ILC Chemistry Ordinary Level Chapter 1 Complete

Chapter Overview

This complete chapter brings together the content from all four chapter parts and fills the syllabus gaps so the notes now cover the full Ordinary Level Chapter 1 specification.

1.1 Periodic Table 1.2 Atomic Structure 1.3 Radioactivity 1.4 Electronic Structure 1.5 Oxidation & Reduction Mandatory Experiment 1.1 Mandatory Experiment 1.2

Added to close syllabus gaps: symbols 1–36, main-group similarities, Li/Na/K reaction with water, conservation of mass, carbon-12 scale, use of the mass spectrometer, history of atomic theory, history of radioactivity, widespread occurrence of radioactivity, food irradiation, atomic radii trends, screening effect, electrochemical series, electrolysis, ionic movement, rusting, electroplating and purification of copper, and the full detail of the two mandatory experiments.

1.1 Periodic Table

Learning Objectives

Explain why the periodic table is useful.
Know the symbols of elements 1–36.
Describe key stages in the development of the periodic table.
State similarities in the main groups, especially alkali metals, alkaline earth metals, halogens and noble gases.
Explain what groups and periods are and how they help us recognise patterns.

Why the Periodic Table Matters

The periodic table is a chart of elements arranged to show repeating patterns in their physical and chemical properties. It is useful because it does two jobs at the same time:

Organisation: it turns a long list of elements into a meaningful pattern.
Prediction: it helps us predict how an element is likely to behave from its position.

Examiner Secret: if you are asked why the periodic table is useful, mention both organisation and prediction of properties.

Symbols of Elements 1–36

Atomic No.	Symbol	Name	Atomic No.	Symbol	Name
1	H	Hydrogen	19	K	Potassium
2	He	Helium	20	Ca	Calcium
3	Li	Lithium	21	Sc	Scandium
4	Be	Beryllium	22	Ti	Titanium
5	B	Boron	23	V	Vanadium
6	C	Carbon	24	Cr	Chromium
7	N	Nitrogen	25	Mn	Manganese
8	O	Oxygen	26	Fe	Iron
9	F	Fluorine	27	Co	Cobalt
10	Ne	Neon	28	Ni	Nickel
11	Na	Sodium	29	Cu	Copper
12	Mg	Magnesium	30	Zn	Zinc
13	Al	Aluminium	31	Ga	Gallium
14	Si	Silicon	32	Ge	Germanium
15	P	Phosphorus	33	As	Arsenic
16	S	Sulfur	34	Se	Selenium
17	Cl	Chlorine	35	Br	Bromine
18	Ar	Argon	36	Kr	Krypton

OL Focus: know the symbols well enough to recognise them in formulas and simple equations.

Development of the Periodic Table

Scientist	Main Idea	Why It Matters
Greeks	Early ideas about matter and the elements.	Important as the starting point in the history of scientific thought.
Robert Boyle	Helped define an element as a substance that cannot be broken down chemically into simpler substances.	Moved chemistry away from vague classical ideas.
Humphry Davy	Used electrolysis to isolate elements such as sodium and potassium.	Showed that new elements could be discovered by experiment.
Döbereiner	Grouped similar elements into triads.	Showed that some properties repeat.
Newlands	Arranged elements by increasing atomic mass and suggested the Law of Octaves.	Recognised a repeating pattern, but it did not work for all elements.
Mendeleev	Arranged elements mainly by increasing atomic mass and left gaps for missing elements.	Predicted properties of undiscovered elements correctly.
Moseley	Showed that elements should be arranged by atomic number.	Led to the modern periodic table.

Groups and Periods

Group

A group is a vertical column. Elements in the same group have similar chemical properties because they have the same number of outer electrons.

Period

A period is a horizontal row. As you move across a period, element properties change gradually in a repeating pattern.

Examiner Trap: a group is vertical and a period is horizontal.

