Hey Tara! Welcome to Topic 2 - the foundation of all chemistry. Everything around you - the air you breathe, the water you drink, the phone in your hand - is made of atoms. In this topic, you will learn what atoms are, how elements differ from compounds and mixtures, how atoms bond together, and the amazing structures they build. Think of this as learning the alphabet of chemistry - once you know these basics, you can read the entire language. Let us get started!
2.1 Elements, Compounds and Mixtures
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What is an Element?
An element is a pure substance that contains only one type of atom. It cannot be broken down into anything simpler by any chemical reaction.
There are about 118 known elements, and each one has its own unique symbol from the Periodic Table. Some symbols are easy to remember because they match the English name - like O for Oxygen, C for Carbon, and N for Nitrogen. Others come from Latin or Greek names - like Fe for Iron (from the Latin word Ferrum), Au for Gold (from Aurum), and Na for Sodium (from Natrium).
Some everyday examples of elements:
- Oxygen (O) - the gas you breathe in
- Iron (Fe) - used in construction, bridges, and railway tracks across India
- Gold (Au) - used in jewellery all around the world
- Carbon (C) - found in diamonds and pencil lead (graphite)
- Copper (Cu) - used in electrical wires and coins
What is a Compound?
A compound is a pure substance made of two or more elements chemically combined in a fixed ratio. The key word here is "chemically combined" - the atoms are bonded together and cannot be separated by physical methods.
Here is the really important part: the properties of a compound are completely different from the properties of the elements that make it up.
For example:
- Water (H₂O) is made from hydrogen (a flammable gas) and oxygen (a gas that supports burning). But water is a liquid that puts out fires! Completely different properties.
- Table salt (NaCl) is made from sodium (a highly reactive metal that explodes in water) and chlorine (a poisonous green gas). But table salt is a harmless white solid you put on your food every day.
- Carbon dioxide (CO₂) is made from carbon (a black solid) and oxygen (a colourless gas). But CO₂ is a colourless gas used in fizzy drinks.
Compounds can only be broken down into their elements by chemical reactions, not by physical methods like filtering or evaporating.
What is a Mixture?
A mixture contains two or more substances (which can be elements or compounds) that are NOT chemically combined. The substances are just mixed together - they are not bonded.
Key features of mixtures:
- Each substance in the mixture keeps its own properties - nothing changes chemically.
- The substances can be in any ratio - unlike compounds, there is no fixed ratio.
- Mixtures can be separated by physical methods such as filtration, distillation, evaporation, chromatography, and using a separating funnel.
- No chemical reaction is needed to separate a mixture.
Examples of mixtures:
- Air - a mixture of nitrogen, oxygen, carbon dioxide, water vapour, and other gases
- Sea water - a mixture of water and dissolved salts
- Alloys like steel (iron + carbon) and brass (copper + zinc)
- Ink - a mixture of different coloured dyes dissolved in a solvent
Elements contain only one type of atom. Compounds have different atoms chemically bonded together in fixed groups. Mixtures have different particles that are NOT bonded to each other.
Comparison Table
| Property | Element | Compound | Mixture |
|---|---|---|---|
| Made of | One type of atom only | Two or more different elements chemically combined | Two or more substances NOT chemically combined |
| Ratio of components | N/A (only one type) | Fixed ratio (e.g. H₂O is always 2:1) | Any ratio - can vary |
| Properties | Unique to that element | Different from the elements that make it up | Each substance keeps its own properties |
| Separation | Cannot be broken down by chemical means | Can only be separated by chemical reactions | Can be separated by physical methods |
| Examples | Iron (Fe), Oxygen (O₂), Gold (Au) | Water (H₂O), Salt (NaCl), CO₂ | Air, sea water, steel, ink |
🧠 Memory Trick
E-C-M - think of it like this: Elements are like a class of students all wearing the same uniform (all same atoms). Compounds are like students from different schools holding hands in a human chain (different atoms bonded together). Mixtures are like a busy train station where people from different places are just walking around together without holding hands (different particles, not bonded).
🧠 Memory Trick
To remember that compounds have different properties from their elements, think of this: Sodium (Na) is a dangerous metal that explodes in water. Chlorine (Cl) is a poisonous gas that was used as a weapon in World War I. But put them together and you get NaCl - table salt, something you happily sprinkle on your chips!
⚠ Exam Tip
A very common exam mistake is confusing an element with a compound. Remember: O₂ (oxygen gas) is still an element because it contains only one type of atom (oxygen). Even though there are two atoms bonded together, they are both the same type. A compound must have different types of atoms.
Similarly, H₂, N₂, Cl₂, Br₂, I₂, F₂ are all elements - not compounds. They are called diatomic elements.
⚠ Exam Tip
When the question asks you to classify a substance, look at the chemical formula. If it has only one type of element symbol (like Fe, O₂, S₈), it is an element. If it has two or more different element symbols (like H₂O, NaCl, CO₂), it is a compound. If the question describes substances "mixed together" or "dissolved in", it is likely a mixture.
Worked Example
Classify each of the following as an element, compound, or mixture: (a) Pure iron (b) Carbon dioxide (c) Salty water (d) Nitrogen gas (e) Bronze
Step 1
Ask yourself: does this substance contain one type of atom, different atoms bonded together, or different substances not bonded?
(a) Pure iron
Iron (Fe) contains only iron atoms. One type of atom = Element.
(b) Carbon dioxide
CO₂ contains carbon and oxygen atoms chemically bonded. Two different elements combined = Compound.
(c) Salty water
This is salt (NaCl) dissolved in water (H₂O). Two different substances mixed but not chemically combined. You can separate them by evaporation = Mixture.
(d) Nitrogen gas
N₂ contains only nitrogen atoms. Even though there are two atoms, they are the same type = Element.
(e) Bronze
Bronze is copper and tin mixed together (an alloy). The metals are not chemically bonded = Mixture.
Final Answers: (a) Element (b) Compound (c) Mixture (d) Element (e) Mixture
Worked Example
A student has a mixture of sand and salt. Describe how she could separate them to get pure dry salt.
Step 1
Add water to the mixture and stir. The salt dissolves in the water, but the sand does not dissolve.
Step 2
Filter the mixture using filter paper and a funnel. The sand stays on the filter paper as the residue. The salt solution passes through as the filtrate.
Step 3
Heat the filtrate (salt solution) in an evaporating dish. The water evaporates, leaving behind pure dry salt crystals.
Final Answer: Dissolve in water, filter to remove sand, then evaporate the water to get pure dry salt.
Apply It: Real-World Chemistry
Chemistry is not just a subject in your textbook - it is everywhere around you. Let us connect what you have learned about elements, compounds, and mixtures to the real world.
1
The air you are breathing right now is about 78% nitrogen, 21% oxygen, 0.04% carbon dioxide, and small amounts of water vapour and noble gases.
Why is air classified as a mixture and not a compound?
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The Chemistry
Air contains several different gases (N₂, O₂, CO₂, H₂O vapour, Ar), but they are NOT chemically bonded to each other. Each gas keeps its own properties - oxygen still supports burning, nitrogen is still unreactive, and CO₂ still turns limewater milky. The gases are simply mixed together in the atmosphere. Also, the composition of air can vary slightly from place to place (more humidity near the sea, more CO₂ in cities), which proves it is a mixture. A compound always has a fixed ratio.
The Big Idea
Mixtures can be separated by physical methods. That is exactly what happens in an air separation plant - air is cooled until it becomes liquid, then the different gases are separated by fractional distillation because they have different boiling points.
2
Water (H₂O) is made from hydrogen and oxygen. Hydrogen is a gas that burns with a squeaky pop. Oxygen is a gas that supports combustion. But water is a liquid that puts out fires.
Why are the properties of water so different from hydrogen and oxygen?
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The Chemistry
When hydrogen and oxygen react to form water, a chemical reaction takes place. The atoms rearrange and form new chemical bonds. This creates a completely new substance - water - with its own unique properties. In a compound, the elements are chemically combined, so the original properties of the individual elements disappear entirely. The hydrogen atoms and oxygen atoms are locked together in H₂O molecules through strong covalent bonds.
The Big Idea
This is the defining feature of compounds: their properties are completely different from the elements that make them up. Chemical bonding creates an entirely new substance. You cannot predict a compound's properties just by looking at the elements it contains.
3
Steel is made by mixing iron with a small amount of carbon (and sometimes other metals). Pure iron is relatively soft. But steel is much harder and stronger - it is used to build skyscrapers, bridges, and the frames of cars.
Is steel an element, a compound, or a mixture? How is it different from pure iron?
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The Chemistry
Steel is a mixture - specifically, it is an alloy (a mixture of a metal with other elements). The iron and carbon atoms are mixed together but not chemically bonded in a fixed ratio. Different types of steel have different amounts of carbon (mild steel has about 0.25% carbon, high-carbon steel has about 1.5%). The carbon atoms sit between the iron atoms and disrupt the regular arrangement, which makes it harder for the layers to slide over each other. Pure iron (Fe) is an element - it contains only iron atoms.
The Big Idea
Alloys are mixtures, not compounds, because their composition can vary and the atoms are not chemically bonded in a fixed ratio. This is exactly why steelmakers can create hundreds of different steel grades by adjusting the recipe.
4
When you dissolve salt (NaCl) in water, you get salt water. The salt seems to disappear - you cannot see it any more. But if you taste the water, it is salty. If you boil the water away, the salt reappears.
Is salt water a mixture or a compound? And is the salt itself (NaCl) a mixture or a compound?
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The Chemistry
Salt water is a mixture. The NaCl and H₂O are mixed together but not chemically bonded to each other. You can separate them by simply evaporating the water - a physical method. The salt and water each keep their own properties (the water is still water, the salt is still salt). However, the salt itself - NaCl - is a compound. Sodium (Na) and chlorine (Cl) are chemically bonded together in a fixed 1:1 ratio. You cannot separate them by physical methods - you would need electrolysis (a chemical process) to break NaCl apart.
The Big Idea
A mixture can contain compounds! Salt water is a mixture of a compound (NaCl) dissolved in another compound (H₂O). The test is simple: can you separate it by physical methods? If yes, it is a mixture.
5
Many ancient temples and statues in India are made of bronze. The famous Dancing Girl statue from the Indus Valley Civilisation (around 2500 BCE) was made of bronze - a mixture of copper and tin. After thousands of years, the statue has turned greenish on the surface.
Why is bronze classified as a mixture (alloy), and what caused the green colour?
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The Chemistry
Bronze is a mixture (an alloy) because the copper and tin atoms are mixed together without being chemically bonded in a fixed ratio. Different bronzes have different ratios of copper to tin. Because it is a mixture, the copper in the bronze still retains some of its own chemical properties. Over thousands of years, the copper on the surface slowly reacted with oxygen, water, and carbon dioxide in the air. This chemical reaction formed copper carbonate (CuCO₃) - a green compound. This is called a patina.
The Big Idea
In a mixture, each substance keeps its own properties. The copper in bronze can still react with air and moisture just like pure copper would - that is why the surface turns green. The green patina is itself a compound (copper carbonate) formed by a chemical reaction.
Test Your Knowledge: 2.1
20 multiple choice questions on Elements, Compounds and Mixtures
Score
0 / 20
QUESTION 1 OF 20
Which of the following is an element?
Iron (Fe) is an element because it contains only one type of atom. Water, carbon dioxide, and salt are all compounds because they contain two or more different types of atoms chemically bonded together.
QUESTION 2 OF 20
A substance contains only one type of atom. It could be:
By definition, an element contains only one type of atom. Compounds contain at least two different types of atoms, and mixtures contain more than one substance.
QUESTION 3 OF 20
What is the key difference between a compound and a mixture?
The key difference is that in a compound, elements are chemically combined (bonded) in a fixed ratio, while in a mixture, substances are simply mixed together without chemical bonding and can be separated by physical methods.
QUESTION 4 OF 20
Which separation method would you use to separate sand from water?
Sand is insoluble in water, so it can be separated by filtration. The sand stays on the filter paper as the residue, while the water passes through as the filtrate. Distillation separates liquids with different boiling points, chromatography separates dissolved substances, and electrolysis is a chemical method.
QUESTION 5 OF 20
Oxygen gas (O₂) is best classified as:
O₂ is an element because it contains only one type of atom (oxygen). Even though there are two atoms bonded together, they are both oxygen atoms. A compound must contain two or more different types of atoms.
QUESTION 6 OF 20
Which of the following is a compound?
Methane (CH₄) is a compound because it contains two different types of atoms (carbon and hydrogen) chemically bonded together. Air and bronze are mixtures. Nitrogen gas (N₂) is an element.
QUESTION 7 OF 20
When sodium and chlorine react to form sodium chloride (NaCl), the product is:
When elements combine to form a compound, the properties of the compound are completely different from those of the individual elements. Sodium is a reactive metal and chlorine is a poisonous gas, but NaCl (table salt) is a harmless white solid.
QUESTION 8 OF 20
Which of the following can be separated by physical methods?
Salt water is a mixture, so it can be separated by physical methods such as distillation or evaporation. Water, carbon dioxide, and calcium carbonate are all compounds - they can only be broken down by chemical reactions, not physical methods.
QUESTION 9 OF 20
The chemical symbol for sodium is Na (from the Latin name Natrium). Which of the following also has a symbol from its Latin name?
Iron has the symbol Fe, which comes from the Latin word "Ferrum." Oxygen (O), Carbon (C), and Nitrogen (N) all have symbols that match their English names. Other elements with Latin-derived symbols include Au (Gold - Aurum), Ag (Silver - Argentum), Cu (Copper - Cuprum), and Pb (Lead - Plumbum).
QUESTION 10 OF 20
A student heats a mixture of iron filings and sulfur powder. After the reaction, what is formed?
When iron and sulfur are heated together, a chemical reaction occurs forming the compound iron sulfide (FeS). The iron and sulfur are now chemically combined. Before heating, you had a mixture that could be separated using a magnet. After heating, the compound FeS cannot be separated by physical methods and is not attracted to a magnet.
QUESTION 11 OF 20
Which statement about mixtures is correct?
In a mixture, the components keep their own properties, can be in any ratio, and can be separated by physical methods (filtration, distillation, chromatography, etc.). No chemical reaction is involved in making or separating a mixture.
QUESTION 12 OF 20
Chromatography is used to separate:
Chromatography separates dissolved substances based on how they move through a medium (like filter paper). It is commonly used to separate the different dyes in ink or food colouring. Filtration separates insoluble solids from liquids. A separating funnel separates immiscible liquids.
QUESTION 13 OF 20
How many elements are present in the compound calcium carbonate (CaCO₃)?
CaCO₃ contains three different elements: calcium (Ca), carbon (C), and oxygen (O). The "3" after the O tells you there are three oxygen atoms, but oxygen is still just one element. Remember: the number of elements is the number of different types of atoms, not the total number of atoms.
QUESTION 14 OF 20
To obtain pure water from sea water, the best method is:
Distillation is used to obtain pure water from sea water. The sea water is heated until the water boils and turns to steam, leaving the dissolved salts behind. The steam is then cooled and condensed back into pure liquid water. Evaporation would give you the salt, not the water. Filtration cannot separate dissolved salts.
QUESTION 15 OF 20
Before heating a mixture of iron filings and sulfur powder, the iron can be removed using a magnet. After heating the mixture to form iron sulfide (FeS), the magnet no longer works. This is because:
After the chemical reaction, iron sulfide (FeS) is a new compound with completely different properties from iron and sulfur. Iron sulfide is not magnetic, even though iron is. This demonstrates the key idea that a compound has different properties from the elements that form it.
QUESTION 16 OF 20
Which of the following is NOT a mixture?
Magnesium oxide (MgO) is a compound, not a mixture. The magnesium and oxygen are chemically combined in a fixed ratio. Brass, sea water, and air are all mixtures - their components are not chemically bonded and can be separated by physical methods.
QUESTION 17 OF 20
A student uses paper chromatography on a sample of black ink. The chromatogram shows four different coloured spots. What does this tell you about the ink?
Each spot on a chromatogram represents a different substance. Four spots mean the ink is a mixture of four different dyes. A pure substance would produce only one spot on a chromatogram.
QUESTION 18 OF 20
A separating funnel is used to separate:
A separating funnel is used to separate two immiscible liquids (liquids that do not mix), such as oil and water. The denser liquid settles to the bottom and can be drained out through the tap. Filtration separates insoluble solids from liquids. Evaporation or distillation separates soluble solids from liquids.
QUESTION 19 OF 20
Which of these formulae represents an element, not a compound?
Cl₂ is an element because it contains only one type of atom (chlorine). The "2" just means there are two chlorine atoms bonded together. H₂SO₄ (sulfuric acid), NH₃ (ammonia), and CaO (calcium oxide) are all compounds because they contain two or more different types of atoms.
QUESTION 20 OF 20
A student has a solution of salt in water. She wants to obtain pure, dry salt crystals. What should she do?
Since the salt is dissolved in the water, filtration will not work (dissolved substances pass through filter paper). To get the salt, you need to evaporate the water by heating the solution in an evaporating dish. The water turns to steam and escapes, leaving pure dry salt crystals behind.
2.2 Atomic Structure and the Periodic Table
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Inside the Atom
For a long time, people thought atoms were the smallest possible things - tiny solid balls that could not be broken down. But scientists discovered that atoms are actually made up of even smaller particles called subatomic particles. There are three types you need to know:
- Protons - found in the nucleus (centre) of the atom. They have a positive charge (+1) and a relative mass of 1.
- Neutrons - also found in the nucleus, right next to the protons. They have no charge (0) and a relative mass of 1.
- Electrons - found orbiting the nucleus in shells (also called energy levels). They have a negative charge (-1) and a negligible mass (about 1/1836 of a proton).
The nucleus is incredibly tiny compared to the whole atom. If the atom were the size of a football stadium, the nucleus would be about the size of a pea in the centre! Most of the atom is empty space where the electrons orbit.
| Particle | Relative Charge | Relative Mass | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | -1 | 1/1836 (negligible) | Shells (orbiting nucleus) |
Structure of a lithium atom (3 protons, 4 neutrons, 3 electrons). The nucleus contains protons and neutrons. Electrons orbit in shells around the nucleus.
