Define proton number/atomic number as the number of protons in the nucleus of an atom

Atoms, Elements and Compounds – Atomic Structure and the Periodic Table

Learning Objectives (Cambridge IGCSE 0620)

• Identify protons, neutrons and electrons and state their relative charges and masses.
• Define **atomic number (Z)** and **mass number (A)** and write nuclear notation ⁽ᴬ⁾₍ᶻ₎X.
• Explain why the number of electrons equals Z only for a neutral atom.
• Write electron‑configuration symbols for the first 20 elements and relate them to groups and periods.
• Use the periodic trends:
  • Group → typical ionic charge (main‑group elements)
  • Period → number of electron shells
  • Achieving a noble‑gas configuration
• Describe isotopes, calculate relative atomic mass and state that isotopes have identical chemical properties.
• Distinguish cations, anions and covalent molecules; sketch dot‑and‑cross diagrams; list characteristic properties of ionic, simple covalent, giant covalent and metallic solids.

1. Atomic Structure

Sub‑atomic particles

• Proton – nucleus, charge +1 e, mass ≈ 1 u.
• Neutron – nucleus, charge 0, mass ≈ 1 u.
• Electron – surrounds nucleus in shells, charge –1 e, mass ≈ 0 u.

Atomic number (Z) and mass number (A)

Atomic number (Z) – the total count of protons in the nucleus. It uniquely identifies an element.

Mass number (A) – the total number of protons + neutrons (A = Z + N).

In nuclear notation the element is written as ⁽ᴬ⁾₍ᶻ₎X, where X is the element symbol.

For a **neutral** atom the number of electrons equals the atomic number (#e⁻ = Z). If the atom carries a charge, the electron count differs by the magnitude of that charge.

Why Z is important

1. Identifies the element – all atoms with the same Z are the same element.
2. Determines the element’s position in the Periodic Table (increasing Z from left to right, top to bottom).
3. Equals the number of electrons in a neutral atom, giving electrical neutrality.
4. Appears in the nuclear notation ⁽ᴬ⁾₍ᶻ₎X used throughout the syllabus.

Notation example – neutral carbon

⁽¹²⁾₍₆₎C Z = 6 (six protons) A = 12 (six protons + six neutrons)

Sample elements and their proton numbers

Element	Symbol	Z (protons)	A (most common isotope)
Hydrogen	H	1	1
Helium	He	2	4
Lithium	Li	3	7
Carbon	C	6	12
Oxygen	O	8	16
Iron	Fe	26	56

Group → typical ionic charge (main‑group elements)

Group	Typical charge of the element’s ion
I	+1 (e.g., Na⁺)
II	+2 (e.g., Mg²⁺)
III	+3 (e.g., Al³⁺)
IV	+4 (e.g., C⁴⁺ in some compounds)
V	+5 (e.g., N⁵⁺ in NO₃⁻)
VI	–2 (e.g., O²⁻)
VII	–1 (e.g., Cl⁻)

Period → number of electron shells

In a neutral atom the period number equals the number of occupied electron shells. Example: Sodium (Na) is in period 3 → three shells (K, L, M).

2. Electron Configuration (Elements 1‑20)

Electrons fill shells in the order 1s, 2s, 2p, 3s, 3p … (Aufbau principle). The table below gives the ground‑state configurations for the first twenty elements.

Z	Element	Electron configuration	Group (I‑VII)	Period
1	H	1s¹	1	1
2	He	1s²	18	1
3	Li	1s² 2s¹	1	2
4	Be	1s² 2s²	2	2
5	B	1s² 2s² 2p¹	13	2
6	C	1s² 2s² 2p²	14	2
7	N	1s² 2s² 2p³	15	2
8	O	1s² 2s² 2p⁴	16	2
9	F	1s² 2s² 2p⁵	17	2
10	Ne	1s² 2s² 2p⁶	18	2
11	Na	1s² 2s² 2p⁶ 3s¹	1	3
12	Mg	1s² 2s² 2p⁶ 3s²	2	3
13	Al	1s² 2s² 2p⁶ 3s² 3p¹	13	3
14	Si	1s² 2s² 2p⁶ 3s² 3p²	14	3
15	P	1s² 2s² 2p⁶ 3s² 3p³	15	3
16	S	1s² 2s² 2p⁶ 3s² 3p⁴	16	3
17	Cl	1s² 2s² 2p⁶ 3s² 3p⁵	17	3
18	Ar	1s² 2s² 2p⁶ 3s² 3p⁶	18	3
19	K	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹	1	4
20	Ca	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²	2	4

Rule for main‑group elements: The group number (I–VII) equals the number of valence electrons, which determines the typical ionic charge shown above.

3. Periodic Trends Linked to Z

• Group → ionic charge (see the table in Section 1).
• Period → number of electron shells (Period n → n shells).
• Noble‑gas configuration – atoms tend to gain, lose or share electrons until they achieve the electron configuration of the nearest noble gas (full outer shell).

