 + 26 x 11) / 100
   = (1896 + 250 + 286) / 100
   = 2432 / 100
   = 24.32
   = 24.3 (1 d.p.)

Worked example 4 — abundances as a ratio. Boron-10 and boron-11 occur in the ratio 1 : 4. Calculate Ar of boron.

total parts = 1 + 4 = 5
Ar = (10 x 1 + 11 x 4) / 5
   = (10 + 44) / 5
   = 54 / 5
   = 10.8

Ar has no units: it is a relative value.

1(d) The Periodic Table

How the table is arranged (1.18)

The elements are arranged in order of increasing atomic number. Each element has one more proton than the element before it.

• A group is a vertical column. Elements in the same group have similar chemical properties. The main groups are numbered 1 to 7, with the noble gases in Group 0 on the far right.
• A period is a horizontal row. Period 1 contains only hydrogen and helium.

Ordering by atomic number, not Ar, puts each element with similar elements. For example, argon (atomic number 18) comes before potassium (atomic number 19), even though argon has the larger Ar. This keeps argon with the noble gases and potassium in Group 1.

Electronic configurations of the first 20 elements (1.19)

Electrons fill shells from the nucleus outwards. For the first 20 elements the first shell holds up to 2 electrons, the second up to 8 and the third up to 8. The fourth shell starts with potassium and calcium. A configuration is written as numbers separated by commas, innermost shell first.

Period	Elements and configurations
1	H 1; He 2
2	Li 2,1; Be 2,2; B 2,3; C 2,4; N 2,5; O 2,6; F 2,7; Ne 2,8
3	Na 2,8,1; Mg 2,8,2; Al 2,8,3; Si 2,8,4; P 2,8,5; S 2,8,6; Cl 2,8,7; Ar 2,8,8
4 (first two only)	K 2,8,8,1; Ca 2,8,8,2

You can build any of these from the element’s position, as the next section shows.

Configuration and position (1.22)

For a main group element:

• Period number = number of occupied electron shells.
• Group number = number of electrons in the outer shell (Groups 1 to 7).
• Group 0 elements have a full outer shell: 2 for helium, 8 for neon and argon.

Worked example 5 — configuration from position. An element is in Period 3 and Group 6. Deduce its electronic configuration and name it.

Period 3      -> three occupied shells
first two shells are full -> 2, 8
Group 6       -> 6 electrons in the outer shell
configuration = 2,8,6
total electrons = 16, so atomic number = 16 -> sulfur

Worked example 6 — position from configuration. An atom has the configuration 2,8,8,2. State its group and period and name it.

four shells       -> Period 4
2 outer electrons -> Group 2
2 + 8 + 8 + 2 = 20 electrons -> atomic number 20 -> calcium

Metals and non-metals by position (1.21)

Metals are on the left of the Periodic Table and in the centre block between Groups 2 and 3. Non-metals are on the right. A stepped (staircase) line separating the two runs from boron down towards the bottom right. Hydrogen is a non-metal, even though it is placed at the top of the table.

Among the first 20 elements, the metals are Li, Be, Na, Mg, Al, K and Ca.

Metals and non-metals from evidence (1.20)

The specification asks you to classify an element from two kinds of evidence.

Electrical conductivity. Test the solid element in a circuit with a bulb or ammeter.

• Metals conduct electricity when solid.
• Non-metals generally do not. Graphite, a form of carbon, is the important exception: it is a non-metal that conducts.

Acid–base character of the oxide. Burn the element in oxygen to make its oxide, then test the oxide.

• Metal oxides are basic. They react with acids to form a salt and water. A metal oxide that dissolves in water gives an alkaline solution (pH above 7), such as sodium oxide forming sodium hydroxide.
• Non-metal oxides are generally acidic. Those that dissolve in water give an acidic solution (pH below 7), such as sulfur dioxide and carbon dioxide.

Worked example 7 — classifying from data. Element E is a dull yellow solid. It does not conduct electricity. It burns to form an oxide that dissolves in water to give a solution of pH 3. Classify E.

does not conduct            -> non-metal property
oxide gives pH 3 (acidic)   -> acidic oxide -> non-metal property
conclusion: E is a non-metal

Use both pieces of evidence: graphite shows one alone can mislead.

Why groups share properties (1.23)

Chemical reactions involve the outer-shell electrons: atoms lose, gain or share them. Elements in the same group have the same number of outer-shell electrons, so they react in similar ways. For example, lithium, sodium and potassium each have one outer electron, and each reacts with water to form a hydroxide and hydrogen. Group trends are taught in Topic 2; see the reactivity and group trends guide.

Why the noble gases are unreactive (1.24)

The noble gases have a full outer shell of electrons (2 for helium; 8 for the others). This is a stable arrangement, so a noble gas atom has no tendency to lose, gain or share electrons. That is why they do not readily react and why they exist as single atoms rather than molecules. Helium is in Group 0 even though it has only 2 outer electrons, because its outer shell (the first shell) is full.

Common errors

• Writing that the electron has “no mass”. Its mass is very small (negligible), not zero.
• Defining isotopes as “different numbers of neutrons” without saying “same element” or “same number of protons”.
• Saying isotopes have different chemical properties.
• Dividing by 100 when abundances are given as a ratio that does not total 100.
• Writing 2,8,9 for potassium. The third shell holds 8 before the fourth starts.
• Saying the noble gases are unreactive “because they have 8 outer electrons”. Helium has 2. The reason is a full outer shell.

Where to go next

Condense this page with the revision notes, then try the practice questions. For wider Topic 1 recall use the Principles of chemistry revision notes and practice questions. Ar values feed straight into mole work in the formulae, equations and calculations practice.

Official syllabus

Pearson Edexcel International GCSE in Chemistry (4CH1) specification, Issue 3, September 2024, published by Pearson Education Limited. Topic 1 Principles of chemistry, sub-topics (c) Atomic structure (points 1.14–1.17) and (d) The Periodic Table (points 1.18–1.24).