Atomic Number and Mass Number
Every element has two important numbers:
Proton Number (Atomic Number), Z
The proton number (also called the atomic number, symbol Z) is the number of protons in the nucleus of an atom. This is the most important number for an element because it defines which element it is. Every atom of carbon has exactly 6 protons. Every atom of oxygen has exactly 8 protons. If you change the number of protons, you change the element.
In a neutral atom (one that has no overall charge), the number of electrons equals the number of protons. So the proton number also tells you how many electrons a neutral atom has.
Nucleon Number (Mass Number), A
The nucleon number (also called the mass number, symbol A) is the total number of protons and neutrons in the nucleus.
A = Z + N
A = nucleon number (mass number)
Z = proton number (atomic number)
N = number of neutrons
This means you can always work out the number of neutrons by rearranging:
Neutrons = A - Z
Number of neutrons = Mass number - Atomic number
In nuclide notation, an atom is written as:
AZX
where A = mass number (top), Z = atomic number (bottom), X = element symbol
For example: 2311Na means sodium with 11 protons and 23 nucleons
🧠 Memory Trick
To remember which number goes where in nuclide notation: "Mass on top, Atomic at the bottom" - or just think "A is Above, Z is below". The mass number (A) is always the bigger number and goes on top.
Electronic Configuration
Electrons do not just float around randomly - they are arranged in shells (also called energy levels) around the nucleus. Each shell can hold a limited number of electrons:
- 1st shell (closest to the nucleus): holds a maximum of 2 electrons
- 2nd shell: holds a maximum of 8 electrons
- 3rd shell: holds a maximum of 8 electrons (for the first 20 elements)
- 4th shell: starts filling after the 3rd shell has 8 electrons
The rules are simple: fill the innermost shell first, then move outward. Only start filling the next shell when the current one is full.
Electronic Configurations of the First 20 Elements
| Element | Symbol | Z | Configuration | Group | Period |
|---|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | 1 | 1 |
| Helium | He | 2 | 2 | 0 (or 18) | 1 |
| Lithium | Li | 3 | 2, 1 | 1 | 2 |
| Beryllium | Be | 4 | 2, 2 | 2 | 2 |
| Boron | B | 5 | 2, 3 | 3 | 2 |
| Carbon | C | 6 | 2, 4 | 4 | 2 |
| Nitrogen | N | 7 | 2, 5 | 5 | 2 |
| Oxygen | O | 8 | 2, 6 | 6 | 2 |
| Fluorine | F | 9 | 2, 7 | 7 | 2 |
| Neon | Ne | 10 | 2, 8 | 0 (or 18) | 2 |
| Sodium | Na | 11 | 2, 8, 1 | 1 | 3 |
| Magnesium | Mg | 12 | 2, 8, 2 | 2 | 3 |
| Aluminium | Al | 13 | 2, 8, 3 | 3 | 3 |
| Silicon | Si | 14 | 2, 8, 4 | 4 | 3 |
| Phosphorus | P | 15 | 2, 8, 5 | 5 | 3 |
| Sulfur | S | 16 | 2, 8, 6 | 6 | 3 |
| Chlorine | Cl | 17 | 2, 8, 7 | 7 | 3 |
| Argon | Ar | 18 | 2, 8, 8 | 0 (or 18) | 3 |
| Potassium | K | 19 | 2, 8, 8, 1 | 1 | 4 |
| Calcium | Ca | 20 | 2, 8, 8, 2 | 2 | 4 |
The Link to the Periodic Table
The electronic configuration of an element tells you exactly where it sits in the Periodic Table:
- Group number = number of electrons in the outer shell (for main group elements). For example, sodium has the configuration 2, 8, 1, so it has 1 electron in its outer shell and is in Group 1.
- Period number = number of occupied electron shells. Sodium has 3 shells that contain electrons, so it is in Period 3.
- Noble gases (Group 0 or Group 18) have full outer shells. This makes them very stable and unreactive. Helium has a full first shell (2 electrons), neon has a full second shell (2, 8), and argon has a full third shell (2, 8, 8).
🧠 Memory Trick
Think of electron shells like rows of seats in a cinema. The first row (closest to the screen/nucleus) is the smallest and has only 2 seats. The second and third rows each have 8 seats. You must fill the front row before moving to the next row - no skipping allowed! If someone asks "which row are you in?" that is the period. If they ask "which seat number in your row?" that is the group.
🧠 Memory Trick
To remember the shell filling pattern 2, 8, 8 for the first 20 elements, think: "2 ate 8" (two ate eight) - then 8 again! First shell takes 2, second takes 8, third takes 8, and the 4th shell starts filling (with K and Ca getting 1 and 2 electrons).
Electron shell diagrams for sodium (2,8,1), chlorine (2,8,7), and calcium (2,8,8,2). Notice how the outer shell electrons determine the group number.
⚠ Exam Tip
A very common exam question gives you the mass number and atomic number and asks for the number of protons, neutrons, and electrons. Always remember: protons = atomic number, electrons = protons (in a neutral atom), and neutrons = mass number - atomic number. Write these three formulas at the top of your exam paper!
⚠ Exam Tip
When drawing electron shell diagrams, examiners are very fussy about getting the right number of electrons in each shell. Always count your dots carefully! Also, do not forget to draw the nucleus and label it with the element symbol or the number of protons. The electrons must be drawn ON the shells (on the circles), not floating between them.
⚠ Exam Tip
Be careful: the mass of an atom is almost entirely in the nucleus (protons + neutrons), because electrons have negligible mass. If a question asks about where most of the mass is concentrated, the answer is always the nucleus.
Worked Example
An atom of sodium is represented as 2311Na. Find the number of protons, neutrons, and electrons in this atom.
Step 1
Identify the numbers. The mass number (A) = 23 (top number). The atomic number (Z) = 11 (bottom number).
Step 2
Number of protons = atomic number = 11
Step 3
Number of electrons = number of protons (in a neutral atom) = 11
Step 4
Number of neutrons = mass number - atomic number = 23 - 11 = 12
Final Answer: 11 protons, 12 neutrons, 11 electrons
Worked Example
Write the electronic configuration of phosphorus (Z = 15). How many electrons are in its outer shell?
Step 1
Phosphorus has Z = 15, so it has 15 electrons to arrange in shells.
Step 2
Fill the 1st shell: 2 electrons. (15 - 2 = 13 electrons remaining)
Step 3
Fill the 2nd shell: 8 electrons. (13 - 8 = 5 electrons remaining)
Step 4
Put the remaining 5 electrons in the 3rd shell.
Step 5
Electronic configuration: 2, 8, 5. There are 5 electrons in the outer shell.
Final Answer: Electronic configuration is 2, 8, 5. There are 5 electrons in the outer shell.
Worked Example
An element has the electronic configuration 2, 8, 7. What group and period is this element in? Name the element.
Step 1
The group number equals the number of electrons in the outer shell. The outer shell has 7 electrons, so the element is in Group 7 (the halogens).
Step 2
The period number equals the number of occupied shells. There are 3 shells with electrons (2, 8, 7), so the element is in Period 3.
Step 3
Total electrons = 2 + 8 + 7 = 17. This means the element has 17 protons (atomic number 17). Looking at the Periodic Table, element 17 is Chlorine (Cl).
Final Answer: Group 7, Period 3. The element is Chlorine (Cl).
Apply It: Real-World Chemistry
Atoms are not just a concept in textbooks - understanding atomic structure has led to some of the most incredible technologies and discoveries in human history. Let us see how.
1
During Diwali, you see fireworks lighting up the sky in brilliant colours - red, green, blue, gold, and white. Different fireworks produce different colours depending on the metal compounds used in them.
How does the electronic configuration of atoms explain why different elements produce different coloured fireworks?
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The Chemistry
When fireworks explode, the heat gives energy to the electrons in the metal atoms. The electrons absorb this energy and jump to higher energy levels (shells further from the nucleus). But they cannot stay there - they quickly fall back down to their original shells. When they fall back, they release that energy as light. The colour of the light depends on exactly how much energy is released, which depends on the electronic configuration of that element. Strontium compounds produce red, barium compounds produce green, copper compounds produce blue, sodium compounds produce yellow, and magnesium or aluminium produce white.
The Big Idea
Every element has a unique electronic configuration, so every element produces a unique colour when its electrons are excited. This is actually how scientists identify elements in distant stars - by looking at the colours of light they emit! It is called emission spectroscopy.
2
Many streets in cities across India and the UK are lit by sodium street lamps, which give off a distinctive yellow-orange glow. These lamps work by passing electricity through sodium vapour.
Why do sodium lamps always produce yellow light, never blue or red?
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The Chemistry
When electricity passes through sodium vapour, the electrons in sodium atoms get excited and jump to higher energy levels. When they fall back to their original shells, they release a very specific amount of energy that corresponds to yellow light (wavelength about 589 nm). Because sodium has a unique electronic configuration (2, 8, 1), the energy gap between its electron shells is always the same, so it always produces exactly the same shade of yellow. This is as predictable as a fingerprint - sodium will always glow yellow, no matter where in the world you find it.
The Big Idea
The colour of light an element emits is determined by its electronic structure. Each element produces its own unique set of colours - like a fingerprint. This is why flame tests work in chemistry: you can identify an unknown metal by the colour of its flame.
3
You have probably seen "neon signs" glowing in bright colours outside shops and restaurants. The original neon signs used neon gas sealed inside glass tubes, and they glow a bright red-orange colour.
Why is neon so stable and unreactive, and how does it produce light?
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The Chemistry
Neon (Ne) has the electronic configuration 2, 8. Its outer shell is completely full with 8 electrons. This full outer shell makes neon extremely stable - it has no need to gain, lose, or share electrons. It does not react with other elements under normal conditions. This is why neon is called a noble gas (Group 0). When electricity is passed through neon gas in a sealed tube, the electrons temporarily jump to higher energy levels and then fall back, releasing red-orange light. Other noble gases produce different colours: argon gives a pale purple, helium gives a pink-orange, and krypton gives a white glow.
The Big Idea
Elements with full outer shells (noble gases) are the most stable and least reactive elements. Understanding electron shell filling explains why some elements are reactive and others are not. The goal of most atoms in chemical reactions is to achieve a full outer shell - just like the noble gases already have.
4
MRI (Magnetic Resonance Imaging) machines in hospitals create detailed images of the inside of your body. The MRI machine works by detecting the behaviour of hydrogen atoms (specifically their protons) inside your body when placed in a powerful magnetic field.
Why are hydrogen atoms particularly useful for MRI, and what property of the proton makes this possible?
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The Chemistry
Your body is about 60% water (H₂O), so hydrogen atoms are everywhere inside you. Hydrogen is the simplest atom - it has just 1 proton and 1 electron (no neutrons in the most common form). The proton has a positive charge (+1) and also has a property called "spin," which makes it act like a tiny magnet. When you lie inside an MRI machine, a powerful magnetic field aligns all the hydrogen protons in your body. Radio waves are then sent in, which knock the protons out of alignment. As the protons snap back into alignment, they release signals that the machine detects and uses to build a detailed image of your organs and tissues.
The Big Idea
The charged subatomic particles inside atoms are not just abstract ideas from a textbook - they have real physical properties that enable life-saving medical technology. The humble proton, sitting in the nucleus of a hydrogen atom, is what makes MRI scans possible.
5
Gold is one of the most prized metals in the world, especially in Indian culture where it plays a central role in weddings and festivals. Its chemical symbol is Au - which does not seem to match the English word "Gold" at all.
Why does gold have the symbol Au instead of Go or Gd, and what does this tell us about the history of chemistry?
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The Chemistry
The symbol Au comes from the Latin word Aurum, which means "shining dawn" - a beautiful name for this glowing yellow metal. When scientists first created the system of chemical symbols in the early 1800s, Latin was the international language of science. Many elements that were known since ancient times kept their Latin-based symbols: Fe for Iron (Ferrum), Ag for Silver (Argentum), Cu for Copper (Cuprum), Pb for Lead (Plumbum), Sn for Tin (Stannum), Na for Sodium (Natrium), and K for Potassium (Kalium). Gold has atomic number 79, meaning it has 79 protons in its nucleus and 79 electrons arranged in shells around it.
The Big Idea
Chemical symbols are an international language that every scientist in the world understands, regardless of whether they speak English, Hindi, Mandarin, or Arabic. The Periodic Table is one of the few truly universal things in all of science - the same 118 elements, with the same symbols, used by every chemist on the planet.
Test Your Knowledge: 2.2
20 multiple choice questions on Atomic Structure and the Periodic Table
Score
0 / 20
QUESTION 1 OF 20
What is the charge of a proton?
A proton has a relative charge of +1. Electrons have a charge of -1, and neutrons have no charge (0).
QUESTION 2 OF 20
Where are protons and neutrons found in an atom?
Protons and neutrons are both found in the nucleus, which is the tiny, dense centre of the atom. Electrons are the particles that orbit the nucleus in shells.
QUESTION 3 OF 20
The relative mass of an electron is approximately:
An electron has a relative mass of about 1/1836 of a proton, which is so small it is considered negligible. This means almost all of the mass of an atom is concentrated in the nucleus (protons + neutrons).
QUESTION 4 OF 20
An atom has a mass number of 40 and an atomic number of 20. How many neutrons does it have?
Number of neutrons = mass number - atomic number = 40 - 20 = 20. This atom is calcium (Ca), which has 20 protons, 20 neutrons, and 20 electrons.
QUESTION 5 OF 20
The atomic number of an element tells you the number of:
The atomic number (Z) tells you the number of protons in the nucleus. It also equals the number of electrons in a neutral atom. The mass number (A) tells you the total number of protons and neutrons.
QUESTION 6 OF 20
What is the electronic configuration of a neutral atom with 12 electrons?
With 12 electrons: first shell gets 2 (max), second shell gets 8 (max), and the remaining 2 go in the third shell. So the configuration is 2, 8, 2. This is magnesium (Mg). Remember: the second shell can only hold a maximum of 8 electrons, not 10.
QUESTION 7 OF 20
An atom has the electronic configuration 2, 8, 3. Which group of the Periodic Table is it in?
The group number equals the number of electrons in the outer shell. The outer shell has 3 electrons, so the element is in Group 3. (This is aluminium, Al.) The element is in Period 3 because it has 3 occupied electron shells.
QUESTION 8 OF 20
Which particle has no electrical charge?
Neutrons have no electrical charge (they are neutral - that is where they get their name!). Protons have a positive charge (+1), and electrons have a negative charge (-1).
QUESTION 9 OF 20
An atom of 3517Cl has:
The bottom number (17) is the atomic number = protons = 17. In a neutral atom, electrons = protons = 17. Neutrons = mass number - atomic number = 35 - 17 = 18.
QUESTION 10 OF 20
What is the maximum number of electrons that can fit in the second electron shell?
The second electron shell can hold a maximum of 8 electrons. The first shell holds a maximum of 2, and the third shell holds a maximum of 8 (for the first 20 elements you need to know).
QUESTION 11 OF 20
Which element has the electronic configuration 2, 8, 8?
2 + 8 + 8 = 18 electrons, which means 18 protons. Element 18 is argon (Ar). It is a noble gas with a full outer shell, making it very stable and unreactive. Neon is 2, 8. Potassium is 2, 8, 8, 1. Calcium is 2, 8, 8, 2.
QUESTION 12 OF 20
In a neutral atom, the number of electrons is always equal to the number of:
In a neutral atom, the positive charges (protons) must balance the negative charges (electrons), so the number of electrons equals the number of protons. If they were not equal, the atom would have an overall charge and would be an ion.
QUESTION 13 OF 20
An element is in Period 3 and Group 1 of the Periodic Table. Its electronic configuration is:
Period 3 means 3 occupied electron shells. Group 1 means 1 electron in the outer shell. So the configuration must be 2, 8, 1. This is sodium (Na). Option A (2, 1) would be Period 2 Group 1 (lithium). Option D (2, 8, 8, 1) would be Period 4 Group 1 (potassium).
QUESTION 14 OF 20
Why are noble gases very unreactive?
Noble gases (Group 0) are very unreactive because they already have full outer electron shells (2 for helium, 8 for the others). They have no need to gain, lose, or share electrons, so they do not form bonds with other atoms under normal conditions.
QUESTION 15 OF 20
Where is most of the mass of an atom concentrated?
Almost all of the mass of an atom is concentrated in the nucleus because protons and neutrons each have a relative mass of 1, while electrons have a negligible mass (1/1836). The nucleus is tiny but extremely dense.
QUESTION 16 OF 20
The electronic configuration of potassium (K) is:
Potassium has 19 electrons: 2 in the first shell, 8 in the second, 8 in the third, and 1 in the fourth shell = 2, 8, 8, 1. The third shell fills to a maximum of 8 before the fourth shell starts filling. Note: 2, 8, 9 is wrong because the third shell can only hold 8 electrons (for the elements you study at IGCSE level).
QUESTION 17 OF 20
An atom has 8 protons, 8 neutrons, and 8 electrons. What is its mass number?
Mass number = protons + neutrons = 8 + 8 = 16. Electrons are NOT included because their mass is negligible. This atom is oxygen-16 (168O).
QUESTION 18 OF 20
Element X has the electronic configuration 2, 8, 6. Element X is most likely:
2 + 8 + 6 = 16 electrons, so this is sulfur (S). It has 6 electrons in its outer shell, placing it in Group 6. Sulfur is a non-metal. It is in Period 3 (3 occupied shells). Elements on the right side of the Periodic Table (Groups 5, 6, 7) are typically non-metals.
QUESTION 19 OF 20
Which of the following pairs of elements are in the same group of the Periodic Table?
Elements in the same group have the same number of electrons in their outer shell. Lithium (2, 1) has 1 outer electron and sodium (2, 8, 1) also has 1 outer electron - they are both in Group 1. The other pairs all have different numbers of outer electrons.
QUESTION 20 OF 20
An atom of aluminium has a mass number of 27 and an atomic number of 13. How many neutrons does it contain?
Number of neutrons = mass number - atomic number = 27 - 13 = 14. Aluminium has 13 protons, 14 neutrons, and 13 electrons. Its electronic configuration is 2, 8, 3, placing it in Group 3, Period 3.
