Chemical Bonds: Ionic, Covalent, and Metallic Bonding Explained

Why Do Atoms Bond?

Atoms are most stable when their outer electron shell is full. Noble gases (Group 18) already have full outer shells, which is why they are unreactive. Every other element needs to gain, lose, or share electrons to reach this stable configuration — and that process is chemical bonding.

The type of bond that forms depends on the electronegativity difference between the atoms involved:

• Large difference (typically > 1.7) → Ionic bond
• Small or zero difference (< 1.7) → Covalent bond
• Metal atoms only → Metallic bond

Learning Goals: By the end of this guide, you should be able to:

1. Explain the formation of ionic, covalent, and metallic bonds using electron transfer/sharing models.
2. Draw dot-and-cross diagrams for simple ionic and covalent compounds.
3. Predict bond type from electronegativity values.
4. Compare the physical properties of ionic, covalent, and metallic substances.
5. Identify common exam mistakes and avoid them.

Ionic Bonding

An ionic bond forms when one atom transfers one or more electrons to another atom. This creates two oppositely charged ions that are held together by electrostatic attraction.

How It Works

1. A metal atom (low electronegativity) loses electrons → becomes a positive cation (e.g., Na⁺).
2. A non-metal atom (high electronegativity) gains those electrons → becomes a negative anion (e.g., Cl⁻).
3. The opposite charges attract, holding the ions together in a giant ionic lattice.

Example: Sodium Chloride (NaCl)

• Sodium (Na): electron configuration $[2, 8, 1]$ → loses 1 electron → $Na^+$ $[2, 8]$
• Chlorine (Cl): electron configuration $[2, 8, 7]$ → gains 1 electron → $Cl^-$ $[2, 8, 8]$
• Both ions now have the stable electron configuration of a noble gas.

Properties of Ionic Compounds

Covalent Bonding

A covalent bond forms when two atoms share one or more pairs of electrons. This usually occurs between two non-metal atoms with similar electronegativities.

Types of Covalent Bonds

Polar vs. Non-Polar Covalent Bonds

When two atoms have different electronegativities, the shared electrons are pulled closer to the more electronegative atom. This creates a polar covalent bond with partial charges ($\delta^+$ and $\delta^-$).

• Non-polar: $H_2$, $Cl_2$, $O_2$ (identical atoms → equal sharing)
• Polar: $H-Cl$, $H-O$ (different atoms → unequal sharing)
• Ionic: $Na-Cl$ (so unequal that transfer occurs)

Think of it as a spectrum: non-polar covalent → polar covalent → ionic, driven by increasing electronegativity difference.

Dative (Coordinate) Covalent Bonds

A special case where both electrons in the shared pair come from the same atom. Once formed, a dative bond is identical to a normal covalent bond.

Example: The ammonium ion $NH_4^+$ — when $NH_3$ donates its lone pair to $H^+$.

Properties of Simple Covalent Molecules

Important distinction: The covalent bonds within each molecule are very strong. It's the forces between molecules that are weak. Always specify whether you mean intramolecular or intermolecular forces in exam answers.

Metallic Bonding

A metallic bond is the electrostatic attraction between a lattice of positive metal cations and a sea of delocalised electrons.

How It Works

Metal atoms lose their outer electrons, which become delocalised (free to move throughout the structure). The resulting positive ions are arranged in a regular lattice, held together by the attraction to the shared electron sea.

Properties of Metals

Comparison: All Three Bond Types

Worked Examples

Example 1: Predicting Bond Type

Question: What type of bond forms between magnesium (electronegativity 1.31) and oxygen (electronegativity 3.44)?

Solution:

Since $\Delta EN > 1.7$, this is an ionic bond. Magnesium transfers 2 electrons to oxygen:

$Mg \rightarrow Mg^{2+} + 2e^-$

$O + 2e^- \rightarrow O^{2-}$

Product: MgO (magnesium oxide) — a giant ionic lattice with very high melting point (2852°C).