Main Group Similarities

Group	What They Are Like	Ordinary Level Points to Know
Alkali metals (Group 1)	Soft, reactive metals	Have 1 outer electron, form +1 ions, react with water.
Alkaline earth metals (Group 2)	Reactive metals	Have 2 outer electrons, form +2 ions, less reactive than Group 1.
Halogens (Group 17)	Reactive non-metals	Have 7 outer electrons, form -1 ions, become less reactive down the group.
Noble gases (Group 18)	Very unreactive gases	Have full outer shells and are stable.

Simple Trends the Table Shows

Properties repeat in a regular pattern.
Metals are mostly on the left; non-metals are mostly on the right.
Elements in the same group behave in similar ways.
Reactivity of Group 1 increases down the group.
Reactivity of Group 17 decreases down the group.

Reaction with Water: Lithium, Sodium and Potassium

The syllabus specifically expects awareness of the reaction of lithium, sodium and potassium with water.

Metal	What You See	Main Idea
Lithium	Floats, fizzes gently, moves slowly.	Reactive, but the least reactive of the three.
Sodium	Floats, fizzes strongly, melts into a ball and moves quickly.	More reactive than lithium.
Potassium	Very vigorous, often ignites with a lilac flame.	Most reactive of the three.

2M + 2H₂O → 2MOH + H₂ (where M = Li, Na or K)

Key conclusion: reactivity increases down Group 1.

1.2 Atomic Structure

Atoms, Molecules and Ions

Matter is made of tiny particles. These particles may be atoms, molecules or ions.

Atom: the smallest particle of an element that still has the properties of that element.
Molecule: two or more atoms chemically joined together.
Ion: a charged particle formed when an atom gains or loses electrons.

Atoms are extremely small. Even though they are tiny, they explain why substances have mass, occupy space and react chemically.

Law of Conservation of Mass

In a chemical reaction, mass is conserved. This means the total mass of the reactants is equal to the total mass of the products, provided nothing escapes from the system.

Total mass before reaction = Total mass after reaction

Why this matters: atoms are rearranged in reactions, but they are not created or destroyed in ordinary chemical changes.

Subatomic Particles

Particle	Relative Charge	Relative Mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	-1	Very small / about 1/1836	Outside nucleus in energy levels

Atomic Number, Mass Number and Isotopes

Atomic number (Z) = number of protons

Mass number (A) = protons + neutrons

Neutrons = mass number − atomic number

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

Hydrogen isotopes

Hydrogen-1: 1 proton, 0 neutrons
Hydrogen-2: 1 proton, 1 neutron
Hydrogen-3: 1 proton, 2 neutrons

Carbon isotopes

Carbon-12: 6 protons, 6 neutrons
Carbon-14: 6 protons, 8 neutrons

Relative Atomic Mass (A_r)

The relative atomic mass of an element is the weighted average mass of its atoms compared with 1/12 of the mass of a carbon-12 atom.

Carbon-12 scale: the carbon-12 isotope is the reference used for relative atomic masses.

Examiner Trap: mass number is for one isotope only, but relative atomic mass is an average value found on the periodic table.

Historical Development of Atomic Theory

Scientist	Main Contribution
Dalton	Suggested matter is made of tiny atoms.
Thomson	Discovered the electron and showed atoms contain negative particles.
Rutherford	Showed that the atom has a tiny dense nucleus using the alpha-particle scattering experiment.
Bohr	Suggested electrons move in fixed energy levels around the nucleus.

Use of the Mass Spectrometer

A mass spectrometer is used to help determine relative atomic mass by separating particles according to their mass-to-charge ratio. At Ordinary Level you do not need the full theory, but you should know its use.

What it is used for

Finding the relative atomic mass of an element
Showing that isotopes exist
Comparing the abundance of isotopes

OL summary

Different isotopes give different peaks, and the combined data can be used to work out the average relative atomic mass.

1.3 Radioactivity

Definition

Radioactivity is the spontaneous emission of radiation from unstable nuclei. “Spontaneous” means it happens by itself and is not controlled by ordinary chemical reactions.