2.3 Isotopes
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What Are Isotopes?
You already know that every atom of a particular element has the same number of protons. That is what makes it that element. But here is a surprising fact: not all atoms of the same element are exactly identical. Some have more neutrons than others. These different versions are called isotopes.
Isotopes are atoms of the same element that have the same proton number (atomic number) but different nucleon numbers (mass numbers). This means they have different numbers of neutrons.
Think of it like this: imagine two brothers from the same family (same element). They have the same surname (same number of protons), but one is heavier than the other (different number of neutrons). They are still brothers - still the same element - but they weigh different amounts.
Nuclide Notation
We write isotopes using nuclide notation. The mass number (nucleon number) goes at the top-left of the element symbol, and the atomic number (proton number) goes at the bottom-left.
mass number →
atomic number →
AZX
A = mass number (protons + neutrons)
Z = atomic number (number of protons)
X = element symbol
Number of neutrons = A - Z
Common Examples of Isotopes
Carbon Isotopes
Carbon has three naturally occurring isotopes:
- Carbon-12 (126C) - 6 protons, 6 neutrons. This is the most common isotope (about 98.9%).
- Carbon-13 (136C) - 6 protons, 7 neutrons. About 1.1% of natural carbon.
- Carbon-14 (146C) - 6 protons, 8 neutrons. Radioactive - used in carbon dating.
Notice: all three have 6 protons (because they are all carbon), but different numbers of neutrons.
Hydrogen Isotopes
Hydrogen is special because its three isotopes actually have their own names:
- Protium (11H) - 1 proton, 0 neutrons. The most common form (99.98%).
- Deuterium (21H) - 1 proton, 1 neutron. Also written as D.
- Tritium (31H) - 1 proton, 2 neutrons. Radioactive. Also written as T.
Chlorine Isotopes
Chlorine has two stable isotopes:
- Chlorine-35 (3517Cl) - 17 protons, 18 neutrons. About 75% of natural chlorine.
- Chlorine-37 (3717Cl) - 17 protons, 20 neutrons. About 25% of natural chlorine.
This is why the relative atomic mass of chlorine on the Periodic Table is 35.5 - it is an average of the two isotopes, weighted by how common each one is.
Both carbon-12 and carbon-14 have 6 protons and 6 electrons (so they are both carbon), but carbon-14 has 2 extra neutrons, making it heavier. Notice the nucleus of C-14 is slightly larger.
Comparison Table: Carbon Isotopes
| Property | Carbon-12 | Carbon-13 | Carbon-14 |
|---|---|---|---|
| Symbol | 126C | 136C | 146C |
| Protons | 6 | 6 | 6 |
| Neutrons | 6 | 7 | 8 |
| Electrons | 6 | 6 | 6 |
| Mass number | 12 | 13 | 14 |
| Electronic configuration | 2, 4 | 2, 4 | 2, 4 |
| Radioactive? | No | No | Yes |
| Natural abundance | 98.9% | 1.1% | Trace |
SUPPLEMENT
Why Isotopes Have the Same Chemical Properties
This is a really important point for exams. Isotopes of the same element have the same chemical properties. Why? Because chemical properties depend on the electrons - specifically, the number and arrangement of electrons in the outer shell.
Since all isotopes of an element have the same number of protons, they also have the same number of electrons (in a neutral atom, protons = electrons). This means they have the same electronic configuration, so they react in exactly the same way.
For example, carbon-12 and carbon-14 both have 6 electrons arranged as 2, 4. They both form 4 covalent bonds. They both react with oxygen to form CO₂. You cannot tell them apart by their chemistry.
Why Isotopes Have Different Physical Properties
However, isotopes have slightly different physical properties because they have different masses. The extra neutrons make the heavier isotope:
- Slightly denser - the heavier atoms pack together with more mass in the same space
- Slightly higher melting and boiling points - heavier particles need a bit more energy to move apart
- Slower rate of diffusion - heavier particles move more slowly (remember, kinetic energy = ½mv², so for the same energy, a heavier particle must move slower)
Calculating Relative Atomic Mass (Ar)
Because most elements exist as a mixture of isotopes, the relative atomic mass you see on the Periodic Table is not a whole number for many elements. It is a weighted average of all the isotopes, taking into account how abundant each one is.
Ar = (mass₁ × %₁ + mass₂ × %₂ + ...) ÷ 100
mass₁, mass₂ = mass numbers of each isotope
%₁, %₂ = percentage abundance of each isotope
The total of all percentages must add up to 100
This is just like calculating a weighted average mark. If you scored 80% in a test worth 75% of your grade and 60% in a test worth 25% of your grade, your average is not simply (80 + 60) ÷ 2 = 70. It is (80 × 75 + 60 × 25) ÷ 100 = 75. The same logic applies to isotopic masses.
Worked Example
Two atoms are represented as 3517Cl and 3717Cl. Explain why these are isotopes. State the number of protons, neutrons, and electrons in each.
Step 1: Check the definition
Isotopes are atoms of the same element with the same proton number but different nucleon numbers. Both atoms have atomic number 17, so they are the same element (chlorine). But their mass numbers are different (35 and 37). So yes, these are isotopes.
Step 2: Find subatomic particles for 3517Cl
Protons = atomic number = 17
Electrons = protons (neutral atom) = 17
Neutrons = mass number - proton number = 35 - 17 = 18
Electrons = protons (neutral atom) = 17
Neutrons = mass number - proton number = 35 - 17 = 18
Step 3: Find subatomic particles for 3717Cl
Protons = atomic number = 17
Electrons = protons (neutral atom) = 17
Neutrons = mass number - proton number = 37 - 17 = 20
Electrons = protons (neutral atom) = 17
Neutrons = mass number - proton number = 37 - 17 = 20
Both are isotopes of chlorine. 35Cl: 17p, 18n, 17e. 37Cl: 17p, 20n, 17e. Same protons and electrons, different neutrons.
Worked Example
SUPPLEMENT Chlorine has two isotopes: 35Cl (75.0% abundant) and 37Cl (25.0% abundant). Calculate the relative atomic mass of chlorine.
Step 1: Write the formula
Ar = (mass₁ × abundance₁ + mass₂ × abundance₂) ÷ 100
Step 2: Substitute the values
Ar = (35 × 75.0 + 37 × 25.0) ÷ 100
Step 3: Calculate
Ar = (2625 + 925) ÷ 100 = 3550 ÷ 100 = 35.5
Answer: The relative atomic mass of chlorine is 35.5
Worked Example
SUPPLEMENT Silicon has three isotopes: 28Si (92.2%), 29Si (4.7%), and 30Si (3.1%). Calculate the relative atomic mass of silicon.
Step 1: Write the formula
Ar = (mass₁ × %₁ + mass₂ × %₂ + mass₃ × %₃) ÷ 100
Step 2: Substitute
Ar = (28 × 92.2 + 29 × 4.7 + 30 × 3.1) ÷ 100
Step 3: Calculate each part
28 × 92.2 = 2581.6
29 × 4.7 = 136.3
30 × 3.1 = 93.0
29 × 4.7 = 136.3
30 × 3.1 = 93.0
Step 4: Add and divide
Ar = (2581.6 + 136.3 + 93.0) ÷ 100 = 2810.9 ÷ 100 = 28.1
Answer: The relative atomic mass of silicon is 28.1 (which matches the Periodic Table value of 28.1)
🧠 Memory Trick
"Same address, different weight" - Isotopes live at the same address on the Periodic Table (same atomic number, same element), but they weigh different amounts (different mass numbers). The postman (chemistry) treats them exactly the same because the address is the same. Only the parcel weighing machine (physical measurements) can tell them apart.
🧠 Memory Trick
To remember the Ar formula, think of it like calculating your weighted exam average. If your maths exam (worth 75%) gave you 35 marks and your science exam (worth 25%) gave you 37 marks, your average is not simply 36. It is (35 × 75 + 37 × 25) ÷ 100 = 35.5. The more common isotope pulls the average towards itself, just like the exam worth more of your grade pulls the average.
⚠ Exam Tip
When asked "What are isotopes?", always include ALL parts of the definition: (1) atoms of the same element, (2) with the same proton number, (3) but different nucleon numbers. Many students lose marks by writing only "same element, different mass" - you need to mention proton number and nucleon number specifically.
⚠ Exam Tip
In Ar calculation questions, always divide by 100 at the end. A very common mistake is forgetting this step and writing a massive number like 3550 instead of 35.5. Also, double-check that your answer makes sense - it should be between the lightest and heaviest isotope masses, and closer to the more abundant one.
⚠ Exam Tip
If a question asks why isotopes have the same chemical properties, your answer must mention electrons. Say: "They have the same number of electrons and the same electronic configuration, so they react in the same way." Do not just say "same protons" - the examiner wants you to link protons → electrons → same chemistry.
Apply It: Real-World Chemistry
Isotopes are not just a textbook concept - they are used in archaeology, medicine, energy, and even keeping your home safe. Let us see how.
1
Archaeologists studying the ancient Indus Valley Civilization in Mohenjo-daro (modern-day Pakistan) wanted to know exactly how old certain wooden artefacts and seeds were. They used a technique called radiocarbon dating, which relies on the isotope carbon-14.
How does carbon-14 help scientists figure out the age of ancient objects?
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The Chemistry
All living things absorb carbon from the atmosphere, including a tiny amount of radioactive carbon-14. While the organism is alive, the ratio of C-14 to C-12 stays constant because it keeps taking in new carbon. But when the organism dies, it stops absorbing carbon. The C-14 then slowly decays (its half-life is about 5,730 years), while C-12 stays the same. By measuring how much C-14 is left compared to C-12, scientists can calculate how long ago the organism died. Less C-14 remaining means the object is older.
The Big Idea
Carbon dating works because C-14 and C-12 are isotopes - they have the same chemical behaviour (so both are absorbed equally by living things) but different nuclear stability (C-14 is radioactive and decays over time). Without isotopes, we would have no way to date ancient civilizations like the Indus Valley, which thrived around 2500 BCE.
2
A patient at a hospital in Mumbai is suspected of having a thyroid problem. The doctor gives her a small dose of iodine-131, a radioactive isotope, and then uses a special camera to take images of her thyroid gland.
Why is a radioactive isotope used instead of normal iodine, and how does it help doctors diagnose the problem?
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The Chemistry
Iodine-131 is an isotope of iodine - it has the same chemical properties as normal iodine (I-127), so the body treats it exactly the same way. The thyroid gland naturally absorbs iodine to make hormones. When the patient swallows I-131, the thyroid absorbs it just like normal iodine. But because I-131 is radioactive, it emits gamma rays that can be detected by a special camera outside the body. This creates an image showing exactly how the thyroid is functioning. An overactive thyroid absorbs too much iodine, while an underactive one absorbs too little.
The Big Idea
Medical tracers work because isotopes have the same chemical properties (so the body processes them normally) but can be detected because of their radioactivity. The isotope acts like an invisible spy inside your body - it goes where normal iodine goes, but it sends out a signal that doctors can track.
3
In nuclear power plants and research reactors, including those in India, a special kind of water called "heavy water" (D₂O) is used. Heavy water looks and tastes almost exactly like normal water, but it contains deuterium instead of ordinary hydrogen.
What makes heavy water "heavy", and why is it useful in nuclear reactors?
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The Chemistry
Normal water is H₂O, where each hydrogen atom has 1 proton and 0 neutrons. Heavy water is D₂O, where each hydrogen is replaced by deuterium (²H), which has 1 proton and 1 neutron. This means each water molecule is about 10% heavier. Heavy water has a slightly higher boiling point (101.4°C vs 100°C), higher density (1.107 g/cm³ vs 1.000 g/cm³), and even freezes at a slightly higher temperature (3.8°C vs 0°C). In nuclear reactors, heavy water is used as a moderator - it slows down fast-moving neutrons to the right speed for sustaining a nuclear chain reaction, without absorbing too many of them.
The Big Idea
Heavy water beautifully demonstrates how isotopes have the same chemical properties (D₂O is still water and behaves chemically like H₂O) but different physical properties (higher density, higher boiling point) because of the extra neutrons. India is one of the world's largest producers of heavy water, with plants in Hazira, Manuguru, and Kota.
4
Almost every home and school in the UK has a small white disc on the ceiling - a smoke detector. Inside it is a tiny amount of americium-241, a radioactive isotope. If you have ever wondered why a smoke detector beeps when toast burns, it is all about isotopes.
How does a radioactive isotope inside a smoke detector actually detect smoke?
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The Chemistry
Americium-241 is an isotope that emits alpha particles. These alpha particles ionise the air molecules inside a small chamber in the detector, creating a tiny electric current that flows continuously. When smoke enters the chamber, the smoke particles absorb or block the alpha particles, reducing the ionisation and causing the current to drop. A microchip detects this drop in current and triggers the alarm. The amount of Am-241 used is incredibly small (less than 1 microgram) and is safely sealed inside the detector.
The Big Idea
This shows isotopes saving lives every day. The radioactive property of Am-241 (emitting alpha particles) is what makes it useful - a non-radioactive isotope of americium would be chemically the same but would not emit the particles needed to ionise air. Different isotopes of the same element can have very different nuclear properties while sharing the same chemistry.
5
India is one of the world's largest producers of mangoes, but mangoes spoil quickly. To extend their shelf life for export to countries like the UK, some mangoes are treated with gamma radiation from cobalt-60, a radioactive isotope. This process is called food irradiation.
How does a radioactive isotope help keep mangoes fresh for longer, and does it make the food radioactive?
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The Chemistry
Cobalt-60 is an isotope of cobalt with 27 protons and 33 neutrons (normal cobalt-59 has 32 neutrons). Co-60 is unstable and emits gamma rays, which are high-energy electromagnetic waves. When mangoes pass through a gamma ray beam, the radiation kills bacteria, moulds, and insects on and inside the fruit. It can also slow down the ripening process by damaging enzymes involved in ripening. Importantly, the food never touches the radioactive source - it only passes through the radiation. The food does NOT become radioactive, just as you do not become radioactive after having an X-ray.
The Big Idea
Food irradiation uses the nuclear property of an isotope (gamma emission from Co-60) rather than its chemical properties. The gamma rays kill microorganisms without leaving any residue or making the food radioactive. India has food irradiation facilities in Mumbai (BARC) and Lasalgaon (near Nashik), helping preserve not just mangoes but also onions, spices, and seafood for both domestic use and export.
Test Yourself: Isotopes
20 questions covering definitions, nuclide notation, properties, and calculations
Your Score:
0 / 20
QUESTION 1 OF 20
What are isotopes?
Isotopes are atoms of the SAME element (same proton number) with DIFFERENT nucleon numbers (mass numbers). This means they differ only in the number of neutrons. Option C is wrong because isotopes have the same number of electrons - it is neutrons that differ.
QUESTION 2 OF 20
An atom of carbon-14 has the symbol 146C. How many neutrons does it have?
Neutrons = mass number - proton number = 14 - 6 = 8. The mass number (14) is the total of protons and neutrons, and the atomic number (6) gives the number of protons.
QUESTION 3 OF 20
Which pair of atoms are isotopes of each other?
Isotopes must be atoms of the SAME element (same atomic number). Only option C has two atoms with the same atomic number (17) but different mass numbers (35 and 37). Option D has the same mass number but different atomic numbers - those are different elements (isobars, not isotopes).
QUESTION 4 OF 20
Hydrogen has three isotopes: protium, deuterium, and tritium. What is the same in all three?
All isotopes of the same element have the same number of protons (that is what makes them the same element). Hydrogen always has 1 proton. The neutrons differ: protium has 0, deuterium has 1, tritium has 2. Option D describes the mass number, which is different for each isotope.
QUESTION 5 OF 20
In nuclide notation, where is the mass number written?
In nuclide notation, the mass number (nucleon number) goes at the top-left and the atomic number (proton number) goes at the bottom-left. Remember: "Mass on top" - the bigger number is usually on top!
QUESTION 6 OF 20
An atom has 8 protons, 8 electrons, and 10 neutrons. What is its nuclide notation?
8 protons means it is oxygen (O). Mass number = protons + neutrons = 8 + 10 = 18. So the notation is 188O. This is oxygen-18, a heavier isotope of oxygen (the common one is oxygen-16 with 8 neutrons).
QUESTION 7 OF 20
Which statement about isotopes is correct?
Isotopes have different mass numbers, which means they have different numbers of neutrons (since protons stay the same). They have the same chemical properties (same electrons), the same number of electrons (in neutral atoms), and the same atomic number (that is what makes them the same element).
QUESTION 8 OF 20
Tritium (31H) is an isotope of hydrogen. How many neutrons does tritium have?
Neutrons = mass number - proton number = 3 - 1 = 2. Tritium is the heaviest hydrogen isotope, with 1 proton and 2 neutrons. (Protium has 0 neutrons, deuterium has 1 neutron.)
QUESTION 9 OF 20
SUPPLEMENT Why do isotopes of the same element have the same chemical properties?
Chemical properties depend on electrons, specifically the electronic configuration. Since isotopes have the same number of protons, they have the same number of electrons (in neutral atoms), and therefore the same electronic configuration. Same electrons = same chemistry.
QUESTION 10 OF 20
SUPPLEMENT Which physical property is different between isotopes of the same element?
Isotopes have different masses, which leads to slightly different physical properties like density, rate of diffusion, and melting/boiling points. Reactivity with acids and number of bonds formed are chemical properties (which are the same for isotopes). Colour is not typically affected by neutron number differences.
QUESTION 11 OF 20
SUPPLEMENT Copper has two isotopes: 63Cu (69.2%) and 65Cu (30.8%). What is the relative atomic mass of copper?
Ar = (63 × 69.2 + 65 × 30.8) ÷ 100 = (4359.6 + 2002.0) ÷ 100 = 6361.6 ÷ 100 = 63.6. Since Cu-63 is more abundant (69.2%), the Ar is closer to 63 than to 65.
QUESTION 12 OF 20
SUPPLEMENT Bromine has two isotopes: 79Br and 81Br. The relative atomic mass of bromine is 79.9. What can you deduce about the abundances?
