IB MYP Sciences – Atomic structure and bonding Study Guide

This study guide teaches atomic structure and bonding for IB Middle Years Programme Sciences. It is written against the International Baccalaureate Organization, Middle Years Programme Subject Brief – Sciences, from 2014, which lists “atomic structure and bonding” among the topics explored in the MYP sciences on-screen examinations. MYP has no SL/HL split, and this page suits MYP years 4 and 5, whether you sit the eAssessment or are assessed by your school.

One thing to be clear about first: MYP has no prescribed content list. Schools design their own science units, so the exact depth your teacher expects may differ from this page. What follows is the standard chemistry that sits behind a topic the IB’s brief names, taught so you can use it in criterion A to D style tasks. Your teacher will share the task-specific clarifications for any assessed work.

Use this page with the revision notes and the practice questions. The course hub and the printable checklist show where the unit fits in the course.

What this unit covers

Area	What you must be able to do	Criterion it mostly trains
The nuclear model	Describe the atom and explain how evidence changed the model	A, D
Subatomic particles	Give relative mass, charge and location of protons, neutrons, electrons	A
Atomic and mass number	Work out numbers of each particle from nuclide notation	A
Isotopes	Define isotopes; calculate relative atomic mass from abundances	A, C
Electron arrangement	Write arrangements for the first 20 elements; link to group and period	A
Ions	Explain how and why atoms form ions; count particles in ions	A
Ionic, covalent, metallic bonding	Describe each bond; draw dot-and-cross diagrams; write formulas	A
Structure and properties	Explain melting point, conductivity and hardness from bonding	A, B, C

The nuclear model of the atom

An atom has a tiny, dense, positively charged nucleus containing protons and neutrons. Electrons move around the nucleus in shells (energy levels). Almost all of an atom’s mass is in the nucleus, but almost all of its volume is empty space.

This picture is a model, one of the related concepts the brief gives for MYP sciences. Models change when evidence demands it:

• J. J. Thomson discovered the electron in 1897 and proposed the “plum pudding” model: negative electrons spread through a ball of positive charge.
• In the gold-foil experiment carried out for Ernest Rutherford by Geiger and Marsden, alpha particles were fired at thin gold foil. Most passed straight through, so the atom is mostly empty space. A very few were deflected through large angles, some bouncing back, so the positive charge and mass must be packed into a tiny nucleus. Rutherford published the nuclear model in 1911.
• Niels Bohr (1913) placed electrons in fixed energy levels.
• James Chadwick identified the neutron in 1932.

Worked example – scale. An atom is about 1 × 10⁻¹⁰ m across and its nucleus about 1 × 10⁻¹⁴ m across. How many times wider is the atom?

ratio = (1 × 10⁻¹⁰) ÷ (1 × 10⁻¹⁴)
      = 10⁴ = 10 000

The atom is about 10 000 times wider than its nucleus.

Subatomic particles

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

An atom has equal numbers of protons and electrons, so it has no overall charge.

Atomic number and mass number

• Atomic number (Z): the number of protons. It defines the element.
• Mass number (A): the number of protons plus neutrons.
• Number of neutrons = A − Z.

In nuclide notation the mass number is written top left and the atomic number bottom left, for example ²³₁₁Na.

Worked example. State the numbers of protons, neutrons and electrons in ²³₁₁Na.

1. Protons = Z = 11.
2. Neutrons = 23 − 11 = 12.
3. Neutral atom, so electrons = protons = 11.

Isotopes and relative atomic mass

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Chlorine-35 and chlorine-37 both have 17 protons; they have 18 and 20 neutrons.

Isotopes of an element have the same electron arrangement, so they have the same chemical properties. They differ in mass, so physical properties that depend on mass (such as density) differ slightly.

Relative atomic mass (Aᵣ) is the weighted mean mass of an element’s atoms compared with 1/12 of the mass of a carbon-12 atom. Weighted means the more common isotope counts for more.

Aᵣ = Σ(mass number × percentage abundance) ÷ 100

Worked example. A sample of chlorine is 75% chlorine-35 and 25% chlorine-37. Calculate Aᵣ.

Aᵣ = (35 × 75 + 37 × 25) ÷ 100
   = (2625 + 925) ÷ 100
   = 3550 ÷ 100
   = 35.5

Check that your answer lies between the lightest and heaviest isotope, and closer to the more abundant one. 35.5 is closer to 35, which matches.

Electron arrangement

For the first 20 elements, electrons fill shells in order: up to 2 in the first shell, then up to 8 in the second, then up to 8 in the third. The 19th and 20th electrons go into the fourth shell.

Element	Z	Arrangement
Lithium	3	2,1
Carbon	6	2,4
Oxygen	8	2,6
Sodium	11	2,8,1
Magnesium	12	2,8,2
Sulfur	16	2,8,6
Chlorine	17	2,8,7
Calcium	20	2,8,8,2

Two links to the periodic table follow directly:

• The number of occupied shells equals the period number. Sodium (2,8,1) has three shells: period 3.
• The number of outer-shell electrons equals the group number for groups 1 and 2, and the last digit for groups 13 to 17 (for example chlorine, group 17, has 7). Noble gases have full outer shells.

Elements in the same group react similarly because they have the same number of outer electrons – the key concept of relationships in action.

Ions

An ion is an atom or group of atoms with an overall charge because electrons have been lost or gained. Atoms react so that they reach a full outer shell, like a noble gas.

