Higher Level Revision Notes

Leaving Cert Higher Level Chemistry

Leaving Cert Higher Level Chemistry Chapter 1: Periodic Table and Atomic Structure

You learn how atoms are built, how the periodic table is organised, why isotopes matter, and how radioactivity is described in simple scientific language.

What this page doesTurns a difficult chapter into clear notes, diagrams, and one simple interactive tool.

Student promisePlain English first, exam language second.

Curriculum

Irish Leaving Certificate (ILC)

Level

Higher Level

Subject

Chemistry

Chapter

Chapter 1 — Periodic Table and Atomic Structure

Simple English Summary

You learn how atoms are built, how the periodic table is organised, why isotopes matter, and how radioactivity is described in simple scientific language.

Teacher voice: Read the ideas first, then use the detailed notes and diagrams to lock in the exam wording.

What To Focus On

Use atomic number and mass number correctly.
Explain isotopes, ions, and atomic structure.
Read basic radioactivity ideas with confidence.
Link periodic trends to the way the table is arranged.

DisclaimerThis publication is an independent educational resource developed by ExamsLogic and compiled for student revision. It is based on publicly available official curricula and is not endorsed by any examination board.

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Quick Simpler Start

In one sentence

You learn how atoms are built, how the periodic table is organised, why isotopes matter, and how radioactivity is described in simple scientific language.

Exam habit

When you revise this chapter, ask yourself: "Can I explain this to a friend in one easy paragraph?" If yes, you are in a good place.

Interactive Simulator

This small tool gives you a quick visual check before you move deeper into the chapter notes.

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Atom Builder

Type in the atomic number and number of neutrons. The tool builds the neutral atom for you in plain language.

Atomic numberNeutrons

Enter values to see the atom breakdown.

Chapter 1: Periodic Table and Atomic Structure

1. Learning Objectives

Describe the development of the periodic table from Döbereiner to Moseley.
Compare Mendeleev's periodic table with the modern periodic table.
Identify protons, neutrons and electrons by charge, mass and location.
Use atomic number and mass number to calculate protons, neutrons and electrons.
Explain isotopes and describe important radioisotope uses.
Compare alpha, beta and gamma radiation.
Write electron configurations for the first 36 elements.
HLExplain orbitals, sublevels, Aufbau, Hund's Rule, Pauli Exclusion Principle and Heisenberg's Uncertainty Principle.
Explain oxidation and reduction in terms of electron transfer.
HLUse oxidation numbers and balance redox equations.

2. The Development of the Periodic Table

The periodic table was not created in one step. It developed as scientists found better ways to arrange elements according to their properties and atomic structure.

The key improvement was Moseley's arrangement by atomic number, not atomic mass.

Mendeleev vs Modern Periodic Table

Mendeleev's Table	Modern Periodic Table
Elements mainly arranged by increasing atomic mass.	Elements arranged by increasing atomic number.
Gaps were left for undiscovered elements.	Elements fit according to proton number and electronic structure.
Some elements were placed out of strict mass order to match chemical properties.	Groups show similar properties because elements have similar outer electron arrangements.
Noble gases were not originally included.	Noble gases are placed in Group 18/Group 0.

Examiner Tip: If asked why Moseley's work was important, write: he arranged elements in order of increasing atomic number, which is the number of protons in the nucleus.

Examiner Trap: Do not say the modern table is arranged by mass number. It is arranged by atomic number.

3. Atomic Structure

Atoms are made of three main subatomic particles: protons, neutrons and electrons.

The nucleus is tiny but contains almost all the mass of the atom.

Particle	Relative Charge	Relative Mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	−1	Very small / 1/1836	Shells around nucleus

Atomic Number, Mass Number and Isotopes

Atomic number = number of protons

Mass number = protons + neutrons

Worked Example: Sodium has mass number 23 and atomic number 11.
Protons = 11
Electrons = 11 in a neutral atom
Neutrons = 23 − 11 = 12

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.