The Ar of 79.9 is between 79 and 81 but slightly closer to 79. This means 79Br is slightly more abundant. If they were exactly equal, the Ar would be exactly 80.0. Since 79.9 is below 80, there must be slightly more of the lighter isotope. (The actual abundances are about 50.7% 79Br and 49.3% 81Br.)
QUESTION 13 OF 20
An element X has two isotopes. Isotope 1 has 20 protons and 20 neutrons. Isotope 2 has 20 protons and 24 neutrons. What are the mass numbers of these isotopes?
Mass number = protons + neutrons. Isotope 1: 20 + 20 = 40. Isotope 2: 20 + 24 = 44. The element is calcium (Ca), and these are calcium-40 and calcium-44.
QUESTION 14 OF 20
Which of the following is NOT true about carbon-12 and carbon-14?
Carbon-12 and carbon-14 are isotopes, so they have the same protons (6), same electrons (6), and therefore the same electronic configuration (2, 4). They differ only in neutrons (6 vs 8). Option C is false because their electronic configurations are identical.
QUESTION 15 OF 20
SUPPLEMENT Neon has three isotopes: 20Ne (90.5%), 21Ne (0.3%), and 22Ne (9.2%). Calculate the relative atomic mass of neon.
Ar = (20 × 90.5 + 21 × 0.3 + 22 × 9.2) ÷ 100 = (1810 + 6.3 + 202.4) ÷ 100 = 2018.7 ÷ 100 = 20.2 (to 1 d.p.). Since Ne-20 is overwhelmingly abundant (90.5%), the average is very close to 20.
QUESTION 16 OF 20
SUPPLEMENT An element has two isotopes with masses 10 and 11. The relative atomic mass is 10.8. What is the percentage abundance of the heavier isotope?
Let x = % of mass 11 isotope, so (100-x) = % of mass 10 isotope. Then: 10.8 = (10(100-x) + 11x) ÷ 100. Multiply both sides by 100: 1080 = 1000 - 10x + 11x = 1000 + x. So x = 80%. This is boron: B-10 (20%) and B-11 (80%).
QUESTION 17 OF 20
A student says: "Chlorine-35 and chlorine-37 are different elements because they have different mass numbers." Is the student correct?
The student is wrong. What defines an element is the number of protons (atomic number), not the mass number. Both Cl-35 and Cl-37 have 17 protons, so they are both chlorine. They are isotopes - the same element with different numbers of neutrons.
QUESTION 18 OF 20
SUPPLEMENT Heavy water (D₂O) has a higher boiling point than normal water (H₂O). This is because deuterium and hydrogen are:
Deuterium (²H) and protium (¹H) are isotopes of hydrogen. They have the same chemical properties (both form water by bonding with oxygen) but different physical properties because deuterium is heavier. This makes D₂O denser and gives it a higher boiling point (101.4°C vs 100°C).
QUESTION 19 OF 20
SUPPLEMENT The relative atomic mass of magnesium is 24.3. Magnesium has three isotopes: 24Mg (79%), 25Mg (10%), and 26Mg (11%). Which calculation gives the correct Ar?
Ar = sum of (each isotope mass × its percentage) ÷ 100. Option A is wrong because it gives a simple average without considering abundances. Option B correctly uses the weighted average formula. Check: (1896 + 250 + 286) ÷ 100 = 2432 ÷ 100 = 24.3. Correct!
QUESTION 20 OF 20
SUPPLEMENT An element has a relative atomic mass of 63.6. It has two isotopes with mass numbers 63 and 65. What is the approximate percentage of the lighter isotope?
Let x = percentage of mass 63. Then: (63x + 65(100-x)) ÷ 100 = 63.6. So 63x + 6500 - 65x = 6360. Therefore -2x = -140, so x = 70%. The lighter isotope (mass 63) makes up 70% of the element. This is copper, matching the data used earlier.
2.4 Ions and Ionic Bonding
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What Are Ions?
You know that atoms are electrically neutral - they have equal numbers of protons (+) and electrons (-). But atoms are not always content to stay that way. Most atoms would be much more stable if they had a full outer electron shell, like the noble gases (Group 0/VIII). To achieve this, atoms can either lose or gain electrons.
When an atom loses or gains electrons, it becomes charged - it is no longer neutral. A charged atom is called an ion.
Metals Form Positive Ions (Cations)
Metal atoms lose electrons to form positive ions (also called cations). When an atom loses electrons, it has more protons than electrons, so the overall charge becomes positive.
- Sodium (Na) has the electronic configuration 2, 8, 1. It loses its 1 outer electron to become Na⁺ with configuration 2, 8 (same as neon).
- Magnesium (Mg) has configuration 2, 8, 2. It loses 2 outer electrons to become Mg²⁺ with configuration 2, 8 (same as neon).
- Aluminium (Al) has configuration 2, 8, 3. It loses 3 outer electrons to become Al³⁺ with configuration 2, 8 (same as neon).
- Calcium (Ca) has configuration 2, 8, 8, 2. It loses 2 outer electrons to become Ca²⁺ with configuration 2, 8, 8 (same as argon).
Notice the pattern: the number of electrons lost equals the group number of the metal in the Periodic Table. Group I metals lose 1 electron (charge +1), Group II lose 2 (charge +2), Group III lose 3 (charge +3).
Non-Metals Form Negative Ions (Anions)
Non-metal atoms gain electrons to form negative ions (also called anions). When an atom gains electrons, it has more electrons than protons, so the overall charge becomes negative.
- Chlorine (Cl) has configuration 2, 8, 7. It gains 1 electron to become Cl⁻ with configuration 2, 8, 8 (same as argon).
- Oxygen (O) has configuration 2, 6. It gains 2 electrons to become O²⁻ with configuration 2, 8 (same as neon).
- Fluorine (F) has configuration 2, 7. It gains 1 electron to become F⁻ with configuration 2, 8 (same as neon).
- Nitrogen (N) has configuration 2, 5. It gains 3 electrons to become N³⁻ with configuration 2, 8 (same as neon).
The pattern here: the number of electrons gained = 8 minus the group number. Group VII gains 1, Group VI gains 2, Group V gains 3.
What Is an Ionic Bond?
An ionic bond is the strong electrostatic attraction between oppositely charged ions. That is the full IGCSE definition - learn it word for word.
Ionic bonds form when a metal atom transfers its outer electrons to a non-metal atom. The metal becomes a positive ion, the non-metal becomes a negative ion, and the opposite charges attract each other strongly.
It is like a transaction: the metal "gives" electrons and the non-metal "takes" them. Both end up with full outer shells and both end up happy.
Dot-and-Cross Diagram: Sodium Chloride (NaCl)
Let us walk through the most important dot-and-cross diagram in IGCSE Chemistry: the formation of sodium chloride.
Na (2,8,1) transfers its outer electron to Cl (2,8,7). Na becomes Na⁺ (2,8) and Cl becomes Cl⁻ (2,8,8). Both now have full outer shells. The dot (purple) among the crosses (blue) on Cl⁻ shows that the electron came from Na.
Properties of Ionic Compounds
Ionic compounds have very distinctive properties, and you need to know these for the exam:
| Property | Ionic Compounds | Why? |
|---|---|---|
| Melting and boiling points | High (typically above 500°C) | Strong electrostatic forces between many ions need lots of energy to overcome |
| Electrical conductivity (solid) | Do NOT conduct | Ions are held in fixed positions - they cannot move to carry charge |
| Electrical conductivity (molten) | DO conduct | When melted, ions become free to move and carry charge |
| Electrical conductivity (dissolved in water) | DO conduct | Ions are free to move in solution and carry charge |
| Solubility | Many are soluble in water | Water molecules are polar and can surround and separate the ions |
| State at room temperature | Solid | Strong forces hold ions in a rigid structure |
The key thing to remember is: for electricity to flow, charged particles must be free to move. In a solid ionic compound, the ions are locked in place. When you melt it or dissolve it, the ions can finally move around, so it conducts.
SUPPLEMENT
Giant Ionic Lattice
Ionic compounds do not exist as individual pairs of ions. Instead, they form a giant ionic lattice - a regular, repeating 3D arrangement of alternating positive and negative ions extending in all directions. There are no individual molecules - just a continuous structure of billions upon billions of ions.
Imagine a 3D chessboard where every white square is a positive ion and every black square is a negative ion, extending outwards in every direction. Every positive ion is surrounded by negative ions, and every negative ion is surrounded by positive ions. The electrostatic attraction acts between ALL neighbouring ions, not just one pair.
The giant ionic lattice of NaCl. Each Na⁺ is surrounded by Cl⁻ ions and vice versa. The Cl⁻ ions are drawn larger because gaining an electron increases the size of the ion, while Na⁺ ions are smaller because losing an electron decreases the ion size.
More Dot-and-Cross Diagrams
Magnesium Oxide (MgO)
Magnesium (Group II) loses 2 electrons. Oxygen (Group VI) gains 2 electrons. A perfect match - one Mg atom donates 2 electrons to one O atom.
Mg (2,8,2) transfers its 2 outer electrons to O (2,6). Both end up with a full outer shell of 8 electrons (like neon). The two purple dots among the blue crosses on O²⁻ show the transferred electrons.
Calcium Chloride (CaCl₂)
Calcium (Group II) loses 2 electrons, but each chlorine (Group VII) can only accept 1 electron. So one calcium atom must donate its electrons to two chlorine atoms - giving 1 electron to each. This is why the formula is CaCl₂, not CaCl.
Sodium Oxide (Na₂O)
This is the opposite problem. Each sodium (Group I) can only donate 1 electron, but oxygen (Group VI) needs 2 electrons. So two sodium atoms are needed to supply 2 electrons to one oxygen atom. The formula is Na₂O.
Magnesium Chloride (MgCl₂)
Magnesium (Group II) loses 2 electrons, each chlorine needs 1. So one magnesium donates 1 electron to each of two chlorine atoms. The formula is MgCl₂.
Explaining Properties in Terms of Structure
Why are melting and boiling points high?
In a giant ionic lattice, there are many strong electrostatic forces acting between all the oppositely charged ions. These forces act in all directions throughout the lattice. It takes a large amount of energy to overcome all these attractions and separate the ions, so the melting and boiling points are very high. NaCl, for example, melts at 801°C.
Why does it conduct when molten or dissolved but not when solid?
In a solid ionic compound, the ions are in fixed positions in the lattice. They can vibrate but they cannot move from place to place. Since they cannot move, they cannot carry an electric current.
When the ionic compound is melted or dissolved in water, the lattice breaks apart and the ions become free to move. These mobile ions can carry charge through the liquid, so it conducts electricity.
Why are ionic compounds brittle?
If a force pushes one layer of ions sideways in the lattice, ions of like charge (positive next to positive, negative next to negative) end up next to each other. Like charges repel, and the crystal shatters along that plane. This is why salt crystals crack cleanly when you hit them - they are brittle.
Worked Example
Draw a dot-and-cross diagram to show the formation of sodium chloride (NaCl). Show the electronic configurations of the atoms and the ions formed.
Step 1: Write the electronic configurations
Sodium (Na): 2, 8, 1 (11 electrons)
Chlorine (Cl): 2, 8, 7 (17 electrons)
Chlorine (Cl): 2, 8, 7 (17 electrons)
Step 2: Identify the electron transfer
Na has 1 electron in its outer shell - it needs to lose this 1 electron to get a full outer shell (2, 8).
Cl has 7 electrons in its outer shell - it needs to gain 1 electron to get a full outer shell (2, 8, 8).
So Na gives its 1 outer electron to Cl. A perfect match!
Cl has 7 electrons in its outer shell - it needs to gain 1 electron to get a full outer shell (2, 8, 8).
So Na gives its 1 outer electron to Cl. A perfect match!
Step 3: Draw the atoms
Draw Na with 2, 8, 1 electrons using dots. Draw Cl with 2, 8, 7 electrons using crosses. Show an arrow from Na's outer electron to Cl's outer shell.
Step 4: Draw the ions
Na⁺ now has configuration 2, 8 (all dots). Put [Na]⁺ with a plus sign.
Cl⁻ now has configuration 2, 8, 8 (crosses plus one dot from Na). Put [Cl]⁻ with a minus sign.
The transferred electron should be shown as a dot among the crosses on Cl⁻.
Cl⁻ now has configuration 2, 8, 8 (crosses plus one dot from Na). Put [Cl]⁻ with a minus sign.
The transferred electron should be shown as a dot among the crosses on Cl⁻.
Step 5: Add charges and brackets
Put square brackets around each ion and write the charge outside: [Na]⁺ and [Cl]⁻.
Answer: Na (2,8,1) transfers 1 electron to Cl (2,8,7). Na becomes Na⁺ (2,8) and Cl becomes Cl⁻ (2,8,8). Both have full outer shells. The ionic bond is the electrostatic attraction between Na⁺ and Cl⁻.
Worked Example
SUPPLEMENT Draw a dot-and-cross diagram for the formation of calcium chloride (CaCl₂). Explain why two chlorine atoms are needed.
Step 1: Electronic configurations
Calcium (Ca): 2, 8, 8, 2 (20 electrons, Group II)
Chlorine (Cl): 2, 8, 7 (17 electrons, Group VII)
Chlorine (Cl): 2, 8, 7 (17 electrons, Group VII)
Step 2: Work out the electron transfer
Ca needs to lose 2 electrons (to get 2, 8, 8 = argon configuration).
Each Cl can only accept 1 electron (to get 2, 8, 8).
So Ca must donate 1 electron to each of 2 Cl atoms. That is why the formula is CaCl₂.
Each Cl can only accept 1 electron (to get 2, 8, 8).
So Ca must donate 1 electron to each of 2 Cl atoms. That is why the formula is CaCl₂.
Step 3: Draw it
Draw Ca in the middle with 2, 8, 8, 2 electrons (dots). Draw two Cl atoms on either side, each with 2, 8, 7 electrons (crosses). Show arrows: one electron from Ca goes to Cl on the left, one goes to Cl on the right.
Step 4: Result
Ca loses 2 electrons → Ca²⁺ (2, 8, 8)
Each Cl gains 1 electron → 2 × Cl⁻ (each 2, 8, 8)
Each Cl⁻ has 7 crosses + 1 dot in its outer shell.
Each Cl gains 1 electron → 2 × Cl⁻ (each 2, 8, 8)
Each Cl⁻ has 7 crosses + 1 dot in its outer shell.
Answer: Ca (2,8,8,2) donates 1 electron to each of 2 Cl atoms. This gives Ca²⁺ (2,8,8) and 2 × Cl⁻ (each 2,8,8). Two Cl atoms are needed because Ca has 2 electrons to lose, but each Cl can only accept 1.
Worked Example
Predict the charge on the ion formed by each of these elements: (a) Potassium (Group I) (b) Sulfur (Group VI) (c) Aluminium (Group III) (d) Bromine (Group VII)
The Rule
Metals (Groups I, II, III) LOSE electrons. The charge = the group number.
Non-metals (Groups V, VI, VII) GAIN electrons. The charge = -(8 minus group number).
Non-metals (Groups V, VI, VII) GAIN electrons. The charge = -(8 minus group number).
(a) Potassium (Group I)
Metal, so it loses electrons. Group I = loses 1 electron. Charge = K⁺ (1+)
(b) Sulfur (Group VI)
Non-metal, so it gains electrons. 8 - 6 = 2 electrons gained. Charge = S²⁻ (2-)
(c) Aluminium (Group III)
Metal, so it loses electrons. Group III = loses 3 electrons. Charge = Al³⁺ (3+)
(d) Bromine (Group VII)
Non-metal, so it gains electrons. 8 - 7 = 1 electron gained. Charge = Br⁻ (1-)
Answers: (a) K⁺ (b) S²⁻ (c) Al³⁺ (d) Br⁻
🧠 Memory Trick
OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons). When sodium LOSES an electron, it is oxidised. When chlorine GAINS an electron, it is reduced. OIL RIG is one of the most useful mnemonics in all of chemistry!
🧠 Memory Trick
Metals are generous givers, non-metals are greedy takers. Metals have only a few outer electrons, so it is easier for them to give them away. Non-metals are close to having a full shell, so they would rather grab a few more electrons than give up the ones they already have. Think of metals as the kid who shares their lunch, and non-metals as the kid who takes food from everyone else's plate!
🧠 Memory Trick
CAT-ions are PAWS-itive. A cat (cat-ion) is positive. An anion is negative - "a negative ion." This helps you remember which is which.
⚠ Exam Tip
In dot-and-cross diagrams, you MUST use different symbols for electrons from different atoms. Use dots for one element and crosses for the other. When an electron is transferred, show it as a dot among crosses (or vice versa) on the receiving ion. This is how the examiner knows you understand that the electron came from the other atom.
⚠ Exam Tip
When explaining why an ionic compound conducts electricity when molten but not when solid, use the phrase "ions are free to move". This is the key phrase examiners look for. Do NOT say "electrons are free to move" - that is metallic bonding! In ionic compounds, it is the IONS that carry the charge, not electrons.
⚠ Exam Tip
Always put square brackets around ions in dot-and-cross diagrams with the charge written outside the bracket. For example: [Na]⁺ and [Cl]⁻. Forgetting the brackets or the charges costs marks in nearly every exam series.
Apply It: Real-World Chemistry
Ionic compounds are everywhere in your daily life - from the salt on your dinner table to the roads you drive on in winter. Let us see ionic bonding in action.
1
Table salt (sodium chloride, NaCl) is something you use every day. But sodium is a silvery metal that explodes violently when dropped in water, and chlorine is a poisonous yellow-green gas that was used as a chemical weapon in World War I. Yet when these two dangerous elements combine, they make something safe enough to sprinkle on your dal and rice.
How do a deadly metal and a poisonous gas combine to make safe, tasty table salt?
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The Chemistry
When sodium reacts with chlorine, each sodium atom transfers its one outer electron to a chlorine atom. Sodium becomes Na⁺ (a stable ion with a full outer shell like neon) and chlorine becomes Cl⁻ (a stable ion with a full outer shell like argon). These ions are completely different from the original atoms. The dangerous reactivity of sodium came from its desperate need to lose that lone outer electron, and chlorine's toxicity came from its reactive nature as it sought one more electron. Once both have achieved full outer shells as ions, they become completely stable and unreactive. The Na⁺ and Cl⁻ ions then arrange themselves in a giant ionic lattice, held together by strong electrostatic attraction.
