What Is Chemical Bonds: Ionic, Covalent And Metallic? A Visual Guide for OxfordAQA IGCSE Chemistry Students

Three Ways Atoms Stick Together

Here is a question worth asking before you memorise anything: what is chemical bonds: ionic, covalent and metallic igcse students are actually being asked to compare? At heart, it is just three different answers to the same question: how do atoms end up stuck together? Sometimes they swap electrons entirely. Sometimes they share them. And sometimes they all pool their outer electrons together into one shared "soup." That is genuinely the whole story, and everything else in oxfordaqa igcse chemical bonds: ionic, covalent and metallic is detail layered onto those three ideas, once each one is oxfordaqa igcse chemistry explained clearly rather than left as a vague memory of a diagram from class.

Definition

Chemical bonding is the process by which atoms achieve a stable, noble-gas electron arrangement, either by transferring electrons between atoms (ionic bonding), by sharing pairs of electrons between atoms (covalent bonding), or by all the atoms in a structure contributing their outer electrons to a shared, delocalised pool (metallic bonding). This is the core oxfordaqa igcse chemistry definition to build everything else on, so it is worth being able to state clearly which of the three mechanisms is operating before you attempt any question on this topic.

Key Facts

• Ionic bonding involves electrons transferring completely from a metal atom to a non-metal atom, producing charged ions held together by electrostatic attraction.
• Covalent bonding involves atoms sharing pairs of electrons; it happens between non-metal atoms.
• Metallic bonding involves delocalised electrons moving freely through a giant structure of positive metal ions.
• The type of bonding a substance has determines its melting point, boiling point, and whether it conducts electricity.
• You should be able to draw or complete diagrams for all three bonding types when given appropriate information.

Ionic Bonding: Chemical Bonds Ionic Covalent and Metallic Explained From the Ground Up

Picture a sodium atom standing next to a chlorine atom. Sodium has one electron in its outer shell that it would honestly rather not have; chlorine has seven and would love an eighth to complete its shell. So sodium hands its spare electron over. Sodium becomes a positively charged ion, Na+, because it now has one more proton than electron; chlorine becomes a negatively charged ion, Cl-, because it now has one more electron than proton. Both ions now have the stable electron arrangement of a noble gas.

This is not a one-off event between two atoms floating alone; in a real ionic compound like sodium chloride, huge numbers of oppositely charged ions arrange themselves into a giant, repeating lattice, held together by strong electrostatic forces of attraction acting in all directions. That lattice structure is exactly why ionic compounds have high melting and boiling points: breaking apart that many strong attractions, all at once, in every direction, takes a great deal of energy.

You should be able to relate the charge on a simple ion to its group number in the periodic table: Group 1 metals form 1+ ions, Group 7 non-metals (the halogens) form 1- ions, and so on. Group 1 elements, the alkali metals, react with non-metals to form ionic compounds where the metal ion carries a single positive charge; Group 7 elements, the halogens, react with metals to form ionic compounds where the halide ion carries a single negative charge.

Worked example: predict the ionic compound formed between magnesium (Group 2) and oxygen (Group 6). Magnesium loses two electrons to form Mg2+; oxygen gains two electrons to form O2-. The compound formed is magnesium oxide, MgO, with a 1:1 ratio of ions because the charges balance exactly.

Covalent Bonding

Now picture two chlorine atoms, each with seven outer electrons, neither one willing to simply hand an electron away. Instead, they share a pair of electrons between them, one contributed by each atom, so that both atoms effectively count that shared pair as part of their own outer shell. That shared pair is a single covalent bond, and it is strong, because both nuclei are attracted to the same pair of electrons sitting between them.

Some covalently bonded substances exist as small, separate molecules: water (H2O), ammonia (NH3), methane (CH4), hydrogen (H2), oxygen (O2) and hydrogen chloride (HCl) are the examples named on the specification, and you should be able to represent their bonding using dot-and-cross or displayed structural diagrams. Other covalent substances form giant covalent structures, or macromolecules, where huge numbers of atoms are linked together entirely by covalent bonds: diamond and silicon dioxide are the two named examples.