Alpha, Beta and Gamma Radiation

Type	Nature	Charge	Penetrating Power	Example Emitter
Alpha (α)	Helium nucleus	+2	Low; stopped by paper or skin	Americium-241 (²⁴¹Am)
Beta (β)	Fast-moving electron	-1	Medium; stopped by thin aluminium	Carbon-14 (¹⁴C)
Gamma (γ)	Electromagnetic radiation	0	High; reduced by thick lead or concrete	Cobalt-60 (⁶⁰Co)

Penetration order: alpha < beta < gamma.

Common Mistake: gamma is the most penetrating, but not the most ionising.

Half-Life

The half-life of a radioactive substance is the time taken for half the radioactive nuclei in a sample to decay. For the syllabus, treat this idea qualitatively rather than mathematically.

History of Radioactivity

Scientist	Contribution
Becquerel	Discovered radiation from uranium salts.
Marie Curie	Worked on radioactivity and helped discover new radioactive elements.
Pierre Curie	Worked with Marie Curie in the study of radioactivity.

Marie and Pierre Curie discovered polonium and radium.

Widespread Occurrence of Radioactivity

Radioactivity is not rare. It occurs naturally in rocks, soil, space radiation and some foods. Small amounts of natural background radiation are always present around us.

Uses of Radioisotopes

Radioisotope / Use	Application
Carbon-14	Age determination of archaeological remains and old organic materials.
Cobalt-60	Cancer treatment using gamma radiation.
Food irradiation	Killing microorganisms and increasing storage life of some foods.

Detection and Penetrating Power

The syllabus expects awareness that the properties of radiation can be demonstrated using a radiation source and detector. The principle of the Geiger–Müller tube is not required.

Examiner Trap: radioactivity is a nuclear change, not a chemical one.

1.4 Electronic Structure of Atoms

Energy Levels

Electrons are arranged in energy levels, also called shells, around the nucleus. For Ordinary Level you should know the shell arrangements of elements 1–20.

Electron Configurations 1–20

Element	Configuration	Outer Electrons	Typical Group Link
H	1	1	Group 1-type behavior
He	2	2	Noble gas stability
Li	2,1	1	Group 1
Be	2,2	2	Group 2
B	2,3	3	Group 13
C	2,4	4	Group 14
N	2,5	5	Group 15
O	2,6	6	Group 16
F	2,7	7	Group 17
Ne	2,8	8	Group 18
Na	2,8,1	1	Group 1
Mg	2,8,2	2	Group 2
Al	2,8,3	3	Group 13
Si	2,8,4	4	Group 14
P	2,8,5	5	Group 15
S	2,8,6	6	Group 16
Cl	2,8,7	7	Group 17
Ar	2,8,8	8	Group 18
K	2,8,8,1	1	Group 1
Ca	2,8,8,2	2	Group 2

Why Elements in the Same Group Behave Similarly

The chemical properties of an element depend mainly on its outer electrons. That is why sodium and potassium are similar, and chlorine and bromine are similar.

Atomic Radii and Trends

Trend	What Happens	Reason
Down a group	Atomic radius increases	Extra shells are added, so the outer electrons are farther from the nucleus.
Across a period	Atomic radius decreases	Nuclear charge increases while electrons are added to the same main shell, so the nucleus pulls the electrons in more strongly.

Screening Effect and Nuclear Charge

Screening effect means inner electrons reduce the pull of the nucleus on outer electrons. Nuclear charge is the positive pull from the protons in the nucleus.

Down Group 1, the outer electron is farther away and more screened, so it is easier to lose. Reactivity increases.
Down Group 17, the incoming electron is farther away and more screened, so halogens become less reactive.

Applications: Sodium Street Lights and Fireworks

When electrons absorb energy, they move to higher energy levels. When they fall back, they emit light of particular colours.