The Big Idea
A compound has completely different properties from the elements that form it. The ionic bond transforms two dangerous elements into a harmless, essential compound. This is one of the most dramatic examples in chemistry of how forming bonds changes everything.
2
In many Indian households, Epsom salt (magnesium sulfate, MgSO₄) is dissolved in warm bath water to relieve sore muscles after a long day. When the white crystals are added to water, they dissolve completely to form a clear solution.
Why do ionic compounds like Epsom salt dissolve so well in water, and what happens to the ions when it dissolves?
▼
The Chemistry
MgSO₄ is an ionic compound containing Mg²⁺ ions and SO₄²⁻ ions arranged in a giant ionic lattice. Water is a polar molecule - the oxygen end is slightly negative and the hydrogen ends are slightly positive. When you drop Epsom salt into water, the water molecules surround the ions: the negative oxygen ends of water molecules are attracted to the Mg²⁺ ions, and the positive hydrogen ends are attracted to the SO₄²⁻ ions. These attractions are strong enough to pull the ions out of the lattice and keep them separated in solution. The ions become "hydrated" - surrounded by a shell of water molecules.
The Big Idea
When an ionic compound dissolves, the lattice breaks apart and the ions become free to move in the solution. This is why salt water conducts electricity but solid salt does not - the dissolved ions can carry charge. The same principle applies to your Epsom salt bath!
3
In a school science lab, a student sets up a circuit with a light bulb and tries to pass electricity through a block of rock salt. The bulb does not light up. She then dissolves the salt in water and dips the electrodes into the salt solution. The bulb glows brightly.
Why does solid rock salt not conduct electricity, but salt water does?
▼
The Chemistry
In solid rock salt, the Na⁺ and Cl⁻ ions are locked into fixed positions in a giant ionic lattice. They can vibrate slightly but cannot move from their positions. For electricity to flow, charged particles must be able to move through the substance and carry charge. Since the ions are stuck, no current flows and the bulb stays dark. When the salt dissolves in water, the lattice breaks apart and the ions become free to move. The Na⁺ ions move towards the negative electrode and the Cl⁻ ions move towards the positive electrode. This movement of ions is what carries the electric current, making the bulb glow.
The Big Idea
The same substance - NaCl - can be a conductor or an insulator depending on its state. The key is whether the ions are free to move. Solid = ions fixed = insulator. Dissolved or molten = ions free = conductor. This is a favourite exam question, so remember the phrase "ions are free to move."
4
During cold winters in the UK, councils spread calcium chloride (CaCl₂) or sodium chloride (NaCl) on roads and pavements to prevent ice from forming. Even when temperatures drop below 0°C, the roads stay ice-free.
Why does spreading an ionic compound on roads prevent ice from forming?
▼
The Chemistry
When CaCl₂ dissolves in the thin film of water on the road surface, it dissociates into Ca²⁺ and 2Cl⁻ ions. These ions interfere with the ability of water molecules to form the regular crystalline structure of ice. To freeze, water molecules need to arrange themselves into a rigid hexagonal pattern, but the dissolved ions get in the way and disrupt this arrangement. This means the water has to be cooled to a lower temperature before it can freeze. CaCl₂ can lower the freezing point to about -29°C, which is more effective than NaCl (which only works down to about -9°C) because CaCl₂ produces three ions per formula unit (1 Ca²⁺ + 2 Cl⁻) compared to NaCl's two ions (1 Na⁺ + 1 Cl⁻).
The Big Idea
When ionic compounds dissolve, they break into individual ions. More ions in solution means a greater effect on the freezing point. CaCl₂ is more effective than NaCl because each formula unit produces 3 ions instead of 2. This is why CaCl₂ is preferred for colder conditions and is worth the extra cost.
5
LED lights have transformed how we light our homes and streets, using a fraction of the electricity of old incandescent bulbs. At the heart of many LEDs are compounds like gallium nitride (GaN) and gallium arsenide (GaAs). Gallium is a metal that can melt in your hand (melting point 29.8°C), but its compounds are incredibly useful in technology.
How does ionic bonding in gallium compounds make LED technology possible?
▼
The Chemistry
Gallium is a Group III metal, so it can lose 3 electrons to form Ga³⁺ ions. Nitrogen is a Group V non-metal that can gain 3 electrons to form N³⁻ ions. In gallium nitride (GaN), the bonding has both ionic and covalent character. The strong bonds create a very stable crystal structure with specific electronic properties. When electricity flows through the crystal, electrons jump between energy levels and release energy as light. The colour of light depends on the exact composition - GaN produces blue and white light, while GaAs produces infrared and red light. By mixing different compounds, engineers can create LEDs of any colour.
The Big Idea
The properties of ionic and partially ionic compounds depend on their structure and bonding. By changing which metals and non-metals are combined, scientists can engineer materials with very specific properties. The LED revolution - which saves billions of kilowatt-hours of energy worldwide - is built on our understanding of how atoms transfer and share electrons to form bonds.
Test Yourself: Ions and Ionic Bonding
20 questions covering ion formation, electron transfer, dot-and-cross diagrams, and properties
Your Score:
0 / 20
QUESTION 1 OF 20
What happens when a sodium atom forms a sodium ion?
Sodium is a metal in Group I, so it LOSES its 1 outer electron to form a positive ion Na⁺. Ions are formed by gaining or losing ELECTRONS, never protons. If an atom lost a proton, it would become a different element entirely.
QUESTION 2 OF 20
What is the electronic configuration of a chloride ion (Cl⁻)?
Chlorine has configuration 2, 8, 7. When it gains 1 electron to form Cl⁻, it becomes 2, 8, 8 - the same electronic configuration as argon (a noble gas). This full outer shell is what makes the ion stable.
QUESTION 3 OF 20
What is an ionic bond?
An ionic bond is the strong electrostatic attraction between oppositely charged ions. Learn this definition word for word - it comes up in almost every exam. Option A describes a covalent bond, and option B describes metallic bonding.
QUESTION 4 OF 20
Magnesium is in Group II. What ion does it form?
Magnesium is a metal in Group II, so it loses its 2 outer electrons to form Mg²⁺. The charge on a metal ion equals its group number: Group I = 1+, Group II = 2+, Group III = 3+.
QUESTION 5 OF 20
Which of these correctly shows an electron transfer in ionic bonding?
In ionic bonding, metal atoms LOSE (give) electrons to non-metal atoms. The metal becomes a positive ion and the non-metal becomes a negative ion. Protons are never transferred - they stay in the nucleus.
QUESTION 6 OF 20
An oxide ion (O²⁻) has 10 electrons. How many protons does it have?
Oxygen always has 8 protons (atomic number 8) - this never changes when ions form. The O²⁻ ion has gained 2 electrons (8 + 2 = 10 electrons) but still has 8 protons. The charge is 2- because there are 2 more electrons than protons (10 - 8 = 2).
QUESTION 7 OF 20
Why do ionic compounds have high melting points?
Ionic compounds form giant lattices with strong electrostatic attractions acting between all the oppositely charged ions. It takes a large amount of energy to overcome these many strong forces, so the melting point is high. Key words: "strong electrostatic forces" and "lots of energy."
QUESTION 8 OF 20
Solid sodium chloride does not conduct electricity. Why not?
Solid NaCl DOES contain ions (Na⁺ and Cl⁻), but they are locked in fixed positions in the giant ionic lattice. They cannot move, so they cannot carry an electric current. When melted or dissolved, the ions become free to move and the compound conducts.
QUESTION 9 OF 20
When molten sodium chloride conducts electricity, what carries the charge?
In molten ionic compounds, it is the IONS (both positive and negative) that move and carry the electric charge. This is different from metals, where electrons carry the charge. Protons and neutrons are locked in nuclei and cannot move to carry current.
QUESTION 10 OF 20
The formula of calcium chloride is CaCl₂. Why are two chlorine atoms needed for each calcium atom?
Calcium (Group II) must lose 2 electrons to get a full outer shell. Each chlorine (Group VII) can only accept 1 electron. So 2 Cl atoms are needed to accept the 2 electrons from 1 Ca atom. Ca gives 1 electron to each Cl: Ca²⁺ + 2Cl⁻.
QUESTION 11 OF 20
Which of these substances would conduct electricity?
Only molten potassium bromide would conduct because the ions (K⁺ and Br⁻) are free to move in the liquid state. All the solid ionic compounds (options A, B, D) have ions locked in fixed positions, so they cannot conduct.
QUESTION 12 OF 20
What is the charge on an aluminium ion?
Aluminium is in Group III, so it has 3 outer electrons. It loses all 3 to form Al³⁺. The charge on a metal ion equals the group number: Group I = 1+, Group II = 2+, Group III = 3+.
QUESTION 13 OF 20
In a dot-and-cross diagram for NaCl, why is the transferred electron shown as a dot among crosses on the chloride ion?
In reality, all electrons are identical. We use dots and crosses purely to show where each electron CAME FROM. The dot among the crosses on Cl⁻ shows that this particular electron was transferred from Na. It helps the examiner see that you understand the electron transfer process.
QUESTION 14 OF 20
The Na⁺ ion has the same electronic configuration as which noble gas?
Na (2, 8, 1) loses 1 electron to become Na⁺ (2, 8). This is the same electronic configuration as neon, which is 2, 8. Ions achieve stability by getting the same electron arrangement as the nearest noble gas.
QUESTION 15 OF 20
SUPPLEMENT Why is the formula of sodium oxide Na₂O?
Oxygen (Group VI) needs to gain 2 electrons to achieve a full outer shell. Each sodium (Group I) can only donate 1 electron. So 2 Na atoms are needed to provide the 2 electrons that 1 O atom needs: 2Na⁺ + O²⁻ = Na₂O.
QUESTION 16 OF 20
SUPPLEMENT What is the structure of a solid ionic compound?
Ionic compounds form giant ionic lattices - a regular 3D arrangement of alternating positive and negative ions extending in all directions. There are no individual molecules. Option C describes metallic bonding, and option A describes simple molecular structures.
QUESTION 17 OF 20
SUPPLEMENT Why are ionic compounds brittle?
When a force pushes one layer of ions in the lattice, ions of like charge (+ next to + or - next to -) end up next to each other. Like charges repel, and this strong repulsion causes the crystal to shatter along that plane. The bonds are not weak - it is the arrangement that causes brittleness.
QUESTION 18 OF 20
A substance has a high melting point, does not conduct electricity as a solid, but conducts when dissolved in water. What type of substance is it most likely to be?
This is a classic property identification question. High melting point = strong forces (not covalent molecules or gas). Does not conduct as solid but conducts when dissolved = ionic compound (ions fixed when solid, free to move when dissolved). A metal would conduct as a solid too. Simple covalent compounds have low melting points.
QUESTION 19 OF 20
SUPPLEMENT In magnesium chloride (MgCl₂), how many electrons does each magnesium atom transfer, and to how many chlorine atoms?
Mg (Group II) has 2 outer electrons to lose. Each Cl (Group VII) needs 1 electron. So Mg donates 1 electron to each of 2 Cl atoms - a total of 2 electrons transferred. This gives Mg²⁺ and 2Cl⁻, hence the formula MgCl₂.
QUESTION 20 OF 20
SUPPLEMENT Which statement correctly explains why molten sodium chloride conducts electricity?
When NaCl is melted, the giant ionic lattice breaks down. The Na⁺ and Cl⁻ ions are no longer held in fixed positions - they become free to move throughout the liquid. These moving ions carry the electric charge. Note: ionic compounds do NOT have delocalised electrons (that is metals) and do NOT contain molecules (they contain ions).
2.5 Simple Molecules and Covalent Bonds
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What Is a Covalent Bond?
In ionic bonding, atoms transfer electrons. But what happens when two non-metal atoms meet? Neither one wants to give up electrons — they both want to gain electrons. So instead of transferring, they share electrons.
A covalent bond is a shared pair of electrons between two atoms. By sharing, both atoms can count the shared electrons as part of their outer shell, so both achieve a full outer shell. It is like two people sharing a textbook — both can read it at the same time.
Covalent bonds form between non-metal atoms. The shared electrons are attracted to the nuclei of both atoms, and this attraction holds the atoms together.
Dot-and-Cross Diagrams for Simple Molecules
Just like with ionic bonding, we use dot-and-cross diagrams to show covalent bonding. Dots represent electrons from one atom, and crosses represent electrons from the other. Where a dot and a cross sit together between two atoms, that is a shared pair — a covalent bond.
Hydrogen (H₂)
Each hydrogen atom has 1 electron and needs 2 to fill its first shell (like helium). Two hydrogen atoms share their electrons so both have 2.
Two hydrogen atoms each contribute 1 electron to form a shared pair. Both atoms now have 2 electrons in their outer shell (like helium).
Chlorine (Cl₂)
Each chlorine atom has 7 outer electrons and needs 8. They share one pair of electrons, giving both a full outer shell of 8.
Each Cl has 7 outer electrons. They share one pair, giving each atom 8 outer electrons (a full shell). The 6 non-bonding electrons on each atom are called lone pairs.
Water (H₂O)
Oxygen has 6 outer electrons and needs 8. It shares one pair with each of two hydrogen atoms. This gives oxygen 8 outer electrons and each hydrogen 2.
Oxygen (2, 6) forms 2 covalent bonds with 2 hydrogen atoms. Oxygen ends up with 8 outer electrons (2 bonding pairs + 2 lone pairs). Each H has 2 electrons.
Methane (CH₄)
Carbon has 4 outer electrons and needs 8. It shares one pair with each of four hydrogen atoms. This gives carbon 8 outer electrons and each hydrogen 2. Carbon forms 4 covalent bonds — the maximum for a Period 2 element.
Carbon (2, 4) forms 4 covalent bonds with 4 hydrogen atoms. Carbon has no lone pairs — all 8 outer electrons are in bonding pairs.
Ammonia (NH₃)
Nitrogen has 5 outer electrons and needs 8. It shares one pair with each of three hydrogen atoms. This gives nitrogen 8 outer electrons (3 bonding pairs + 1 lone pair) and each hydrogen 2.
Hydrogen Chloride (HCl)
Hydrogen has 1 outer electron and chlorine has 7. They share one pair, giving hydrogen 2 and chlorine 8 outer electrons.
Properties of Simple Covalent Molecules
| Property | Simple Covalent Molecules | Why? |
|---|---|---|
| Melting and boiling points | Low | Weak intermolecular forces between molecules need little energy to overcome |
| Electrical conductivity | Do NOT conduct | No charged particles (ions or free electrons) are available to carry charge |
| State at room temperature | Often gases or liquids, some are low-melting-point solids | Weak forces between molecules means they are easily separated |
⚠ Exam Tip
This is one of the most common mistakes in IGCSE Chemistry: when explaining why simple covalent substances have low melting points, students often say "the covalent bonds are weak." This is WRONG. Covalent bonds are very strong! The intermolecular forces (forces between the molecules) are weak. When a substance like water boils, the covalent O-H bonds inside each molecule do NOT break — the weak forces between the molecules are overcome.
⚠ Exam Tip
Simple covalent substances do not conduct electricity because they have no ions and no free electrons. The electrons are all locked in covalent bonds. Even when melted or dissolved, there are no charged particles to carry a current (unlike ionic compounds where ions become free).
SUPPLEMENT
Multiple Bonds: Double and Triple Bonds
Sometimes sharing one pair of electrons is not enough to give both atoms a full outer shell. In these cases, atoms share two pairs (a double bond) or even three pairs (a triple bond).
Oxygen (O₂) — Double Bond
Each oxygen atom has 6 outer electrons and needs 8. By sharing two pairs of electrons, both oxygen atoms achieve 8 outer electrons.
Each O contributes 2 electrons to the bond. Both atoms have 8 outer electrons: 2 bonding pairs + 2 lone pairs each.
Carbon Dioxide (CO₂) — Two Double Bonds
Carbon has 4 outer electrons. It forms a double bond with each of two oxygen atoms (sharing 2 pairs with each). This gives carbon 8 outer electrons (all in bonding pairs) and each oxygen 8 outer electrons.
Nitrogen (N₂) — Triple Bond
Each nitrogen atom has 5 outer electrons and needs 8. They share three pairs, forming a very strong triple bond. This is why nitrogen gas is so unreactive — the triple bond is extremely difficult to break.
Each N contributes 3 electrons to the bond. Both atoms have 8 outer electrons: 3 bonding pairs + 1 lone pair each. The triple bond makes N₂ one of the most stable molecules in chemistry.
Ethene (C₂H₄) — Carbon-Carbon Double Bond
Ethene has a double bond between the two carbon atoms. Each carbon also bonds to two hydrogen atoms. The double bond consists of 2 shared pairs of electrons between the carbons.
Methanol (CH₃OH)
Methanol has 3 C-H single bonds, 1 C-O single bond, and 1 O-H single bond. Oxygen has 2 lone pairs. Carbon has 4 bonds total (no lone pairs), and oxygen has 2 bonds + 2 lone pairs.
Explaining Properties of Simple Molecules
When you heat a simple covalent substance, you are not breaking the covalent bonds inside each molecule. You are overcoming the weak intermolecular forces (forces between separate molecules). Think of it like this: each molecule is a tightly bonded group of atoms, but the forces holding different molecules near each other are very weak.
Since the intermolecular forces are weak:
- Low melting and boiling points — only a small amount of energy is needed to overcome the weak intermolecular forces and separate the molecules.
- Many are gases or liquids at room temperature — because the forces between molecules are so easily overcome.
- As molecules get bigger (more electrons, more mass), the intermolecular forces get slightly stronger, so boiling points increase. For example, methane (CH₄) is a gas at room temperature, but octane (C₈H₁₈) is a liquid.
Simple covalent substances do not conduct electricity because:
- There are no ions — the atoms are held together by shared electrons, not by ionic bonds.
- There are no free (delocalised) electrons — all electrons are locked in covalent bonds or lone pairs.