Example 2: Drawing Dot-and-Cross for Water

Question: Draw the dot-and-cross diagram for $H_2O$.

Solution:

1. Oxygen has 6 outer electrons, each hydrogen has 1.
2. Oxygen shares 1 electron with each hydrogen → 2 bonding pairs.
3. Oxygen has 2 remaining lone pairs.
4. Total: 2 bonding pairs + 2 lone pairs = 4 electron domains → bent shape (VSEPR).

The O–H bonds are polar covalent because oxygen ($EN = 3.44$) is more electronegative than hydrogen ($EN = 2.20$), creating $\delta^+$ on H and $\delta^-$ on O.

Example 3: Why Does Diamond Have a High Melting Point?

Question: Explain why diamond has a very high melting point despite being a covalent substance.

Solution: Diamond is a giant covalent structure (not a simple molecule). Each carbon atom forms 4 strong covalent bonds to other carbon atoms in a tetrahedral arrangement, creating a rigid 3D network. To melt diamond, you must break many strong C–C covalent bonds, which requires enormous energy. This is different from simple covalent molecules where only weak intermolecular forces need to be overcome.

Common Mistakes

1. "Ionic bonds are stronger than covalent bonds" — This is not always true. Individual covalent bonds (e.g., C–C in diamond) can be very strong. The comparison depends on the specific substances.

2. Confusing intramolecular and intermolecular forces — Water has strong O–H covalent bonds within molecules but weak hydrogen bonds between molecules. The low boiling point is due to weak intermolecular forces, not weak bonds.

3. "NaCl is a molecule" — Ionic compounds do not form discrete molecules. NaCl is a formula unit representing the simplest ratio of ions in the giant lattice.

4. Forgetting dative bonds exist — In $NH_4^+$, one of the four N–H bonds is dative, but once formed it is identical to the others. Don't label it differently in structural formulas.

5. Assuming metals always have high melting points — Mercury (Hg) is a liquid at room temperature. Metallic bonding strength depends on the number of delocalised electrons and ion size.

Exam Tips (A-Level / AP)

• When asked to explain a property, always link back to the type of bonding and the structure (giant vs. simple molecular).
• For comparison questions, use a table format in your answer — examiners reward structured responses.
• If the question asks about electrical conductivity, always specify the state (solid, liquid, aqueous) — many students lose marks by not doing this.
• Remember: giant covalent structures (diamond, silicon dioxide, graphite) behave very differently from simple covalent molecules. Don't generalise "covalent = low melting point."

Frequently Asked Questions

What is the difference between ionic and covalent bonding?

Ionic bonding involves the complete transfer of electrons from a metal to a non-metal, creating charged ions held together by electrostatic attraction. Covalent bonding involves the sharing of electron pairs between two non-metal atoms.

Why do ionic compounds conduct electricity when dissolved but not when solid?

In solid ionic compounds, ions are locked in a rigid lattice and cannot move. When dissolved or molten, the ions are free to move, allowing them to carry charge and conduct electricity.

Can a bond be both ionic and covalent?

All bonds exist on a spectrum. A polar covalent bond has some ionic character, and ionic bonds can have some covalent character (called polarisation). The classification depends on the degree of electron transfer or sharing.

Why is diamond hard but graphite is soft?

Both are giant covalent structures of carbon, but diamond has a 3D rigid network of C–C bonds, while graphite has layers with weak van der Waals forces between them. The layers in graphite can slide, making it soft.

What determines whether a bond is polar or non-polar?

The electronegativity difference between the two bonded atoms. If the difference is zero or very small, the bond is non-polar. If significant (but < 1.7), it is polar covalent.

Related Topics

• VSEPR Theory — Predict the 3D shape of molecules from their electron pairs.
• Intermolecular Forces — Understand the forces between molecules that determine physical properties.
• Orbital Hybridization — How atomic orbitals combine to form bonding orbitals.