HL Only Brief History of Atomic Theory

Scientist	Main Contribution	Why it matters
Dalton	Suggested matter is made of tiny atoms.	First modern atomic theory model.
Thomson	Discovered the electron.	Showed atoms are divisible.
Rutherford	Gold-foil experiment led to nuclear model.	Atoms have a tiny dense nucleus.
Bohr	Electrons occupy fixed energy levels.	Helped explain line spectra.
Wave-mechanical model	Electrons described by probability orbitals.	Better HL model than simple circular shells.

HL focus: Each model improved because new experimental evidence showed limits in the previous one.

HL Only Relative Atomic Mass and Isotopic Abundance

For HL, relative atomic mass is based on the carbon-12 scale. In real samples, the value of A_r depends on the abundance of each isotope.

A_r = (isotope mass × % abundance + isotope mass × % abundance + ... ) / 100

Worked Example: Chlorine contains about 75% chlorine-35 and 25% chlorine-37.
A_r = (35 × 75 + 37 × 25) / 100 = 35.5

HL Only Mass Spectrometer

A mass spectrometer is used to identify isotopes and determine relative atomic mass. It separates particles according to their mass-to-charge ratio, often written as m/z.

Stage	What happens?	Purpose
Vaporisation	Sample is converted to a gas.	Particles must be in gas phase.
Ionisation	Gas particles are bombarded so positive ions form.	Only charged particles can be accelerated and detected.
Acceleration	Ions are accelerated by an electric field.	Gives ions kinetic energy.
Deflection	Ions are bent by a magnetic field.	Lighter ions or ions with bigger charge are deflected more.
Detection	Ions strike detector and produce a signal.	Creates a mass spectrum.

Mass spectrometry is one of the key HL links between isotopes and relative atomic mass.

4. Radioactivity

Radioactivity is the spontaneous emission of radiation from unstable nuclei. The three main types are alpha, beta and gamma radiation.

Alpha is least penetrating; gamma is most penetrating.

Radiation	Nature	Charge	Penetration	Ionising Power
Alpha (α)	Helium nucleus	+2	Low	High
Beta (β)	Fast electron	−1	Medium	Medium
Gamma (γ)	Electromagnetic radiation	0	High	Low

Half-Life

Half-life is the time taken for half of the radioactive nuclei in a sample to decay, or for the activity to fall to half its original value.

Each half-life halves the amount remaining: 100% → 50% → 25% → 12.5%.

Worked Example: Carbon-14 has a half-life of 5730 years. How long for 80 g to become 20 g?
80 g → 40 g → 20 g = 2 half-lives
Time = 2 × 5730 = 11,460 years.

Radioisotopes and Uses

Carbon-14

Used in archaeological dating of once-living materials.

Cobalt-60

Used in radiotherapy and sterilising medical equipment.

Tracers

Radioisotopes can track movement of substances in industry and medicine.

HL Only Nuclear Equations and Nuclear vs Chemical Change

In a chemical reaction, electrons are rearranged. In a nuclear reaction, the nucleus changes, so a new element may form.

Type	What changes?	Example
Chemical reaction	Electrons / bonding	Mg + 2HCl → MgCl₂ + H₂
Nuclear reaction	Nucleus	Alpha or beta decay

²³⁸₉₂U → ²³⁴₉₀Th + ⁴₂He

¹⁴₆C → ¹⁴₇N + ⁰₋₁e

Examiner Tip: In nuclear equations, both the mass numbers and the atomic numbers must balance.

HL Only Scientists Behind Radioactivity

Becquerel
Discovered natural radioactivity while working with uranium salts.

Marie and Pierre Curie
Studied radioactive substances in detail and isolated new radioactive elements.

5. Electronic Structure

Electrons occupy energy levels around the nucleus. For OL and HL, students should be able to write electron configurations for the first 36 elements.

Outer electrons explain many chemical properties.

Examiner Tip: For elements up to calcium, students often use 2,8,8,2. For elements after calcium up to krypton, configurations may involve the third shell holding more than 8 electrons.

HL Only Orbitals, Sublevels and Quantum Ideas

Bohr Model: Electrons move in fixed energy levels or shells.