- Without mobile charged particles, no electric current can flow.
🧠 Memory Trick
Covalent = Co-operating to Achieve a goal. Non-metal atoms cooperate by sharing electrons, unlike metals in ionic bonding which just hand them over. Think of two students sharing a textbook in class — both can use it at the same time.
🧠 Memory Trick
To remember how many bonds each element forms: H=1, O=2, N=3, C=4. Think HONC-1234. Hydrogen always makes 1 bond. Oxygen makes 2. Nitrogen makes 3. Carbon makes 4. This works because each element needs to fill its outer shell: H needs 1 more, O needs 2 more, N needs 3 more, C needs 4 more.
Worked Example
Draw a dot-and-cross diagram for a molecule of water (H₂O). State the type of bonding present.
Step 1: Count outer electrons
Oxygen has 6 outer electrons. Each hydrogen has 1 outer electron.
Step 2: Work out the bonds
Oxygen needs 2 more electrons to have 8 (full shell). Each H needs 1 more to have 2 (full shell like He). So O forms 2 covalent bonds, one with each H atom.
Step 3: Draw the diagram
Place O in the centre. Show its 6 outer electrons using crosses. Place 2 H atoms either side. Show each H electron using a dot. The shared pairs (one dot + one cross between O and each H) are the covalent bonds. O also has 2 lone pairs (4 non-bonding electrons).
Answer: H₂O has covalent bonding (shared pairs of electrons between non-metal atoms). O has 2 bonding pairs and 2 lone pairs. Each H has 1 bonding pair.
Worked Example
SUPPLEMENT Explain why water (H₂O) has a much lower boiling point than sodium chloride (NaCl), even though both contain oxygen/chlorine from the same period.
Step 1: Identify the structures
Water is a simple covalent molecule. NaCl is a giant ionic lattice.
Step 2: Compare the forces that must be broken on boiling
To boil water, you must overcome the weak intermolecular forces between water molecules (NOT the O-H covalent bonds). To boil NaCl, you must overcome the strong electrostatic forces between many Na⁺ and Cl⁻ ions throughout the giant lattice.
Step 3: Compare energy needed
The weak intermolecular forces in water require much less energy to overcome than the strong ionic bonds in NaCl. So water boils at 100°C but NaCl boils at 1413°C.
Answer: Water has weak intermolecular forces between its molecules, requiring little energy to overcome. NaCl has strong electrostatic attractions between ions in a giant lattice, requiring much more energy. So water has a much lower boiling point.
Apply It: Real-World Chemistry
Covalent molecules are everywhere in your daily life — from the water you drink to the oxygen you breathe to the petrol that powers cars. Let us see covalent bonding in action.
1
When you swim in a pool, you are surrounded by trillions of water molecules. Water is a liquid at room temperature and boils at 100°C. But hydrogen sulfide (H₂S), which has a similar structure to water and is actually heavier, is a gas at room temperature and boils at -60°C.
Both H₂O and H₂S are simple covalent molecules. Why does water have a much higher boiling point than H₂S?
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The Chemistry
Both H₂O and H₂S are simple covalent molecules with weak intermolecular forces between the molecules. However, water has unusually strong intermolecular forces called hydrogen bonds, which form because oxygen is very electronegative and the O-H bond is very polar. H₂S cannot form these hydrogen bonds because sulfur is much less electronegative. The stronger intermolecular forces in water mean more energy is needed to separate the molecules, so water has a higher boiling point. At IGCSE level, the key point is: the intermolecular forces between water molecules are stronger than those between H₂S molecules.
The Big Idea
Boiling points of simple covalent substances depend on the strength of forces BETWEEN molecules, not the strength of covalent bonds WITHIN molecules. Stronger intermolecular forces = higher boiling point. This is why water is a liquid at room temperature while many similar-sized molecules are gases.
2
Nitrogen gas (N₂) makes up 78% of the air. Despite being all around us, nitrogen is incredibly unreactive — it does not burn, does not support combustion, and does not react with most substances at normal temperatures. It is used to fill crisp packets in India and the UK to keep them fresh.
Why is nitrogen gas so unreactive compared to oxygen gas?
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The Chemistry
Nitrogen gas (N₂) contains a triple covalent bond (N≡N) — three shared pairs of electrons between the two nitrogen atoms. This triple bond is extremely strong and requires a huge amount of energy to break (945 kJ/mol). Oxygen (O₂) has only a double bond (O=O), which is weaker (498 kJ/mol). Because the triple bond in N₂ is so hard to break, nitrogen does not easily react with other substances. It takes extreme conditions — like the very high temperatures inside a car engine or a lightning strike — to break the N≡N bond.
The Big Idea
More shared pairs = stronger bond. Single bond < double bond < triple bond in strength. The triple bond in N₂ is one of the strongest covalent bonds known, which is why nitrogen is used as an inert gas to prevent oxidation in food packaging and industrial processes.
3
In Indian homes, LPG (liquefied petroleum gas) is widely used for cooking. LPG is mainly a mixture of propane (C₃H₈) and butane (C₄H₁₀). At room temperature and normal pressure, both of these are gases. But inside the blue LPG cylinder, they are stored as liquids under pressure.
Why are propane and butane gases at room temperature, and how does pressure convert them to liquids in the cylinder?
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The Chemistry
Propane and butane are simple covalent molecules. The intermolecular forces between these molecules are weak, so at room temperature the molecules have enough kinetic energy to overcome these forces and exist as gases. When you compress the gas inside the cylinder, the molecules are forced closer together, and the intermolecular forces become strong enough to hold them as a liquid. When you open the valve, the pressure drops and the liquid rapidly boils back into a gas — this is why you can feel the cylinder getting cold when you use gas (the liquid absorbs heat as it evaporates).
The Big Idea
Simple covalent molecules have low boiling points because the intermolecular forces are weak. By increasing pressure, you can force these molecules close enough together to become a liquid. This is the principle behind LPG storage, refrigeration, and even aerosol cans.
4
Pure water does not conduct electricity. You can safely put your hand in a beaker of distilled water with electrodes and no current will flow. However, if you add a pinch of salt, the water immediately begins to conduct and can give you a shock.
Why does pure water not conduct electricity, even though it is a liquid?
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The Chemistry
Pure water (H₂O) is a simple covalent molecule. It contains no ions and no free electrons — all the electrons are locked in covalent O-H bonds or in lone pairs on the oxygen. For a substance to conduct electricity, it needs mobile charged particles (either ions or free electrons). Pure water has neither, so it does not conduct. When you add salt (NaCl), the salt dissolves into Na⁺ and Cl⁻ ions, which are free to move through the water and carry charge — so the solution now conducts.
The Big Idea
Being a liquid is not enough to conduct electricity. The substance needs mobile charged particles. Molten ionic compounds and ionic solutions conduct because they have free ions. Simple covalent liquids (water, ethanol, oil) do not conduct because they have no ions or free electrons. This is a key distinction the examiner loves to test.
5
Anaesthetics used in hospitals contain molecules like sevoflurane (C₄H₃F₇O), which is a liquid that evaporates easily. When a patient breathes in the vapour, they lose consciousness within seconds. After surgery, the anaesthetic wears off quickly as the molecules leave the body through the lungs.
Why does sevoflurane evaporate so easily, and why does its effect wear off quickly?
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The Chemistry
Sevoflurane is a simple covalent molecule. The intermolecular forces between sevoflurane molecules are relatively weak (despite the molecule being fairly large), partly because the C-F bonds are not very polar in a way that promotes strong intermolecular interactions. This means the liquid evaporates easily — the molecules readily escape from the liquid surface into the gas phase. Because the molecules are small enough and have weak intermolecular forces, they can also easily pass through the thin walls of the alveoli in the lungs (both going in and coming back out).
The Big Idea
The low boiling point and easy evaporation of simple covalent molecules is not just a textbook fact — it is the reason modern surgery is possible. Anaesthetics must be volatile (evaporate easily) so they can be inhaled as gases, and they must wear off quickly so patients recover. The weak intermolecular forces in these molecules make all of this possible.
Test Yourself: Simple Molecules and Covalent Bonds
20 questions covering covalent bonding, dot-and-cross diagrams, and properties
Your Score:
0 / 20
QUESTION 1 OF 20
What is a covalent bond?
A covalent bond is a shared pair of electrons between two atoms. Option A describes ionic bonding, C describes metallic bonding, and D is the definition of an ionic bond.
QUESTION 2 OF 20
How many covalent bonds does a carbon atom form?
Carbon has 4 outer electrons and needs 8 for a full shell. It must share 4 more electrons, so it forms 4 covalent bonds. Remember HONC-1234: Hydrogen=1, Oxygen=2, Nitrogen=3, Carbon=4.
QUESTION 3 OF 20
In a dot-and-cross diagram of H₂O, how many lone pairs does the oxygen atom have?
Oxygen has 6 outer electrons. 2 are used in bonding pairs (one with each H), leaving 4 non-bonding electrons = 2 lone pairs. Total: 2 bonding pairs + 2 lone pairs = 8 outer electrons (full shell).
QUESTION 4 OF 20
Why do simple covalent substances have low boiling points?
Covalent bonds within molecules are STRONG. It is the intermolecular forces BETWEEN separate molecules that are weak. When a substance boils, it is these weak intermolecular forces that are overcome, not the covalent bonds. This is one of the most important distinctions in IGCSE Chemistry.
QUESTION 5 OF 20
Which of these molecules contains a double covalent bond?
O₂ contains a double bond (O=O) — two shared pairs of electrons. H₂ has a single bond, H₂O has two single O-H bonds, and HCl has a single H-Cl bond.
QUESTION 6 OF 20
Covalent bonds form between:
Covalent bonds form between two non-metal atoms, which share electrons. Metal + non-metal = ionic bonding (electron transfer). Metal + metal = metallic bonding. Noble gases generally do not form bonds at all.
QUESTION 7 OF 20
In a molecule of methane (CH₄), how many electrons are in the outer shell of the carbon atom?
Carbon contributes 4 electrons and shares 1 from each of 4 hydrogen atoms = 4 + 4 = 8 electrons in carbon's outer shell. Each shared pair consists of 1 carbon electron + 1 hydrogen electron, and carbon counts all 8 shared electrons as part of its outer shell.
QUESTION 8 OF 20
Why does pure water not conduct electricity?
Pure water is a simple covalent molecule with no ions and no free electrons. For electricity to flow, you need mobile charged particles. Water molecules are neutral — they have no charge. Adding an ionic compound (like NaCl) provides the free ions needed for conduction.
QUESTION 9 OF 20
How many bonding pairs and lone pairs does nitrogen have in a molecule of ammonia (NH₃)?
Nitrogen has 5 outer electrons. It forms 3 bonds (one with each H), using 3 electrons for bonding. The remaining 2 electrons form 1 lone pair. Total: 3 bonding pairs + 1 lone pair = 8 outer electrons.
QUESTION 10 OF 20
Which substance has the highest boiling point?
NaCl is a giant ionic lattice with strong electrostatic forces throughout, giving it a very high boiling point (1413°C). The other three are all simple covalent molecules with weak intermolecular forces and low boiling points. Among covalent molecules, larger ones (more electrons) tend to have higher boiling points.
QUESTION 11 OF 20
SUPPLEMENT What type of bond is found in a molecule of nitrogen (N₂)?
Nitrogen has 5 outer electrons and needs 8. Each N atom shares 3 electrons with the other, forming 3 shared pairs = a triple bond (N≡N). This is one of the strongest bonds in chemistry and explains why N₂ is so unreactive.
QUESTION 12 OF 20
SUPPLEMENT In a molecule of CO₂, carbon forms:
Carbon has 4 outer electrons. In CO₂, it forms 2 double bonds — one with each oxygen atom. Each double bond involves 2 shared pairs. Carbon shares a total of 8 electrons (4 bonds worth), filling its outer shell. Each oxygen has 2 bonding pairs + 2 lone pairs = 8 outer electrons.
QUESTION 13 OF 20
What happens to the covalent bonds inside water molecules when water boils?
When water boils, the covalent O-H bonds within each molecule remain completely intact. The molecules themselves are unchanged — still H₂O. What is overcome are the weak intermolecular forces between the water molecules. The water molecules simply separate from each other and enter the gas phase. If the covalent bonds broke, you would get hydrogen and oxygen gas (electrolysis, not boiling!).
QUESTION 14 OF 20
Which of these dot-and-cross diagrams is correct for HCl?
HCl has one shared pair of electrons (a single covalent bond). H contributes 1 electron and Cl contributes 1, giving H a total of 2 (full like He) and Cl a total of 8 outer electrons (full shell). Option C describes ionic bonding (which does not happen because both are non-metals in this molecule). H only needs 1 bond.
QUESTION 15 OF 20
Which of these is NOT a property of simple covalent substances?
Simple covalent substances do NOT conduct electricity in any state — solid, liquid, or dissolved. They have no ions or free electrons. Conducting when dissolved describes ionic compounds, not covalent ones. When sugar (a covalent substance) dissolves in water, the solution does not conduct.
QUESTION 16 OF 20
SUPPLEMENT In the molecule C₂H₄ (ethene), how many shared pairs of electrons are between the two carbon atoms?
Ethene has a C=C double bond, meaning 2 shared pairs of electrons between the two carbon atoms. Each carbon also forms 2 single bonds with hydrogen atoms, giving each carbon 4 bonds total (2 in the double bond + 2 with H atoms).
QUESTION 17 OF 20
SUPPLEMENT Which correctly explains why chlorine (Cl₂) has a low boiling point?
Cl₂ is a simple covalent molecule. The covalent bond WITHIN each Cl₂ molecule is strong, but the intermolecular forces BETWEEN separate Cl₂ molecules are weak. It is these weak forces that are overcome when chlorine boils (-34°C). The covalent bonds stay intact.
QUESTION 18 OF 20
In a molecule of Cl₂, the shared pair of electrons is attracted to:
The shared pair of electrons sits between the two nuclei and is attracted to BOTH positively charged nuclei. This mutual attraction is what holds the atoms together in a covalent bond. The electrons do not float freely — they are strongly attracted to both nuclei.
QUESTION 19 OF 20
SUPPLEMENT As the size of simple covalent molecules increases, what happens to their boiling points?
Larger molecules have more electrons and more surface area, which means stronger intermolecular forces between them. Stronger intermolecular forces require more energy to overcome, so the boiling point increases. This is why methane (CH₄) is a gas but octane (C₈H₁₈) is a liquid at room temperature.
QUESTION 20 OF 20
A substance has a very low melting point, does not conduct electricity in any state, and is insoluble in water. It is most likely:
Low melting point = weak intermolecular forces (simple covalent). Does not conduct = no ions or free electrons (not ionic or metallic). Insoluble in water = non-polar covalent (many covalent substances are insoluble). Ionic compounds have high melting points and conduct when molten. Metals conduct. Giant covalent substances have very high melting points.
2.6 Giant Covalent Structures
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What Are Giant Covalent Structures?
In simple covalent substances (like H₂O and CH₄), atoms are bonded in small individual molecules. But some covalent substances are very different — they form giant covalent structures (also called macromolecular structures), where millions of atoms are all bonded together by strong covalent bonds in a continuous network. There are no separate molecules — just one enormous structure.
The three giant covalent substances you need to know are diamond, graphite, and (for the Supplement) silicon(IV) oxide (SiO₂). Diamond and graphite are both made entirely of carbon atoms, but the way the atoms are arranged is completely different — giving them very different properties.
Diamond
In diamond, every carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. This creates a rigid, 3D network of strong C-C covalent bonds extending in all directions.
In diamond, each carbon atom is bonded to 4 other carbon atoms by strong covalent bonds. All 4 outer electrons of each carbon are used in bonding — there are no free electrons and no weak forces anywhere in the structure.
Properties of Diamond
- Very hard — each carbon is bonded to 4 others in all directions, forming an extremely rigid structure. Diamond is the hardest natural substance known.
- Very high melting point (about 3550°C) — many strong covalent bonds must be broken to melt diamond, requiring an enormous amount of energy.
- Does not conduct electricity — all 4 outer electrons on each carbon are locked in covalent bonds. There are no free (delocalised) electrons or ions to carry charge.
- Transparent and colourless — pure diamond allows light to pass through.
Uses of Diamond
- Cutting tools and drill bits — because diamond is the hardest known substance, it can cut through almost anything. Industrial diamonds are used in drill bits for oil exploration and in saw blades for cutting stone and concrete.
- Jewellery — diamonds sparkle brilliantly because of the way they refract light. India is one of the world's largest centres for diamond cutting and polishing, particularly in Surat, Gujarat.
- Glass cutting — a diamond-tipped tool can score and cut glass cleanly.
Graphite
In graphite, each carbon atom is covalently bonded to three other carbon atoms, forming flat layers (sheets) of hexagonal rings. Between the layers, there are only weak intermolecular forces.
In graphite, each carbon bonds to 3 others in flat hexagonal layers. The 4th electron from each carbon is delocalised (free to move) between the layers, allowing graphite to conduct electricity. The weak forces between layers allow them to slide over each other.
Properties of Graphite
- Soft and slippery — the layers can slide over each other because only weak intermolecular forces hold the layers together. This is why graphite is used as a lubricant and leaves a mark on paper (pencils).
- Very high melting point (about 3730°C) — even though the layers slide easily, the covalent bonds within each layer are very strong. You must break these C-C bonds to melt graphite, which requires enormous energy.
- Conducts electricity — each carbon is bonded to only 3 others (using 3 of its 4 outer electrons). The fourth electron from each carbon is delocalised — free to move along the layers. These mobile electrons can carry an electric current. Graphite is the only non-metal element that conducts electricity well.
- Black and opaque — graphite absorbs light.
Uses of Graphite
- Pencils — the soft, slippery layers slide off and leave a mark on paper.
- Lubricant — the sliding layers reduce friction between moving parts in machinery. Graphite lubricant is used in locks, hinges, and high-temperature machinery where oil would burn off.
- Electrodes — graphite conducts electricity and has a very high melting point, so it is used as electrodes in electrolysis and in electric arc furnaces used in steel production.