• Metals lose outer electrons and form positive ions (cations). Group 1 forms 1+, group 2 forms 2+, aluminium forms 3+.
• Non-metals gain electrons and form negative ions (anions). Group 17 forms 1−, group 16 forms 2−, nitrogen forms 3−.

Worked example. Give the numbers of protons, neutrons and electrons in ¹⁶₈O²⁻ and its electron arrangement.

1. Protons = 8 (the atomic number never changes when an ion forms).
2. Neutrons = 16 − 8 = 8.
3. Charge 2− means two electrons gained: 8 + 2 = 10 electrons.
4. Arrangement 2,8 – the same as neon.

The same method gives ⁴⁰₂₀Ca²⁺: 20 protons, 20 neutrons, 18 electrons, arrangement 2,8,8.

Ionic bonding

When a metal reacts with a non-metal, electrons transfer from the metal atoms to the non-metal atoms. The oppositely charged ions attract each other strongly in all directions. This electrostatic attraction is the ionic bond.

The ions pack into a giant ionic lattice: a regular three-dimensional arrangement of alternating positive and negative ions.

In a dot-and-cross diagram, show the electrons of one atom as dots and the other as crosses. Draw each ion in square brackets with its charge outside the top right. For sodium chloride: [Na]⁺ with arrangement 2,8 and [Cl]⁻ with a full outer shell of 8, one of which is a dot from sodium.

Worked example – formula. Write the formula of aluminium oxide.

1. Ions: Al³⁺ and O²⁻.
2. Total positive charge must equal total negative charge. The lowest common multiple of 3 and 2 is 6.
3. Two Al³⁺ give +6; three O²⁻ give −6.
4. Formula Al₂O₃.

Properties of ionic compounds

• High melting and boiling points: many strong attractions must be overcome throughout the lattice.
• Conduct electricity when molten or dissolved, because the ions are then free to move and carry charge. They do not conduct as solids: the ions are fixed in place.
• Often soluble in water; hard but brittle – a shift in the layers puts like charges next to each other and they repel.

Covalent bonding

When non-metal atoms bond, they share pairs of electrons. A covalent bond is a shared pair of electrons, held by the attraction of both nuclei to that pair.

Molecule	Shared pairs	Lone pairs on central atom
H₂	1 (single bond)	–
H₂O	2	2 on O
NH₃	3	1 on N
CH₄	4	0 on C
O₂	2 (double bond)	–
N₂	3 (triple bond)	–
CO₂	2 double bonds (4 pairs)	0 on C

In a dot-and-cross diagram, draw overlapping shells with the shared pairs in the overlap. Every atom except hydrogen should end with 8 outer electrons; hydrogen ends with 2.

Simple molecular substances

Water, carbon dioxide and methane are made of small molecules. The covalent bonds inside each molecule are strong, but the forces between molecules are weak. Melting or boiling only overcomes these weak intermolecular forces, not the covalent bonds, so these substances have low melting and boiling points. They do not conduct electricity: there are no ions or free electrons.

Giant covalent structures

Some substances are huge networks of atoms, all joined by covalent bonds.

• Diamond: each carbon forms 4 covalent bonds. Very hard, very high melting point, does not conduct. Used in cutting tools.
• Graphite: each carbon forms 3 bonds, in flat layers. The layers are held by weak forces and slide, so graphite is soft and slippery. The fourth electron on each carbon is delocalised, so graphite conducts electricity.
• Silicon dioxide (silica, in sand): very high melting point and hard.

Diamond and graphite are the same element. Their different properties come from different structures – a clear example of the key concept of systems.

Metallic bonding

A metal is a lattice of positive ions in a “sea” of delocalised electrons. The metallic bond is the strong attraction between the positive ions and the delocalised electrons.

• Conductors of electricity and heat: the delocalised electrons move through the structure.
• Malleable and ductile: layers of ions can slide over each other while the electron sea holds them together.
• Usually high melting points: strong attractions throughout the lattice.

An alloy is a mixture of a metal with other elements. The atoms of a different size distort the layers, so they slide less easily. Alloys such as steel are therefore harder than the pure metal.

Using this content in criterion B, C and D tasks

The brief describes three eAssessment tasks: Knowing and understanding (criterion A, 25 marks), Investigation skills (criteria B and C, 50 marks) and Applying science (criterion D, 25 marks). This topic feeds all three.

• Criterion B: design a fair test of how salt concentration affects the current a solution carries, keeping voltage and electrodes the same.
• Criterion C: classify unknown substances from melting-point and conductivity data, quoting the data.
• Criterion D: weigh a material’s uses against its costs, such as synthetic against mined diamond for cutting tools.

The criteria in practice guide and the investigation skills preparation page explain how each criterion is judged in general.

Common errors

• Saying ions form because atoms “want” a full shell. Say they gain or lose electrons to reach a stable full outer shell.
• Changing the number of protons when forming an ion. Only electrons change.
• Writing that simple molecular substances melt at low temperatures because “covalent bonds are weak”. The covalent bonds are strong; the intermolecular forces are weak.
• Saying ionic solids conduct because they contain ions. The ions must be free to move.
• Forgetting square brackets and charges on ions in dot-and-cross diagrams.
• Giving Aᵣ as a whole number or a value outside the range of the isotopes’ masses.

Where to go next

Recap with the revision notes, test yourself with the practice questions, and see the syllabus guide for the course framework.

Official syllabus

International Baccalaureate Organization, Middle Years Programme Subject Brief – Sciences, from 2014. It lists this topic for the on-screen examinations but prescribes no detailed content.