This distinction between simple molecules and giant covalent structures explains a property that trips a lot of students up: simple molecular substances have low melting and boiling points, while giant covalent structures have extremely high ones, even though both types are held together internally by strong covalent bonds. The difference is not in the strength of the covalent bonds themselves, but in what actually has to be overcome to melt or boil the substance.

Metallic Bonding

Metals take a third approach entirely. Picture the metal atoms as a stack of positive ions, arranged in a regular, repeating pattern, with the electrons from their outer shells no longer belonging to any one atom at all. Instead, those electrons are delocalised: free to drift through the entire structure like a shared electron "sea" surrounding fixed positive ions. This is why metals conduct electricity and heat so well (the free electrons carry charge and energy through the structure) and why they can be bent or hammered into new shapes without shattering (the layers of positive ions can slide past each other while the delocalised electrons keep the whole thing held together).

Linking Bonding to Properties: The Exam's Favourite Question

By far the most common exam pattern on this topic gives you a substance, or a diagram of its structure, and asks you to explain a property using the correct type of bonding. A quick reference, worth memorising as a table rather than as separate facts:

Notice the ionic row carefully: a solid ionic compound does not conduct electricity, because the ions are locked in place in the lattice and cannot move to carry charge. It is only when the compound is melted or dissolved in water, freeing the ions to move, that it conducts.

Why This One Topic Keeps Reappearing

You will notice this same bonding-and-structure reasoning quietly showing up again later, in totally different-looking topics. Whether a metal conducts and can be extracted, why an alkane is a gas while a longer hydrocarbon is a thick liquid, why a polymer softens on heating: each of these questions is really the bonding-and-structure argument above, applied to a new substance. Getting genuinely comfortable with the three bonding types now, rather than memorising them as an isolated block for one test, pays off across the rest of the specification, not just on questions that use the word "bonding" directly in the stem.

Diamond and graphite are a particularly instructive pair to hold in mind together, because they are both giant covalent structures made entirely of carbon, yet they behave very differently: diamond is hard and does not conduct electricity, because every carbon atom forms four strong covalent bonds in a rigid three-dimensional lattice with no free electrons; graphite is soft, slippery, and does conduct electricity, because its carbon atoms form only three bonds each in flat layers, leaving one delocalised electron per atom free to move within each layer. The same element, the same type of bonding overall, but a different arrangement produces strikingly different properties, which is exactly the kind of comparison a good exam question likes to set.

Common Mix-Ups to Watch For

• Describing ionic bonding as "sharing" electrons, which is the language of covalent bonding, not ionic.
• Saying a simple molecular substance melts because its "covalent bonds break," when actually it is the weak intermolecular forces between separate molecules that are overcome; the covalent bonds inside each molecule stay intact.
• Forgetting that a solid ionic compound does not conduct electricity, and only stating the molten or dissolved case.
• Mixing up which elements go where: metals lose electrons and form positive ions, non-metals gain electrons and form negative ions, and this never reverses.

Self-Check Questions

1. Explain, in terms of electron transfer, how sodium chloride forms an ionic bond, and describe the structure of solid sodium chloride.
2. Draw a dot-and-cross diagram showing the covalent bonding in a molecule of ammonia, NH3.
3. Explain, in terms of delocalised electrons, why metals conduct electricity.
4. Explain why a simple molecular substance has a low boiling point despite having strong covalent bonds within each molecule.
5. Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride does.

Once you can sort any given substance into one of the three bonding types and connect it straight to a property, chemical bonds: ionic, covalent and metallic explained this way should stop feeling like a memory test and start feeling like a single, reusable piece of reasoning you can apply to any unfamiliar substance the exam throws at you. Keep these oxfordaqa igcse chemistry notes close whenever you revise structure and bonding, since almost every other topic on the specification, from metals to organic chemistry, quietly assumes you already have this one solid.