Wave Mechanical Model: Electrons are found in orbitals, which are regions where there is a high probability of finding an electron.

Aufbau Principle
Electrons fill the lowest available energy level first.
Pauli Exclusion Principle
A maximum of two electrons can occupy one orbital, with opposite spins.
Hund's Rule
Electrons occupy equal-energy orbitals singly before pairing.
Heisenberg Uncertainty Principle
It is impossible to know both the exact position and exact momentum of an electron at the same time.

5A. Higher-Level Electronic Structure Extensions

HL Only Emission and Absorption Spectra

When electrons absorb energy, they move to a higher energy level. When they fall back down, they emit light of a specific wavelength. This gives a line spectrum.

Absorption spectrum
Electrons absorb certain wavelengths and jump to higher levels.

Emission spectrum
Electrons fall to lower levels and emit certain wavelengths.

Each element has its own pattern of spectral lines.

HL Only Balmer Series

The Balmer series is the set of visible lines in the hydrogen emission spectrum. These lines arise when excited electrons fall to the second energy level.

HL focus: You do not need deep quantum mathematics here, but you should know that spectral lines support the idea of fixed energy differences between levels.

HL Only First and Successive Ionisation Energies

The first ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

Na(g) → Na⁺(g) + e⁻

Successive ionisation energies remove further electrons one at a time.

Trend / idea	Explanation
Across a period, first ionisation energy usually increases.	Nuclear charge increases and atomic radius usually decreases.
Down a group, first ionisation energy usually decreases.	Outer electron is farther from nucleus and more screened.
Large jump in successive ionisation energies	Shows that the next electron comes from an inner shell.

Atomic radius

Smaller radius means the nucleus attracts the outer electron more strongly.

Screening effect

Inner electrons reduce the pull felt by outer electrons.

Nuclear charge

More protons means stronger attraction for electrons.

Worked Idea: A very large jump after removing two electrons suggests the atom had two outer electrons, so it was in Group 2.

Examiner Trap: Do not explain ionisation-energy trends using only one factor. HL answers are strongest when they mention radius + screening + nuclear charge together.

6. Oxidation and Reduction

Oxidation and reduction can be described in terms of electron transfer.

Oxidation = loss of electrons

Reduction = gain of electrons

Memory Trick: OIL RIG
Oxidation Is Loss; Reduction Is Gain.

Naming Transition Metal Compounds

Transition metals can have more than one possible charge, so Roman numerals are used in compound names.

Formula	Name	Meaning
FeCl₂	Iron(II) chloride	Iron has charge +2
FeCl₃	Iron(III) chloride	Iron has charge +3
CuO	Copper(II) oxide	Copper has charge +2

HL Only Oxidation Numbers and Balancing Redox Equations

Oxidation number increases: oxidation has occurred.

Oxidation number decreases: reduction has occurred.

In redox equations, the number of electrons lost must equal the number of electrons gained.

7. Worked Examples

Example 1: Isotope Calculation
An atom has atomic number 17 and mass number 35.
Protons = 17
Electrons = 17 if neutral
Neutrons = 35 − 17 = 18

Example 2: Half-Life
A radioactive sample has activity 640 counts per minute. After 3 half-lives, what is its activity?
640 → 320 → 160 → 80
Final activity = 80 counts per minute

Example 3: Redox
Mg → Mg²⁺ + 2e−
Magnesium loses electrons, so magnesium is oxidised.

8. Examiner Secrets, Mistakes and Traps

Examiner Secret: For atomic structure questions, always state both the number and the location of the particle if asked to describe it.

Common Mistake: Saying isotopes have different numbers of protons. They do not. Isotopes of the same element always have the same number of protons.

Common Mistake: Confusing mass number with atomic mass. Mass number is protons + neutrons for one isotope.

Examiner Trap: Gamma radiation is not a particle. It is electromagnetic radiation.

Examiner Trap: Oxidation is not always simply adding oxygen at HL. In redox chemistry, oxidation means loss of electrons or increase in oxidation number.

9. Practical Skills

Flame Tests
Used to identify metal ions from flame colours. For example, sodium gives a yellow flame.