Sodium street lights give a strong yellow-orange light.
Fireworks use metal salts to produce different colours.

1.5 Oxidation and Reduction

Basic Meaning

At Ordinary Level, oxidation and reduction are explained in terms of electron transfer.

Oxidation = loss of electrons

Reduction = gain of electrons

Memory trick: OIL RIG = Oxidation Is Loss, Reduction Is Gain.

Simple Redox Examples

Na with Cl₂: sodium loses electrons, chlorine gains electrons.
Mg with O₂: magnesium is oxidised, oxygen is reduced.
Zn with Cu²⁺: zinc loses electrons to copper ions.

Oxidising and Reducing Agents

Type	What It Does	What Happens to It
Oxidising agent	Causes oxidation by accepting electrons	Itself gets reduced
Reducing agent	Causes reduction by donating electrons	Itself gets oxidised

Electrochemical Series

The electrochemical series is a list of metals arranged in order of how easily they are oxidised. More reactive metals lose electrons more easily.

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au

At this level, you mainly use it to explain simple displacement reactions.

Electrolysis of Copper Sulfate with Copper Electrodes

At the anode (+)

Copper atoms lose electrons and enter solution as Cu²⁺ ions.

Cu → Cu²⁺ + 2e⁻

At the cathode (−)

Copper ions gain electrons and are deposited as copper metal.

Cu²⁺ + 2e⁻ → Cu

The copper concentration stays nearly the same because copper dissolves at one electrode and is deposited at the other.

Electrolysis of Acidified Water with Inert Electrodes

Hydrogen is formed at the cathode and oxygen is formed at the anode.

Cathode: reduction gives hydrogen gas.
Anode: oxidation gives oxygen gas.

The gases are produced in a 2:1 volume ratio of hydrogen to oxygen.

Ionic Movement

In electrolysis, positive ions move to the cathode and negative ions move to the anode.

Applications and Social Aspects

Topic	Main Idea
Rusting of iron	Iron reacts with oxygen and water to form rust. This is corrosion.
Swimming-pool water treatment	Chlorine-based chemicals act as oxidising agents to kill microorganisms.
Use of scrap iron to extract copper	Iron, being more reactive, displaces copper from copper ion solution.
Electroplating	A thin layer of one metal is placed onto another by electrolysis.
Purification of copper	Electrolysis removes impurities and gives very pure copper.
Chrome and nickel plating	Used for protection and appearance.
Cutlery	Often plated for appearance and resistance to corrosion.

Mandatory Experiments

Experiment 1.1 – Flame Tests

Aim: identify metal ions from the colour they give in a flame.

Ion / Salt	Flame Colour
Lithium	Crimson red
Sodium	Yellow / yellow-orange
Potassium	Lilac
Barium	Apple green
Strontium	Red
Copper	Blue-green / green

Method

Clean a nichrome wire by dipping it in hydrochloric acid and heating it until no colour appears.
Dip the wire into the sample.
Place it in the non-luminous Bunsen flame.
Observe and record the flame colour.

Why colours appear

Electrons are excited by heat and then fall back to lower energy levels, releasing light of characteristic colours.

Exam tip: sodium contamination is common and can hide other colours.

Experiment 1.2 – Redox Reactions of Group VII Elements and Metal Displacement

Aim: show that halogens act as oxidising agents and that more reactive metals displace less reactive metals from solution.

Part A: Halogens as Oxidising Agents

Test halogen solutions with bromides, iodides, Fe²⁺ and sulfites.

Reaction Type	Example	What to Notice
Halogen displaces bromine	Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂	Brown/orange bromine appears.
Halogen displaces iodine	Cl₂ + 2I⁻ → 2Cl⁻ + I₂	Brown iodine appears.
Oxidation of Fe²⁺	Cl₂ oxidises Fe²⁺ to Fe³⁺	Shows halogen is an oxidising agent.
Oxidation of sulfite	Cl₂ oxidises SO₃²⁻ to sulfate	Again shows halogen is an oxidising agent.

