IB MYP Sciences – Atomic structure and bonding Revision Notes

These revision notes cover atomic structure and bonding for IB Middle Years Programme Sciences, aligned to the International Baccalaureate Organization, Middle Years Programme Subject Brief – Sciences, from 2014. The brief lists “atomic structure and bonding” as a topic explored in the MYP sciences on-screen examinations. There is no SL/HL split in MYP, and the notes suit MYP years 4 and 5.

MYP has no prescribed content list – schools design their own units – so check the depth with your teacher, who will share the task-specific clarifications for assessed work. For full explanations and worked examples, read the atomic structure and bonding study guide first. Then try the practice questions. Course context: course hub and printable checklist.

Key definitions

• Atom: the smallest particle of an element; a nucleus of protons and neutrons with electrons in shells around it.
• Atomic number (Z): number of protons in the nucleus. It identifies the element.
• Mass number (A): number of protons + neutrons.
• Isotopes: atoms of the same element with the same number of protons but different numbers of neutrons.
• Relative atomic mass (Aᵣ): weighted mean mass of an element’s atoms, relative to 1/12 of the mass of a carbon-12 atom.
• Ion: an atom or group of atoms with an overall charge, formed by losing or gaining electrons.
• Ionic bond: the strong electrostatic attraction between oppositely charged ions.
• Covalent bond: a shared pair of electrons between two non-metal atoms.
• Metallic bond: the attraction between a lattice of positive metal ions and delocalised electrons.
• Delocalised electrons: electrons not tied to one atom, free to move through a structure.
• Alloy: a mixture of a metal with at least one other element.

The three particles

Particle	Relative mass	Relative charge	Where
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus
Electron	about 1/1840	−1	Shells

Formulas and rules

What	Rule
Neutrons	A − Z
Electrons in a neutral atom	Z
Electrons in an ion	Z − charge (so 2+ means Z − 2; 2− means Z + 2)
Relative atomic mass	Σ(mass number × % abundance) ÷ 100
Shell capacities (first 20 elements)	2, 8, 8, then the 4th shell
Period	number of occupied shells
Group (1, 2, 13-17)	outer electrons (groups 13-17: last digit)
Ionic formula	total + charge = total − charge

Method in steps

Particles in an atom or ion

1. Protons = atomic number (bottom number).
2. Neutrons = mass number − atomic number.
3. Electrons = protons, then subtract for a positive charge or add for a negative charge.
4. Arrangement: fill 2, 8, 8 in order.

Relative atomic mass

1. Multiply each isotope’s mass number by its percentage.
2. Add the products.
3. Divide by 100 (or by the total abundance if it is not 100).
4. Check: the answer lies between the isotope masses, nearer the most common one.

Formula of an ionic compound

1. Write both ion charges (group 1: +1, group 2: +2, Al: +3; group 17: −1, group 16: −2, N: −3).
2. Find the lowest common multiple of the charge sizes.
3. Use enough of each ion to reach it. Write subscripts, not charges, in the formula.

Explaining a property from bonding

1. Name the structure (giant ionic, simple molecular, giant covalent, giant metallic).
2. Name the particles and the force that must be overcome, or the particles that carry charge.
3. Say whether that force is strong or weak, or whether the charged particles can move.
4. Link back to the property in the question.

Small worked reminders

• ¹⁹₉F⁻: 9 protons, 10 neutrons, 10 electrons (2,8).
• ²⁴₁₂Mg²⁺: 12 protons, 12 neutrons, 10 electrons (2,8).
• Lithium: 7.5% lithium-6 and 92.5% lithium-7 gives Aᵣ = (6 × 7.5 + 7 × 92.5) ÷ 100 = 6.925, so 6.9 to 1 d.p.
• Magnesium chloride: Mg²⁺ and Cl⁻, so MgCl₂. Aluminium oxide: Al³⁺ and O²⁻, so Al₂O₃.
• Ammonia NH₃: 3 bonding pairs, 1 lone pair on nitrogen.
• Sodium chloride dot-and-cross: sodium’s one outer electron (a dot) moves to chlorine. Draw [Na]⁺ as 2,8 and [Cl]⁻ with 8 outer electrons, seven crosses and one dot, each ion in square brackets with its charge.
• Water dot-and-cross: oxygen shares one pair with each hydrogen, leaving two lone pairs; oxygen ends with 8 outer electrons and each hydrogen with 2.