4. Isotopes

An isotope of an element has the same atomic number Z (same number of protons) but a different number of neutrons N, giving a different mass number A. Because chemical behaviour depends on electron arrangement, isotopes have **identical chemical properties**.

Example – Chlorine

• ⁽³⁵⁾₁₇Cl Z = 17, N = 18 (A = 35)
• ⁽³⁷⁾₁₇Cl Z = 17, N = 20 (A = 37)

Natural chlorine is a mixture of these two isotopes (≈ 75 % ³⁵Cl, 25 % ³⁷Cl). The relative atomic mass is the weighted average:

Ar(Cl) = (0.75 × 35) + (0.25 × 37) = 35.5 u

Worked example – carbon

1. Isotopes: ¹²C (98.9 % abundance, A = 12) and ¹³C (1.1 % abundance, A = 13).
2. Calculate the relative atomic mass:

   Ar(C) = (0.989 × 12) + (0.011 × 13)
         = 11.868 + 0.143
         = 12.011 u
   

5. Ions and Ionic Bonding (2.4)

Formation of ions

• Cation – a positively charged ion formed by loss of electrons (e.g., Na → Na⁺).
• Anion – a negatively charged ion formed by gain of electrons (e.g., Cl → Cl⁻).

Quick checklist for ionic‑bond formation

1. Metal from Group I or II (or a transition metal that can lose electrons).
2. Non‑metal from Group VI or VII.
3. Transfer of electrons equal to the magnitude of the charge required for each ion.

Example – formation of NaCl

1. Na (Z = 11) loses one electron → Na⁺ (10 e⁻).
2. Cl (Z = 17) gains one electron → Cl⁻ (18 e⁻).
3. Opposite charges attract, producing an ionic lattice.

Dot‑and‑cross diagram (textual)

Na·   :   :Cl⁻
   +   -

Properties of ionic compounds (exact syllabus wording)

• High melting and boiling points (strong electrostatic forces).
• Hard and brittle solids.
• Conduct electricity when molten or dissolved in water (ions are free to move).
• Do **not** conduct electricity in the solid state.

6. Simple Covalent Molecules (2.5)

Formation

Covalent bonds are formed by the sharing of electrons between non‑metal atoms so that each atom attains a noble‑gas configuration.

Examples and dot‑and‑cross diagrams

Water – H₂O

   H : O : H
      ..

Methane – CH₄

       H
       .
   H – C – H
       .
       H

Properties of simple covalent compounds (syllabus phrasing)

• Low to moderate melting and boiling points.
• Usually gases or liquids at room temperature.
• Do **not** conduct electricity (no free ions).

7. Giant Covalent Structures (2.6)

Diamond

• Each carbon atom is covalently bonded to four others in a three‑dimensional tetrahedral network.
• Properties: extremely hard, very high melting point, electrical insulator.
• Uses: cutting tools, abrasives, jewellery.

Graphite

• Carbon atoms form layers of hexagonal sheets; each carbon is bonded to three neighbours, with one electron delocalised over each layer.
• Properties: soft, slippery, good conductor of electricity along the layers.
• Uses: pencil “lead”, lubricants, electrodes.

Comparative table – diamond vs. graphite

Feature	Diamond	Graphite
Structure	3‑D tetrahedral network	Layered sheets of hexagons
Bonding per atom	4 strong covalent bonds	3 strong covalent bonds + 1 delocalised electron
Hardness	Very hard	Soft, lubricating
Electrical conductivity	Insulator	Conductor (within layers)
Melting point	Very high	Very high (but layers can slide)

8. Metallic Bonding (2.7)

Metals consist of positively charged metal ions surrounded by a “sea of delocalised electrons”. This model explains the characteristic properties of metals.

Properties of metallic solids (syllabus wording)

• High electrical and thermal conductivity (mobile electrons).
• Malleable and ductile (ions can slide while the electron sea holds the structure together).
• High melting and boiling points (strong metallic bonds).
• Lustrous appearance (reflection of light by the electron sea).

9. Suggested Classroom Activities

1. Build the Periodic Table – using element cards, arrange them in order of increasing Z and colour‑code by period and group.
2. Calculate unknown Z – give students the mass number A and number of neutrons N; ask them to find Z = A – N.
3. Bead model of an atom – red beads = protons, blue = neutrons, green = electrons. Counting the red beads reinforces the concept of Z.
4. Ion‑formation practice – start from neutral nuclear notation (e.g., ⁽²³⁾₍₁₁₎Na) and write the ion notation after loss/gain of electrons (⁽²³⁾₍₁₁₎Na⁺).
5. Isotope mass‑average – give percentages of two isotopes and have students compute the relative atomic mass (use the carbon example).
6. Dot‑and‑cross drawing relay – teams draw the diagrams for H₂O, CH₄, NH₃ and CO₂ on the board, checking for correct electron sharing.
7. Compare giant covalent structures – using model kits, construct diamond and graphite lattices and fill in the comparative table.
8. Metallic‑bond demonstration – melt a small piece of metal (under supervision) and discuss why it conducts electricity in the liquid state.