2Br⁻ − 2e⁻ → Br₂

Only half-equation level detail such as the one above is needed at OL.

Part B: Displacement Reactions of Metals

Reaction	Meaning
Zn + Cu²⁺ → Zn²⁺ + Cu	Zinc is more reactive than copper, so zinc displaces copper.
Mg + Cu²⁺ → Mg²⁺ + Cu	Magnesium is more reactive than copper, so magnesium displaces copper.

Main conclusion: the more reactive species is more easily oxidised.

Interactive Simulators / Virtual Activities

Build an Atom: practise protons, neutrons, electrons, isotopes and ions.
Periodic Table Explorer: compare elements in Groups 1, 2, 17 and 18.
Flame Test Virtual Lab: identify ions by flame colour.
Electrolysis Simulation: track ion movement to the correct electrode.

Exam Practice Questions

Q1. State two reasons why the periodic table is useful. [2]

Q2. Explain one difference between Mendeleev’s table and the modern periodic table. [2]

Q3. Name the scientist associated with each of the following: atomic number, Law of Octaves, alpha-particle scattering, discovery of radiation from uranium salts. [4]

Q4. Define atomic number and mass number. [4]

Q5. Carbon-14 and carbon-12 are isotopes. Explain why. [3]

Q6. State the relative charge, relative mass and location of a proton, neutron and electron. [9]

Q7. What is meant by the law of conservation of mass? [2]

Q8. Name one alpha emitter, one beta emitter and one gamma emitter from the syllabus. [3]

Q9. Which type of radiation is most penetrating? Which is most ionising? [2]

Q10. State one use each of carbon-14 and cobalt-60. [2]

Q11. Write the electronic configurations of sodium, chlorine and calcium. [3]

Q12. Explain why potassium is more reactive than lithium. [3]

Q13. Define oxidation and reduction in terms of electron transfer. [4]

Q14. In the reaction Zn + Cu²⁺ → Zn²⁺ + Cu, identify the oxidised substance and the reduced substance. [4]

Q15. Give the flame colours for sodium, potassium and copper. [3]

MCQs with Answers

1. Which scientist arranged elements by atomic number?
A. Dalton B. Moseley C. Davy D. Thomson
Answer: B

2. Which particle has relative mass 1 and charge 0?
A. Electron B. Proton C. Neutron D. Ion
Answer: C

3. Which radiation is stopped by paper?
A. Alpha B. Beta C. Gamma D. X-rays
Answer: A

4. Which group contains the halogens?
A. 1 B. 2 C. 17 D. 18
Answer: C

5. Which process can be used to purify copper?
A. Neutralisation B. Electrolysis C. Filtration D. Distillation
Answer: B

Mark Scheme / Answer Guide

Q1. Any two: organises elements [1]; shows patterns [1]; predicts properties [1].

Q2. Mendeleev arranged mainly by atomic mass [1]; modern table by atomic number [1].

Q3. Atomic number – Moseley [1]; Law of Octaves – Newlands [1]; alpha-particle scattering – Rutherford [1]; uranium-salt radiation – Becquerel [1].

Q4. Atomic number = number of protons [2]. Mass number = protons + neutrons [2].

Q5. Same number of protons [1]; different numbers of neutrons [1]; therefore different mass numbers [1].

Q6. Proton: +1, mass 1, nucleus. Neutron: 0, mass 1, nucleus. Electron: -1, very small mass / 1/1836, outside nucleus in energy levels. [9 total]

Q7. Total mass of reactants equals total mass of products [2].

Q8. α: ²⁴¹Am [1]; β: ¹⁴C [1]; γ: ⁶⁰Co [1].

Q9. Most penetrating: gamma [1]. Most ionising: alpha [1].

Q10. Carbon-14: age determination [1]. Cobalt-60: cancer treatment [1].

Q11. Na 2,8,1 [1]; Cl 2,8,7 [1]; Ca 2,8,8,2 [1].