Comparison: Diamond vs Graphite
| Property | Diamond | Graphite |
|---|---|---|
| Bonds per carbon | 4 (tetrahedral) | 3 (flat layers) + 1 delocalised |
| Structure | 3D rigid network | Flat layers with weak forces between |
| Hardness | Very hard (hardest natural substance) | Soft and slippery |
| Melting point | Very high (~3550°C) | Very high (~3730°C) |
| Electrical conductivity | Does not conduct | Conducts (delocalised electrons) |
| Appearance | Transparent, colourless, sparkly | Black, opaque, shiny |
| Uses | Cutting tools, jewellery | Pencils, lubricant, electrodes |
🧠 Memory Trick
Diamond = 4 bonds = HARD. Think of 4 as "all locked up" — every electron is bonded, nothing is free, everything is rigid. Graphite = 3 bonds + 1 free = SOFT + CONDUCTS. Think of 3 as "one hand free" — the fourth electron is not bonded and can move around, letting graphite conduct and making the layers slide.
🧠 Memory Trick
"PENCIL = PEEL off layers." When you write with a pencil, you are literally peeling off layers of graphite onto the paper. The layers slide because the forces between them are weak. If pencils were made of diamond, you could not write with them — diamond is far too hard!
SUPPLEMENT
Silicon(IV) Oxide (SiO₂)
Silicon(IV) oxide (also called silicon dioxide or silica) has a giant covalent structure similar to diamond. Each silicon atom is bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms. This creates a continuous 3D network of strong Si-O covalent bonds.
Properties of SiO₂:
- Very high melting point (about 1710°C) — many strong covalent bonds must be broken.
- Very hard — rigid 3D network, similar to diamond.
- Does not conduct electricity — no free electrons or ions.
SiO₂ is the main component of sand and quartz. It is used in making glass, optical fibres, and is abundant in the Earth's crust.
Explaining Properties in Terms of Structure
The key to explaining properties of giant covalent structures is always the same: strong covalent bonds throughout the entire structure.
- High melting points: To melt any of these substances, you need to break many strong covalent bonds throughout the giant structure. This requires a very large amount of energy. There are no weak intermolecular forces here — it is all strong bonds.
- Hardness (diamond, SiO₂): The rigid 3D network of bonds in all directions makes the structure extremely resistant to being deformed.
- Softness (graphite): The layers within graphite are strong, but the weak forces between layers allow them to slide. The structure is only soft in one direction — perpendicular to the layers.
- Electrical conductivity: Diamond and SiO₂ do not conduct because all outer electrons are locked in covalent bonds. Graphite conducts because each carbon only uses 3 of its 4 outer electrons for bonding — the fourth electron is delocalised (free to move along the layers and carry charge).
⚠ Exam Tip
Both diamond and graphite have very high melting points. Students often think graphite has a low melting point because it is soft, but that is wrong! Graphite's softness comes from layers sliding, NOT from weak bonds. The covalent bonds within the layers are extremely strong. You need to break those bonds to melt it.
⚠ Exam Tip
When explaining why graphite conducts electricity, you MUST say "delocalised electrons". Do not say "free ions" (that is ionic compounds) or "metallic bonds" (that is metals). Say: "Each carbon atom is bonded to only 3 others. The fourth outer electron is delocalised and free to move along the layers, carrying charge."
Worked Example
Explain why diamond is very hard but graphite is soft and slippery, even though both are made entirely of carbon atoms.
Step 1: Describe diamond's structure
In diamond, each carbon atom is covalently bonded to 4 other carbon atoms in a rigid 3D tetrahedral arrangement. Strong covalent bonds extend in all directions throughout the structure, making it extremely difficult to deform.
Step 2: Describe graphite's structure
In graphite, each carbon atom is covalently bonded to only 3 other carbon atoms, forming flat layers of hexagonal rings. Between these layers, there are only weak intermolecular forces.
Step 3: Explain the difference
Diamond is hard because bonds extend rigidly in all 3 dimensions — there is no weak point. Graphite is soft because the weak forces between layers can be easily overcome, allowing the layers to slide over each other.
Answer: Diamond has each C bonded to 4 others in a rigid 3D network with no weak points, making it very hard. Graphite has C bonded to 3 others in flat layers with weak forces between layers, allowing them to slide, making it soft.
Worked Example
SUPPLEMENT Explain why graphite can conduct electricity but diamond cannot.
Step 1: Diamond
In diamond, every carbon atom uses all 4 of its outer electrons to form covalent bonds with 4 other carbon atoms. There are no free (delocalised) electrons available to carry charge, so diamond does not conduct electricity.
Step 2: Graphite
In graphite, each carbon atom uses only 3 of its 4 outer electrons for covalent bonding. The fourth outer electron on each carbon atom is delocalised — it is free to move along the layers. These delocalised electrons can carry an electric charge, so graphite conducts electricity.
Answer: In diamond, all outer electrons are used in bonding so none are free to move. In graphite, each C has 1 delocalised electron that is free to move along the layers and carry charge.
Apply It: Real-World Chemistry
Diamond and graphite are both pure carbon, yet one is the hardest natural substance and the other is soft enough to write with. Let us see how their structures determine their incredible real-world uses.
1
India is the world's largest centre for diamond cutting and polishing, with around 90% of the world's diamonds being processed in Surat, Gujarat. Diamond-tipped drill bits are used in oil exploration across Rajasthan and offshore near Mumbai. Industrial diamonds cost a fraction of gem-quality ones.
Why are diamonds used for cutting and drilling rather than any other material?
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The Chemistry
Diamond is the hardest natural substance known. Its extreme hardness comes from its structure: every carbon atom is bonded to 4 others by strong covalent bonds in a rigid 3D tetrahedral arrangement. There are no weak points in any direction. This means diamond can scratch or cut virtually any other material. In oil drilling, diamond-tipped drill bits can cut through extremely hard rock formations deep underground. In the gem industry, only diamond can cut and polish another diamond.
The Big Idea
The extreme hardness of diamond is a direct consequence of its structure — strong covalent bonds in all directions with no weak points. Structure determines properties, and properties determine uses. A material's usefulness comes from the arrangement of its atoms.
2
The pencil you use to write in exams contains graphite mixed with clay (not lead, despite being called "pencil lead"). A standard HB pencil has about 68% graphite. When you write, a thin grey line appears on the paper. The line is actually microscopic layers of graphite that have been rubbed off.
Why does graphite leave a mark on paper when you write, and why does diamond not?
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The Chemistry
In graphite, the carbon atoms are arranged in flat layers of hexagonal rings. Within each layer, strong covalent bonds hold the atoms together. But between the layers, there are only weak intermolecular forces. When the pencil presses against the rough surface of paper, the friction is enough to overcome these weak forces, causing layers to slide off and stick to the paper. In diamond, every carbon is bonded to 4 others in all directions — there are no layers that can slide. Instead, diamond just scratches the surface without leaving a mark of its own material.
The Big Idea
Graphite and diamond are both 100% carbon, yet one leaves marks and the other does not. The only difference is how the atoms are arranged. This is one of the most striking examples in chemistry of how structure determines properties.
3
In aluminium smelting plants (like those run by HINDALCO in India and RUSAL in Russia), aluminium is extracted from aluminium oxide (Al₂O₃) by electrolysis. The electrodes used in this process are made of graphite. These electrodes must conduct electricity and withstand temperatures above 900°C.
Why is graphite chosen as the electrode material for aluminium electrolysis?
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The Chemistry
Graphite has two key properties that make it ideal for electrodes. First, it conducts electricity because each carbon atom has one delocalised electron that is free to move along the layers and carry charge. Second, it has a very high melting point (about 3730°C), so it can withstand the extreme temperatures inside the electrolysis cell (typically around 950°C). Most metals would melt or react at these temperatures. Graphite is also relatively cheap and chemically inert (unreactive).
The Big Idea
Graphite is one of the very few materials that combines electrical conductivity with an extremely high melting point. This unique combination of properties — which comes directly from its layered structure with delocalised electrons and strong covalent bonds within layers — makes it irreplaceable in high-temperature electrolysis.
4
Glass is made primarily from silicon dioxide (SiO₂), which is the main component of sand. When sand is heated to about 1700°C, it melts and can be shaped into windows, bottles, lenses, and optical fibres. The optical fibres that carry internet data across oceans are made from ultra-pure SiO₂.
Why does sand (SiO₂) need such a high temperature to melt, and why is it suitable for making glass?
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The Chemistry
Silicon dioxide has a giant covalent structure similar to diamond. Each silicon atom is covalently bonded to 4 oxygen atoms, and each oxygen atom is bonded to 2 silicon atoms, creating a continuous 3D network of strong Si-O covalent bonds. To melt SiO₂, you must break many of these strong bonds, which requires an enormous amount of energy — hence the very high melting point of about 1710°C. Once melted and then cooled slowly, SiO₂ forms glass, which is actually an amorphous solid (it does not crystallise back into a regular structure). Glass is transparent because the structure allows visible light to pass through.
The Big Idea
Giant covalent structures like SiO₂ have high melting points because of strong covalent bonds throughout the structure, not because of any special property of silicon or oxygen individually. The same principle applies to diamond — strong bonds throughout = high melting point.
5
Car mechanics and cyclists use graphite-based lubricants on chains, gears, and locks. Unlike oil-based lubricants, graphite lubricant works well at very high temperatures and does not attract dust. A thin coating of graphite powder on a bicycle chain can last for weeks without reapplication.
Why does graphite work as a lubricant, and why is it better than oil in some situations?
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The Chemistry
Graphite works as a lubricant because of its layered structure. The carbon atoms within each layer are held together by strong covalent bonds, but the layers are held to each other by only weak intermolecular forces. When graphite is placed between two moving surfaces, the layers slide over each other with very little friction, reducing wear. Unlike oil, graphite does not evaporate at high temperatures (its melting point is 3730°C), does not become viscous in cold temperatures, and does not decompose or catch fire. It also stays dry, so it does not attract dust and grit.
The Big Idea
The ability of graphite layers to slide is a direct consequence of the weak intermolecular forces between the layers. This is the same reason pencils work — layers sliding onto paper. The combination of strong bonds within layers (high temperature resistance) and weak forces between layers (sliding ability) makes graphite unique.
Test Yourself: Giant Covalent Structures
20 questions covering diamond, graphite, silicon dioxide, and their properties
Your Score:
0 / 20
QUESTION 1 OF 20
In diamond, each carbon atom is bonded to how many other carbon atoms?
In diamond, each carbon atom forms 4 covalent bonds with 4 other carbon atoms in a tetrahedral arrangement. All 4 outer electrons are used in bonding.
QUESTION 2 OF 20
Why does diamond not conduct electricity?
In diamond, all 4 outer electrons on each carbon are used in covalent bonds. There are no free (delocalised) electrons and no ions, so there are no mobile charged particles to carry an electric current. Note: diamond does NOT have layers — that is graphite.
QUESTION 3 OF 20
In graphite, each carbon atom is bonded to how many other carbon atoms?
In graphite, each carbon atom forms covalent bonds with only 3 other carbon atoms in flat hexagonal layers. The 4th outer electron is delocalised (free to move between the layers).
QUESTION 4 OF 20
Why is graphite soft and slippery?
The covalent bonds within each layer of graphite are very strong. But the forces between layers are weak intermolecular forces. These weak forces are easily overcome, allowing the layers to slide over each other, making graphite soft and slippery.
QUESTION 5 OF 20
Why does graphite conduct electricity?
Each carbon in graphite bonds to only 3 others, using 3 of its 4 outer electrons. The 4th electron is delocalised — free to move along the layers. These mobile electrons can carry an electric charge. Graphite is NOT a metal — it is a non-metal that conducts because of delocalised electrons.
QUESTION 6 OF 20
Which of the following is a use of diamond?
Diamond is used for cutting tools and drill bits because it is the hardest natural substance. Pencils, electrodes, and lubricants all use graphite, not diamond.
QUESTION 7 OF 20
Both diamond and graphite have very high melting points. This is because:
Both diamond and graphite are giant covalent structures with many strong C-C covalent bonds throughout their structures. To melt either substance, many of these strong bonds must be broken, requiring a very large amount of energy. Even graphite, despite being soft, has very strong bonds within its layers.
QUESTION 8 OF 20
Which use of graphite relies on its ability to conduct electricity?
Electrodes need to conduct electricity to allow current to flow into the electrolyte. Graphite is used as electrodes because it conducts (delocalised electrons) AND has a very high melting point. Pencils use graphite's softness (sliding layers), lubricants use its slipperiness, and jewellery uses diamond.
QUESTION 9 OF 20
Diamond and graphite are both made entirely of carbon. They are examples of:
Diamond and graphite are allotropes of carbon — different physical forms of the same element. Isotopes have different numbers of neutrons, isomers are different arrangements of the same molecular formula in compounds, and compounds contain more than one element.
QUESTION 10 OF 20
A student says: "Graphite has a low melting point because it is soft." Is the student correct?
The student is wrong. Graphite actually has a VERY high melting point (~3730°C) because the covalent bonds within each layer are extremely strong. Its softness is due to weak forces between layers allowing them to slide, NOT due to weak bonds. Softness and melting point are unrelated in this case.
QUESTION 11 OF 20
SUPPLEMENT Silicon dioxide (SiO₂) has a structure similar to:
SiO₂ has a giant covalent structure similar to diamond, with each Si bonded to 4 O atoms and each O bonded to 2 Si atoms in a 3D network. Like diamond, it is very hard, has a very high melting point, and does not conduct electricity.
QUESTION 12 OF 20
SUPPLEMENT In silicon dioxide, each silicon atom is bonded to:
In SiO₂, each silicon atom is bonded to 4 oxygen atoms tetrahedrally, and each oxygen atom is bonded to 2 silicon atoms. The formula SiO₂ reflects the ratio (1 Si : 2 O) but each Si is actually bonded to 4 O atoms in the giant structure.
QUESTION 13 OF 20
Which property of graphite makes it suitable for use in pencils?
Pencils work because graphite layers slide off onto the paper due to the weak forces between layers. The layers are thin enough to appear as a grey/black mark. Electrical conductivity and high melting point are useful for other applications but not for pencils.
QUESTION 14 OF 20
SUPPLEMENT Which statement correctly explains why diamond has a very high melting point?
Diamond is a giant covalent structure — there are no molecules, so "intermolecular forces" is wrong. There are no ions (not ionic) and no metallic bonds. The high melting point is because many strong covalent bonds extend throughout the entire 3D structure, and a lot of energy is needed to break them.
QUESTION 15 OF 20
Which of these substances does NOT have a giant covalent structure?
Carbon dioxide (CO₂) is a simple covalent molecule — it exists as individual O=C=O molecules with weak intermolecular forces between them. It is a gas at room temperature. Diamond, graphite, and silicon dioxide are all giant covalent structures with high melting points.
QUESTION 16 OF 20
Which of the following is NOT a property of giant covalent structures?
Giant covalent structures generally do NOT conduct electricity at all (except graphite, which conducts due to delocalised electrons). They do not contain ions, so they cannot conduct when dissolved. Conducting when dissolved is a property of ionic compounds.
QUESTION 17 OF 20
Why is graphite used as a lubricant rather than diamond?
Graphite is a lubricant because its layers can slide over each other due to weak intermolecular forces between them. Diamond has a rigid 3D structure with no layers and no sliding ability. While cost is also a factor, the scientific reason is the layered structure.
QUESTION 18 OF 20
SUPPLEMENT Which of these correctly compares diamond and graphite?
In diamond, all 4 outer electrons per carbon are in covalent bonds. In graphite, only 3 electrons per carbon are in bonds; the 4th is delocalised. Option A is reversed (graphite conducts, not diamond). Both have similar melting points. Both are pure carbon — graphite does not contain other elements.
QUESTION 19 OF 20
A substance is very hard, has a very high melting point, and does not conduct electricity. It could be:
Very hard + very high mp + does not conduct = diamond (or SiO₂). Graphite conducts and is soft. NaCl conducts when molten and is not especially hard. Iron conducts electricity. Diamond matches all three criteria perfectly.
QUESTION 20 OF 20
SUPPLEMENT Silicon dioxide does not conduct electricity because:
SiO₂ is a giant covalent structure (NOT ionic). All outer electrons on both Si and O are involved in covalent bonds. There are no free electrons or ions to carry charge. SiO₂ does not have layers (unlike graphite) and is a solid at room temperature (mp ~1710°C).
SUPPLEMENT 2.7 Metallic Bonding
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This entire section is Supplement (Extended) content. It will only be tested on the Extended paper. If you are sitting the Core paper, you can skip this section — but understanding metallic bonding will deepen your understanding of how different types of bonding work.
What Is Metallic Bonding?
You already know about two types of bonding:
- Ionic bonding — electrons are transferred from metal to non-metal, forming ions held by electrostatic attraction.
- Covalent bonding — electrons are shared between non-metal atoms.
But what holds metal atoms together in a piece of iron, copper, or gold? Metal atoms do not transfer electrons to each other (they are all metals), and they do not share pairs of electrons like non-metals do. Instead, metals have their own special type of bonding: metallic bonding.
In metallic bonding:
- Each metal atom loses its outer shell electrons, becoming a positive metal ion (cation).
- The lost electrons do not go to any particular atom — instead, they form a "sea" of delocalised electrons that are shared among ALL the metal ions.
- The positive metal ions are arranged in a regular lattice (a repeating pattern).
- The positive ions are held together by the strong electrostatic attraction between the positive ions and the sea of negative delocalised electrons.
This attraction between positive ions and delocalised electrons is what we call a metallic bond.
In a metal, positive ions (M⁺) are arranged in a regular lattice. The outer shell electrons are delocalised — they form a "sea" of electrons that moves freely throughout the structure. The strong electrostatic attraction between the positive ions and the sea of negative electrons is the metallic bond.
Key Points About Metallic Bonding
- The metal atoms lose their outer electrons and become positive ions.
- The electrons do not belong to any particular ion — they are delocalised (free to move throughout the whole structure).
- The bonding is the electrostatic attraction between the positive metal ions and the sea of delocalised electrons.
- This is NOT the same as covalent bonding (no shared pairs) and NOT the same as ionic bonding (no transfer to a specific atom).