Bonding and properties at a glance

	Giant ionic	Simple molecular	Giant covalent	Giant metallic
Example	NaCl, MgO	H₂O, CO₂, CH₄	Diamond, graphite, SiO₂	Cu, Fe, Al
Bond type	Ionic	Covalent in molecule; weak forces between	Covalent throughout	Metallic
Melting point	High	Low	Very high	Usually high
Conducts as solid	No	No	No (graphite: yes)	Yes
Conducts molten	Yes	No	No	Yes
Other	Brittle	Often gases or liquids at room temperature	Diamond hard; graphite soft	Malleable, ductile

Must-know distinctions

• Mass number vs relative atomic mass: mass number is a whole number for one isotope; Aᵣ is an average for the element and is often not a whole number.
• Atomic number vs number of electrons: equal in an atom, different in an ion.
• Covalent bonds vs intermolecular forces: covalent bonds hold atoms together inside a molecule and are strong; intermolecular forces act between molecules and are weak. Boiling water separates molecules; it does not break O-H bonds.
• Ionic vs metallic conduction: ionic compounds conduct by moving ions (only when molten or dissolved); metals conduct by moving delocalised electrons (solid or liquid).
• Diamond vs graphite: 4 bonds per carbon and rigid (hard, non-conductor) vs 3 bonds per carbon in sliding layers with delocalised electrons (soft, conductor).
• Pure metal vs alloy: regular layers slide easily vs distorted layers resist sliding.
• Losing vs gaining electrons: metals lose electrons (positive ions); non-metals gain electrons (negative ions).

Models and concepts

The brief names models as a related concept for MYP sciences, and change, relationships and systems as key concepts. This unit gives you ready examples:

• Models: the plum pudding model was replaced by the nuclear model after the gold-foil experiment showed most alpha particles pass straight through and a few bounce back.
• Relationships: electron arrangement explains group, period and the charge of an ion.
• Systems: the same element (carbon) behaves very differently as diamond and graphite.

Quick self-test

1. State the relative charge and relative mass of a neutron.
2. How many protons, neutrons and electrons are in ³¹₁₅P?
3. Write the electron arrangement of phosphorus.
4. How many electrons does a K⁺ ion formed from ³⁹₁₉K have, and what is its arrangement?
5. How many electrons are in ¹⁴₇N³⁻?
6. Boron is 20% boron-10 and 80% boron-11. Calculate Aᵣ.
7. Write the formulas of sodium oxide and calcium fluoride.
8. Why do isotopes of an element react in the same way?
9. Silicon has Z = 14. State its group and period.
10. Explain why solid sodium chloride does not conduct but molten sodium chloride does.
11. Explain why methane has a low boiling point.
12. Why is an alloy usually harder than the pure metal?

Answers

1. Charge 0; mass 1.
2. 15 protons, 16 neutrons, 15 electrons.
3. 2,8,5.
4. 18 electrons; 2,8,8.
5. 10 electrons (7 + 3).
6. (10 × 20 + 11 × 80) ÷ 100 = 1080 ÷ 100 = 10.8.
7. Na₂O; CaF₂.
8. They have the same number of electrons and the same electron arrangement, and chemical reactions involve electrons.
9. Arrangement 2,8,4: group 14, period 3.
10. In the solid the ions are fixed in the lattice. When molten the ions are free to move and carry charge.
11. Methane is a simple molecule. Only weak intermolecular forces are overcome on boiling, which needs little energy. The covalent bonds are not broken.
12. Atoms of different sizes distort the regular layers, so the layers cannot slide over each other as easily.

Where marks are usually lost

• Saying “the ionic bond is between the atoms” – it is between oppositely charged ions.
• Writing ionic formulas with charges left in (Mg²⁺Cl⁻₂) or unbalanced (AlO).
• Changing the proton number when an ion forms.
• Mixing up the direction: a 2− ion has two more electrons than protons, not two fewer.
• Saying covalent bonds break when a simple molecular substance melts or boils.
• Explaining graphite’s conductivity with “free ions” rather than delocalised electrons.
• Giving Aᵣ without working, or rounding to a whole number when the data give a decimal.
• Drawing dot-and-cross diagrams for ionic compounds as if the electrons were shared.
• Describing metals as “atoms in a sea of electrons” – the lattice is of positive ions.
• In a data task, classifying a substance without quoting the evidence (melting point, conductivity) that supports it.

For practice tied to each criterion, go to the practice questions. For the particle model, changes of state and related properties, see the separate states and properties of matter revision notes. For how the criteria work in general, see the assessment revision notes and the criteria in practice revision notes.

Official syllabus

International Baccalaureate Organization, Middle Years Programme Subject Brief – Sciences, from 2014. The brief lists “atomic structure and bonding” among the topics of the MYP sciences on-screen examinations and names models as a related concept; detailed unit content is set by schools.