Q12. Potassium has an outer electron farther from the nucleus [1]; more screening [1]; easier to lose an electron, so more reactive [1].

Q13. Oxidation = loss of electrons [2]. Reduction = gain of electrons [2].

Q14. Zinc is oxidised [2]; copper ion is reduced [2].

Q15. Sodium yellow [1]; potassium lilac [1]; copper blue-green / green [1].

Last-Minute Revision Sheet

The modern periodic table is arranged by increasing atomic number.
Elements in the same group have similar properties because they have the same number of outer electrons.
Group 1 reactivity increases down the group. Group 17 reactivity decreases down the group.
Atomic number = protons. Mass number = protons + neutrons.
Relative atomic mass is an average based on the carbon-12 scale.
Radioactivity comes from unstable nuclei.
Alpha is least penetrating but most ionising. Gamma is most penetrating.
Electrons are arranged in energy levels.
Oxidation is loss of electrons. Reduction is gain of electrons.
More reactive metals displace less reactive metals from solution.
Flame tests identify certain metal ions by colour.

Self-Assessment Checklist

I can explain why the periodic table is useful.
I can recognise the symbols of elements 1–36.
I can compare Mendeleev’s table with the modern periodic table.
I can state similarities of alkali metals, alkaline earth metals, halogens and noble gases.
I can define atomic number, mass number and isotopes.
I can explain the carbon-12 scale for relative atomic mass.
I can describe the contributions of Dalton, Thomson, Rutherford and Bohr.
I can identify alpha, beta and gamma radiation and their properties.
I can explain half-life qualitatively.
I can write electron configurations up to element 20.
I can explain the trends in atomic radius down a group and across a period.
I can define oxidation and reduction in terms of electrons.
I can explain the electrochemical series at simple OL level.
I can describe the key steps and purpose of flame tests and simple displacement reactions.

PREMIUM REVISION NOTES

Chapter Overview

Contents

1.1 Periodic Table

Learning Objectives

Why the Periodic Table Matters

Symbols of Elements 1–36

Development of the Periodic Table

Groups and Periods

Group

Period

Main Group Similarities

Simple Trends the Table Shows

Reaction with Water: Lithium, Sodium and Potassium

1.2 Atomic Structure

Atoms, Molecules and Ions

Law of Conservation of Mass

Subatomic Particles

Atomic Number, Mass Number and Isotopes

Hydrogen isotopes

Carbon isotopes

Relative Atomic Mass (Ar)

Historical Development of Atomic Theory

Use of the Mass Spectrometer

What it is used for

OL summary

1.3 Radioactivity

Definition

Alpha, Beta and Gamma Radiation

Half-Life

History of Radioactivity

Widespread Occurrence of Radioactivity

Uses of Radioisotopes

Detection and Penetrating Power

1.4 Electronic Structure of Atoms

Energy Levels

Electron Configurations 1–20

Why Elements in the Same Group Behave Similarly

Atomic Radii and Trends

Screening Effect and Nuclear Charge

Applications: Sodium Street Lights and Fireworks

1.5 Oxidation and Reduction

Basic Meaning

Simple Redox Examples

Oxidising and Reducing Agents

Electrochemical Series

Electrolysis of Copper Sulfate with Copper Electrodes

At the anode (+)

At the cathode (−)

Electrolysis of Acidified Water with Inert Electrodes

Ionic Movement

Applications and Social Aspects

Mandatory Experiments

Experiment 1.1 – Flame Tests

Method

Why colours appear

Experiment 1.2 – Redox Reactions of Group VII Elements and Metal Displacement

Part A: Halogens as Oxidising Agents

Part B: Displacement Reactions of Metals

Interactive Simulators / Virtual Activities

Exam Practice Questions

MCQs with Answers

Mark Scheme / Answer Guide

Last-Minute Revision Sheet

Self-Assessment Checklist

Relative Atomic Mass (A_r)