How Metals and Non-metals React — Ionic Bonding
A metal atom gives up one or more electrons to become a positive ion; a non-metal atom accepts them to become a negative ion. The compound the two ions form together — held by pure electrical attraction — is an ionic compound.
In this lesson you will
- Explain reactivity as the tendency of an atom to reach a stable, noble-gas-like outer shell.
- Show the formation of sodium chloride and magnesium chloride as electron transfers, with correct charges.
- Define ionic (electrovalent) compounds and name the cation and anion in a given formula.
- List the four characteristic properties of ionic compounds and explain each from the strength of the ionic bond.
- Explain why an ionic compound conducts electricity when molten or dissolved but not as a solid.
Why atoms react at all
Noble gases, with a completely filled outermost shell, show almost no chemical activity at all. Every other element reacts in a way that moves it toward that same filled-shell stability. A sodium atom has just one electron in its outermost shell — losing it leaves a complete, stable shell underneath. A chlorine atom has seven electrons in its outermost shell — gaining just one completes an octet. Put a sodium atom and a chlorine atom together, and the sodium atom’s spare electron transfers to chlorine’s nearly-full shell: both atoms end up stable.
Forming sodium chloride
Sodium loses its one outer electron to become a sodium cation, Na⁺. Chlorine gains that electron to become a chloride anion, Cl⁻. The nucleus of each atom never changes — only the electron count — so the resulting charges come purely from the mismatch between protons and electrons. The oppositely charged Na⁺ and Cl⁻ ions then attract each other electrostatically and pack together as sodium chloride, NaCl. Crucially, NaCl does not exist as individual molecules — it is a vast, repeating lattice of alternating positive and negative ions.
Forming magnesium chloride
Magnesium has two electrons in its outer shell and must lose both to reach stability, becoming Mg²⁺. Since each chlorine atom can only accept one electron, it takes two chlorine atoms to receive magnesium’s two electrons, each becoming a separate Cl⁻ ion. The result is one Mg²⁺ ion for every two Cl⁻ ions — magnesium chloride, MgCl2.
Compounds formed this way, by the transfer of electrons from a metal to a non-metal, are called ionic or electrovalent compounds.
Properties of ionic compounds
- Physical nature: hard, because of the strong attraction between oppositely charged ions — but brittle, breaking apart when struck rather than bending.
- Melting and boiling points: both very high — a considerable amount of energy is needed to break the strong attraction holding the ion lattice together.
- Solubility: generally soluble in water, insoluble in solvents like petrol and kerosene.
- Conduction of electricity: conducting electricity means moving charged particles. In the solid state the ions are locked in a rigid lattice and cannot move, so a solid ionic compound does not conduct. Melt it, or dissolve it in water, and the ions become free to move toward the electrodes — and it conducts well.
Key terms
- Cation
- A positively charged ion, formed when an atom loses one or more electrons — typically from a metal.
- Anion
- A negatively charged ion, formed when an atom gains one or more electrons — typically a non-metal.
- Ionic (electrovalent) compound
- A compound formed by the transfer of electrons from a metal atom to a non-metal atom, held together by the electrostatic attraction between the resulting oppositely charged ions.
Key relationships
- Na loses 1 electron → Na⁺ (2,8,1 → 2,8); Cl gains 1 electron → Cl⁻ (2,8,7 → 2,8,8); Na⁺ + Cl⁻ → NaCl.
- Mg loses 2 electrons (one to each of two Cl atoms) → Mg²⁺; each Cl becomes Cl⁻ → MgCl2.
- Ionic compounds exist as aggregates (lattices) of ions, never as individual molecules.
- Ionic compounds are hard but brittle solids with high melting and boiling points.
- Ionic compounds are generally soluble in water, insoluble in kerosene or petrol.
- Ionic compounds do not conduct electricity as solids, but do conduct when molten or dissolved in water.
Formulae
Sodium chloride formation
Na → Na⁺ + e⁻ ; Cl + e⁻ → Cl⁻ ; Na⁺ + Cl⁻ → NaCl
Magnesium chloride formation
Mg → Mg²⁺ + 2e⁻ ; 2Cl + 2e⁻ → 2Cl⁻ ; Mg²⁺ + 2Cl⁻ → MgCl2
Activities
Activity 3.13 — Testing the properties of ionic compounds
What you do: Heat samples of sodium chloride, potassium iodide and barium chloride directly on a flame; try dissolving each in water, petrol and kerosene; then test a solution’s conductivity with graphite electrodes in a simple circuit.
What you see: The samples resist melting on an ordinary flame (their real melting points are far higher); they dissolve in water but not in petrol or kerosene; the bulb lights up once electrodes are placed in the solution.
What it shows: Ionic compounds are high-melting, water-soluble solids that conduct electricity only when their ions are free to move — in solution or molten, never as a rigid solid.
What you should be able to do
Recognise
- Which element loses electrons (becomes a cation) and which gains them (becomes an anion) in a given reaction.
- The correct charge on a cation or anion from the number of electrons transferred.
Explain
- Why ionic compounds have high melting and boiling points.
- Why an ionic compound conducts electricity molten or in solution, but not as a solid.
- Why ionic compounds are brittle even though they are hard.
Draw
- Electron-dot (Lewis) structures for sodium, chlorine, magnesium and oxygen, and the ions they form.
State
- The four properties of ionic compounds: physical nature, melting/boiling point, solubility, and electrical conduction.
NCERT alignment
This topic corresponds to NCERT Class 10 Science, Chapter “Metals and Non-metals”, sections 3.3, 3.3.1; activities 3.13. Explanations, visuals and worked material here are original.
Quick revision
- Metals lose electrons to form cations (positive); non-metals gain electrons to form anions (negative).
- Na → Na⁺ + e⁻; Cl + e⁻ → Cl⁻; together → NaCl.
- Mg → Mg²⁺ + 2e⁻, shared with TWO chlorine atoms → MgCl2.
- Ionic compounds: hard but brittle, high melting/boiling point, soluble in water not kerosene/petrol, conduct only molten or dissolved.
- Ionic compounds exist as ion lattices, never as individual molecules.