Properties of Metals Explained by Metallic Bonding
1. Metals Conduct Electricity
When a voltage is applied across a piece of metal, the delocalised electrons can flow through the structure towards the positive terminal. Because these electrons are free to move, metals are excellent electrical conductors. This is why copper is used for electrical wires.
When a voltage is applied, the delocalised electrons drift through the metal towards the positive terminal. The positive ions stay in their fixed positions in the lattice. This flow of electrons is an electric current.
2. Metals Conduct Heat
The delocalised electrons also transfer kinetic energy quickly through the structure. When one end of a metal bar is heated, the delocalised electrons near the hot end gain energy and move faster, colliding with electrons further along and transferring energy rapidly. This is why metals feel cold to the touch — they conduct heat away from your warm hand very efficiently.
3. Metals Are Malleable (Can Be Hammered into Shape)
When a force is applied to a metal, the layers of ions can slide over each other into new positions. Crucially, the sea of delocalised electrons is not directional — it can rearrange around the ions in their new positions immediately. So the metallic bond is maintained even after the ions move. This is why metals can be bent, hammered, and pressed into sheets without breaking.
When a force is applied to a metal, the layers of positive ions slide over each other. But the sea of delocalised electrons flows around the ions in their new positions, maintaining the metallic bond. The metal changes shape without breaking. Compare this with an ionic lattice, where sliding brings like charges together, causing the lattice to shatter.
4. Metals Are Ductile (Can Be Drawn into Wires)
For the same reason metals are malleable, they are also ductile — they can be pulled and stretched into thin wires without breaking. The layers of ions slide past each other, and the delocalised electrons rearrange to maintain the bond. This is why copper can be drawn into thin electrical wires.
5. Metals Have High Melting Points (Usually)
The metallic bond (electrostatic attraction between positive ions and delocalised electrons) is strong. To melt a metal, you need to provide enough energy to overcome these strong attractions and allow the ions to move freely. Most metals therefore have high melting points. The more outer electrons each atom contributes to the "sea", the stronger the metallic bond and the higher the melting point.
Comparing All Three Types of Bonding
| Feature | Ionic | Covalent (Simple) | Metallic |
|---|---|---|---|
| What happens to electrons? | Transferred from metal to non-metal | Shared between non-metal atoms | Delocalised into a "sea" |
| Particles present | Positive and negative ions | Molecules (atoms sharing electrons) | Positive ions + delocalised electrons |
| Force holding them | Electrostatic attraction between ions | Shared pair attraction to both nuclei | Attraction between ions and electron sea |
| Melting point | High | Low (simple); Very high (giant) | Usually high |
| Conducts solid? | No (ions fixed) | No (no charged particles) | Yes (delocalised electrons) |
| Malleable? | No (brittle — shatters) | N/A | Yes (layers slide) |
🧠 Memory Trick
"Metals are generous — they GIVE UP their electrons to everyone." In ionic bonding, a metal gives electrons to one specific non-metal. But in metallic bonding, each metal atom gives up its outer electrons to the ENTIRE structure. Nobody owns the electrons — they belong to everybody. That is what "delocalised" means — not localised to any one atom.
🧠 Memory Trick
"SWIM to remember metallic bonding." Sea of electrons, Wires (ductile), I have Ions (positive), Malleable (layers slide). SWIM through metals!
⚠ Exam Tip
The exam may ask you to explain why ionic compounds are brittle but metals are malleable. The key difference: in an ionic lattice, if layers slide, positive ions end up next to positive ions and negative next to negative — the repulsion shatters the crystal. In a metal, when layers slide, the delocalised electrons simply rearrange around the new positions — there is nothing directional to cause repulsion.
⚠ Exam Tip
Do not confuse the delocalised electrons in metals with those in graphite. Both involve delocalised electrons, but: metals have metallic bonding (positive ions + sea of electrons), while graphite has covalent bonding within layers and weak forces between layers, with delocalised electrons between the layers. Graphite is NOT metallic bonding!
Worked Example
Describe metallic bonding and use it to explain why metals conduct electricity.
Step 1: Describe the structure
In metals, the atoms lose their outer shell electrons to form positive ions (cations) arranged in a regular lattice. The lost electrons become delocalised — they are free to move throughout the entire structure, forming a "sea" of electrons.
Step 2: Describe the bond
The metallic bond is the strong electrostatic attraction between the positive metal ions and the sea of delocalised negative electrons.
Step 3: Explain conductivity
When a voltage is applied, the delocalised electrons can flow through the metal structure towards the positive terminal. Because these electrons are free to move, they can carry an electric charge, making metals good conductors of electricity.
Answer: Metallic bonding involves positive metal ions in a regular lattice surrounded by a sea of delocalised electrons. The bond is the electrostatic attraction between the ions and the electron sea. Metals conduct because the delocalised electrons are free to move and carry charge when a voltage is applied.
Worked Example
Explain why metals are malleable but ionic compounds are brittle.
Step 1: Metals — what happens when force is applied
In metals, when a force is applied, the layers of positive ions can slide over each other into new positions. The sea of delocalised electrons is not directional — it rearranges around the ions in their new positions immediately, so the metallic bond is maintained. The metal changes shape without breaking.
Step 2: Ionic compounds — what happens when force is applied
In ionic compounds, the ions are held in fixed positions in a lattice by electrostatic attraction between positive and negative ions. If a force causes the layers to slide, ions of the same charge end up next to each other (positive next to positive, negative next to negative). The resulting repulsion causes the lattice to shatter.
Answer: Metals are malleable because layers of ions slide and the electron sea rearranges to maintain the bond. Ionic compounds are brittle because sliding brings like charges together, causing repulsion that shatters the crystal.
Apply It: Real-World Chemistry
Metals are everywhere in our lives — from electrical wiring to jewellery to the steel in bridges. Every useful property of metals can be traced back to metallic bonding. Let us explore how.
1
Copper is used for almost all electrical wiring worldwide. India's copper demand exceeded 1.5 million tonnes in 2025, driven by the expansion of the electrical grid to rural areas. Copper wires are drawn to be as thin as 0.025 mm for use in transformer coils.
Why is copper used for electrical wiring rather than other materials?
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The Chemistry
Copper is an excellent electrical conductor because it has delocalised electrons that are free to flow through the metal when a voltage is applied. This is due to metallic bonding — copper atoms lose their outer electrons, which form a "sea" around the positive copper ions. When a voltage is applied, these electrons drift towards the positive terminal, carrying charge. Copper is also ductile — it can be drawn into very thin wires without breaking. This is because the layers of copper ions can slide over each other while the delocalised electrons rearrange to maintain the metallic bond. Copper is preferred over other metals because it has the second-highest electrical conductivity of any metal (after silver, which is too expensive), is relatively cheap, and does not corrode easily.
The Big Idea
Every copper wire in your house — from the charger for your phone to the wires inside the walls — relies on two properties of metallic bonding: conductivity (delocalised electrons carry charge) and ductility (layers slide without breaking the bond). Both properties come directly from the sea of delocalised electrons.
2
Gold is hammered into incredibly thin sheets called "gold leaf," which can be as thin as 0.1 micrometres (about 400 atoms thick). In India, gold leaf is used to decorate temples, sweets (vark on barfi and mithai), and in Ayurvedic medicine. A single gram of gold can be beaten into a sheet covering about 1 square metre.
Why can gold be hammered into such incredibly thin sheets without breaking?
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The Chemistry
Gold is the most malleable metal known. Its extreme malleability is due to metallic bonding. In gold, the positive ions are arranged in a regular lattice, surrounded by a sea of delocalised electrons. When gold is hammered, the layers of positive ions slide over each other into new positions. The sea of delocalised electrons is not directional — it simply rearranges around the ions wherever they move, maintaining the strong metallic bond. No matter how thin the sheet becomes, as long as there are layers of ions with electrons around them, the metallic bond holds. Gold is particularly malleable because it has a face-centred cubic crystal structure, which has many slip planes allowing easy sliding.
The Big Idea
The reason metals can be hammered and stretched is fundamentally different from other bonding types. In an ionic crystal, the same sliding would bring like charges together and shatter the material. In a metal, the non-directional electron sea simply flows around the new arrangement. The metallic bond is amazingly forgiving of structural rearrangement.
3
Steel (iron with carbon) is used to construct bridges, buildings, and railway tracks across the world. The Howrah Bridge in Kolkata, built in 1943, uses over 26,000 tonnes of steel. Unlike glass or ceramics, steel bends under stress rather than shattering — a critical safety feature.
Why does steel bend rather than shatter when a large force is applied?
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The Chemistry
Steel is an alloy based on iron, which has metallic bonding. When a large force is applied to steel, the layers of positive iron ions slide over each other. The sea of delocalised electrons redistributes around the displaced ions, maintaining the metallic bond throughout the deformation. The structure changes shape but does not break. This is called ductility or malleability. In contrast, materials with ionic bonds (like ceramics) or giant covalent structures (like glass, which is based on SiO₂) shatter when stressed, because their bonds are directional — any displacement disrupts the bond arrangement and causes catastrophic failure.
The Big Idea
The ability of metals to deform without breaking is not just convenient — it is a critical safety feature. A steel bridge that bends under overload gives warning; a ceramic bridge that shatters gives none. This difference in behaviour is entirely due to the difference in bonding: the non-directional electron sea in metals vs the directional bonds in ionic and covalent materials.
4
In British kitchens, aluminium and stainless steel saucepans are used for cooking. When you place a metal saucepan on a gas hob, the handle quickly becomes hot even though the flame only touches the base. Metal spoons left in hot soup also become too hot to hold. By contrast, wooden spoons remain cool.
Why do metals conduct heat so much better than non-metals like wood?
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The Chemistry
Metals conduct heat efficiently because of their delocalised electrons. When the base of a saucepan is heated, the delocalised electrons near the flame gain kinetic energy and move faster. These energetic electrons collide with other delocalised electrons further along the metal, transferring kinetic energy rapidly through the structure. This is much faster than the conduction that happens in non-metals (like wood), where heat can only be transferred by vibrations passing from atom to atom through fixed bonds. Wood has no delocalised electrons, so heat travels very slowly through it. This is why wooden spoons are used for cooking — they do not conduct heat to your hand.
The Big Idea
Both thermal and electrical conductivity in metals come from the same source: delocalised electrons. Good electrical conductors are almost always good thermal conductors too, because in both cases it is the mobile electrons doing the carrying. This is not a coincidence — it is a direct consequence of metallic bonding.
5
Tungsten has the highest melting point of all metals — 3422°C. It is used in light bulb filaments because it glows white-hot without melting. Sodium, on the other hand, has a melting point of only 98°C and can be cut with a knife. Both are metals with metallic bonding.
Why does tungsten have a much higher melting point than sodium, even though both have metallic bonding?
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The Chemistry
Both tungsten and sodium have metallic bonding (positive ions in a sea of delocalised electrons), but the strength of the metallic bond differs greatly. Tungsten has more outer shell electrons to contribute to the "sea" (it is a transition metal with electrons in multiple shells available for delocalisation). It also has a smaller ionic radius and higher charge density. The more delocalised electrons there are, and the greater the charge on the ions, the stronger the electrostatic attraction between the ions and the electron sea. Sodium has only 1 outer electron to contribute, and its ion is relatively large with a low charge (Na⁺). So sodium's metallic bond is much weaker, giving it a low melting point. Tungsten's many delocalised electrons and high ionic charge create very strong metallic bonds, giving it the highest melting point of any metal.
The Big Idea
Not all metallic bonds are equal. The strength depends on how many electrons each atom contributes to the sea and the charge/size of the resulting ion. More delocalised electrons and higher ionic charge = stronger metallic bonding = higher melting point. The same model explains all metals, from the softest (sodium) to the toughest (tungsten).
Test Yourself: Metallic Bonding
20 questions — all Supplement content
Your Score:
0 / 20
QUESTION 1 OF 20
In metallic bonding, the outer electrons of metal atoms:
In metallic bonding, the outer electrons are delocalised — they leave their original atoms and form a "sea" of electrons that is shared throughout the entire metal structure. They do not belong to any particular atom.
QUESTION 2 OF 20
What is a metallic bond?
A metallic bond is the strong electrostatic attraction between the lattice of positive metal ions and the sea of delocalised electrons. Option A describes ionic bonding. Option B describes covalent bonding. Option D describes the forces between layers in graphite.
QUESTION 3 OF 20
Metals conduct electricity because:
Metals conduct electricity because their delocalised electrons are free to flow through the structure towards the positive terminal when a voltage is applied. The positive ions do NOT move — they stay in their fixed positions in the lattice. It is only the electrons that move.
QUESTION 4 OF 20
Metals are malleable because:
Metals are malleable because the layers of positive ions can slide over each other when a force is applied. The sea of delocalised electrons is not directional, so it rearranges around the ions in their new positions. The metallic bond is maintained even after the ions have moved to new positions.
QUESTION 5 OF 20
In a metallic structure, the particles in the lattice are:
In a metal, each atom has lost its outer electron(s) to become a positive ion (cation). These positive ions are arranged in a regular lattice. A mixture of positive and negative ions describes an ionic lattice, not a metallic one.
QUESTION 6 OF 20
Why are ionic compounds brittle but metals are malleable?
In ionic compounds, if layers slide, positive ions end up next to positive ions (and negative next to negative). The strong repulsion between like charges shatters the crystal. In metals, the sea of delocalised electrons is not directional — it flows around the ions in their new positions, maintaining the bond. That is the key difference.
QUESTION 7 OF 20
Which type of bonding involves a "sea of delocalised electrons"?
The "sea of delocalised electrons" is a defining feature of metallic bonding. In ionic bonding, electrons are transferred. In covalent bonding, electrons are shared in pairs. Intermolecular forces are weak forces between molecules.
QUESTION 8 OF 20
Copper is used for electrical wires because it:
Copper is used for wires because: (1) it has delocalised electrons that make it an excellent conductor of electricity, and (2) it is ductile — it can be drawn into thin wires because layers of ions slide while the electron sea maintains the bond. Copper is NOT an insulator and does NOT have a low melting point.
QUESTION 9 OF 20
Why do metals generally have high melting points?
Metals generally have high melting points because the metallic bond (electrostatic attraction between positive ions and the sea of delocalised electrons) is strong. A lot of energy is needed to overcome this attraction and allow the ions to move freely. Metals do NOT have covalent bonds or intermolecular forces.
QUESTION 10 OF 20
Which of these is NOT a property explained by metallic bonding?
Metallic bonding explains electrical conductivity (delocalised electrons flow), malleability (layers slide, electrons rearrange), and thermal conductivity (delocalised electrons transfer kinetic energy). Forming coloured compounds is a property of transition metals related to their d-orbitals, not directly related to metallic bonding.
QUESTION 11 OF 20
A student writes: "Metals conduct electricity because the positive ions move through the metal." Is this correct?
The student is wrong. In solid metals, the positive ions are fixed in their positions in the regular lattice. They vibrate but do not move from place to place. It is the delocalised electrons that flow through the structure to carry charge. This is a very common exam mistake — always say "delocalised electrons move," NOT "ions move."
QUESTION 12 OF 20
Metals are good conductors of heat because:
Metals conduct heat efficiently primarily because the delocalised electrons can gain kinetic energy when heated and quickly transfer this energy through the structure by colliding with other electrons. While ion vibrations contribute slightly, the main mechanism is the mobile electrons — this is why metals conduct heat much better than non-metals.
QUESTION 13 OF 20
"Ductile" means a metal can be:
Ductile = can be drawn (pulled/stretched) into thin wires. Malleable = can be hammered into thin sheets. Both properties exist because layers of metal ions can slide while the electron sea maintains the metallic bond. Breaking into pieces is called "brittle" — the opposite of ductile and malleable.
QUESTION 14 OF 20
Which comparison between metals and ionic compounds is correct?
Metals conduct in the solid state because their delocalised electrons are free to move. Ionic compounds do NOT conduct as solids (ions are fixed in lattice) — they only conduct when molten or dissolved (ions become free to move). Metals are malleable (not brittle), and ionic compounds do NOT have delocalised electrons.
QUESTION 15 OF 20
Both graphite and copper conduct electricity. The conducting particles are:
Both graphite and copper conduct through delocalised electrons. In copper, these come from metallic bonding (outer electrons form a sea). In graphite, each carbon's 4th electron is delocalised between the layers. Different types of bonding, but the conducting particles are delocalised electrons in both cases.
QUESTION 16 OF 20
Gold can be hammered into sheets only 0.1 micrometres thick. This property is called:
Being hammered into thin sheets is malleability. Being drawn into wires is ductility. Both are due to layers of ions sliding while the electron sea rearranges. Conductivity is the ability to carry charge or heat. Brittleness is the opposite — shattering when hit.
QUESTION 17 OF 20
Sodium has a low melting point compared to iron because:
Both sodium and iron have metallic bonding, but sodium has only 1 outer electron to contribute to the sea (forming Na⁺), while iron can contribute more electrons from multiple shells. More delocalised electrons and higher ionic charge = stronger metallic bond = higher melting point. Sodium IS a metal (Group I).
QUESTION 18 OF 20
Which of the following correctly describes the structure of a metal?
A metal consists of positive ions (the atoms have lost their outer electrons) arranged in a regular lattice, surrounded by a sea of delocalised electrons. Option D describes an ionic lattice. Option B describes covalent bonding. Option A has the charges reversed.
QUESTION 19 OF 20
A metal rod feels cold when you touch it. This is because:
A metal rod at room temperature is the same temperature as a wooden table, but the metal feels colder because it conducts heat away from your warm skin much faster. The delocalised electrons in the metal transfer kinetic energy rapidly through the structure, quickly moving heat away from your hand. Wood conducts heat slowly, so it feels warmer even at the same temperature.
QUESTION 20 OF 20
Which row correctly matches the bonding type with the particles and properties?
Metallic: positive ions in a lattice with a sea of delocalised electrons; generally high melting points (strong metallic bonds); malleable (layers slide, electrons rearrange); conducts as a solid (delocalised electrons). The other options have incorrect matches — ionic has ions but is brittle (not malleable) and does not conduct as a solid.