Radiation Detection
A Geiger-Müller tube can detect ionising radiation and measure count rate.

Halogen Displacement / Redox Practical
Chlorine can displace bromide and iodide ions from solution because it is a stronger oxidising agent. Bromine can displace iodide ions. These colour changes help compare halogen reactivity and redox behaviour.

10. Exam Practice Questions

Q1. Explain why Mendeleev left gaps in his periodic table. [3 marks]

Q2. An isotope of potassium has mass number 39 and atomic number 19. Calculate the number of protons, neutrons and electrons in a neutral atom. [3 marks]

Q3. Compare alpha, beta and gamma radiation in terms of charge and penetrating power. [6 marks]

Q4. A radioactive sample decreases from 160 g to 20 g in 24 days. Calculate the half-life. [4 marks]

Q5. Explain oxidation and reduction using electron transfer, using magnesium reacting with copper(II) ions as an example. [5 marks]

MCQs with Explanations

1. Which particle determines the atomic number of an element?
A. Electron B. Proton C. Neutron D. Nucleus

Answer: B. Atomic number equals the number of protons. Electrons can change when ions form, but the element remains defined by proton number.

2. Which radiation has the greatest penetrating power?
A. Alpha B. Beta C. Gamma D. Neutron only

Answer: C. Gamma radiation is most penetrating and is reduced by thick lead.

3. Oxidation in electron transfer means:
A. Gain of electrons B. Loss of electrons C. Loss of neutrons D. Gain of protons

Answer: B. OIL RIG: Oxidation Is Loss; Reduction Is Gain.

11. Higher-Level Challenge

Challenge: Explain why the electronic configuration of an atom is more useful than atomic mass when predicting chemical properties.

Guidance: Chemical reactions involve outer electrons. Elements in the same group have similar outer electron arrangements, so they show similar chemical properties.

Redox Challenge: Identify oxidation and reduction in:

Zn + Cu²⁺ → Zn²⁺ + Cu

Zn loses electrons and is oxidised. Cu²⁺ gains electrons and is reduced.

12. Last-Minute Revision Sheet

Mendeleev arranged elements mainly by atomic mass and left gaps.
Moseley arranged elements by atomic number.
Atomic number = protons.
Mass number = protons + neutrons.
Isotopes have same protons but different neutrons.
Alpha: +2, least penetrating, most ionising.
Beta: −1, medium penetration.
Gamma: no charge, most penetrating, least ionising.
Half-life means the sample or activity halves each interval.
Electron configuration explains group properties.
OIL RIG: Oxidation Is Loss; Reduction Is Gain.
HLAufbau fills lowest energy first; Hund fills singly first; Pauli allows two opposite-spin electrons per orbital.

13. Self-Assessment Checklist

I can explain how the periodic table developed.
I can compare Mendeleev's table with the modern periodic table.
I can identify protons, neutrons and electrons by charge, mass and location.
I can calculate protons, neutrons and electrons from atomic and mass numbers.
I can define isotopes.
I can compare alpha, beta and gamma radiation.
I can solve half-life problems.
I can write electron configurations for the first 36 elements.
I can explain oxidation and reduction using electron transfer.
I can answer HL questions on orbitals, sublevels and redox.

14. Mark Scheme

Q1. Mendeleev left gaps for undiscovered elements [1]; predicted their properties [1]; placed elements with similar chemical properties in groups even if mass order was not perfect [1].

Q2. Protons = 19 [1]; electrons = 19 [1]; neutrons = 39 − 19 = 20 [1].

Q3. Alpha has +2 charge and low penetration [2]; beta has −1 charge and medium penetration [2]; gamma has no charge and high penetration [2].

Q4. 160 → 80 → 40 → 20 = 3 half-lives [2]; 24 ÷ 3 = 8 days [2].

Q5. Oxidation is loss of electrons [1]; reduction is gain of electrons [1]; magnesium loses electrons to form Mg²⁺ [1]; copper(II) ions gain electrons to form copper [1]; correct example equations or explanation [1].