The main difference between an ionic and a covalent bond is how electrons are involved. An ionic bond is the electrostatic attraction between oppositely charged ions, usually formed after electrons are transferred from a metal atom to a non-metal atom. A covalent bond is the attraction between two nuclei and a shared pair of electrons, usually occurring between non-metal atoms.
That distinction explains the structures and properties produced by each model. Ionic compounds generally form extended lattices, while covalent substances may exist as small molecules or giant covalent networks. For IB Chemistry, identifying the bond type is only the beginning: strong answers connect bonding to structure, structure to forces, and forces to observable properties.
The current IB Chemistry course, first assessed in 2025, places these ideas in Structure 2: Models of bonding and structure. The ionic and covalent models are treated in Structure 2.1 and Structure 2.2, while Structure 2.4 emphasizes that bonding is better understood as a continuum than as a set of perfectly separate categories. For wider syllabus context, use the topic-wide IB Chemistry Structure Explained exam guide; this article focuses specifically on the ionic versus covalent comparison.
Ionic vs covalent bond at a glance
| Feature | Ionic bonding | Covalent bonding |
|---|---|---|
| Fundamental attraction | Attraction between oppositely charged ions | Attraction between both nuclei and a shared electron pair |
| Typical elements involved | Metal and non-metal | Two non-metals |
| Treatment of valence electrons | Electron transfer produces ions | Electron pairs are shared |
| Particles in the resulting substance | Cations and anions | Molecules, atoms in a network, or polyatomic ions within a larger structure |
| Typical structure | Giant ionic lattice | Simple molecular or giant covalent structure |
| Direction of bonding | Electrostatic attraction acts in all directions through the lattice | Covalent bonds have specific directions between bonded atoms |
| Typical melting point | High | Often low for molecular substances, but very high for giant covalent structures |
| Electrical conductivity | Not as a solid; often conducts when molten or aqueous | Usually poor, with important exceptions such as graphite |
| Formula meaning | Simplest whole-number ratio of ions | Actual molecular composition or simplest ratio in a network |
| Example | NaCl, MgO | H₂O, CH₄, SiO₂ |
These are patterns rather than absolute rules. Saying that all covalent substances have low melting points, for example, ignores covalent networks such as diamond and silicon dioxide.
How does an ionic bond form?
Ionic bonding is commonly introduced through electron transfer. A metal atom loses one or more valence electrons to form a positive ion, or cation, while a non-metal atom gains electrons to form a negative ion, or anion.
For sodium chloride:
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
The transferred electron creates Na⁺ and Cl⁻ ions. The ionic bond is not the transfer itself. It is the electrostatic force of attraction between the resulting oppositely charged ions.
This distinction matters in exams. If asked to define ionic bonding, writing only “electrons are transferred” describes how the ions may form but does not define the bond. A more complete answer is:
An ionic bond is the electrostatic attraction between oppositely charged ions.
Ionic compounds form lattices, not isolated pairs
A diagram of Na⁺ beside Cl⁻ can make sodium chloride appear to contain individual NaCl units. In reality, solid sodium chloride forms a giant three-dimensional ionic lattice in which each ion is surrounded by ions of opposite charge.
Consequently, NaCl is described by a formula unit, not as a discrete NaCl molecule. Its formula gives the simplest whole-number ratio required for electrical neutrality. In MgCl₂, for example, one Mg²⁺ ion balances two Cl⁻ ions.
The lattice also explains why ionic attractions should not be described as a single bond between one chosen cation and one chosen anion. Electrostatic attractions extend throughout the structure and act in all directions.
What determines the strength of ionic attraction?
The attraction becomes stronger when:
- the ions have larger charges
- the distance between their charged centers is smaller
This follows from electrostatic principles. Mg²⁺ and O²⁻ have larger charges than Na⁺ and Cl⁻, so magnesium oxide generally has stronger lattice attractions and a higher melting point than sodium chloride, although ionic radius must also be considered.
IB questions may ask students to compare compounds rather than recite a general property. A strong comparison names both relevant variables: ionic charge and ionic radius.
How does a covalent bond form?
A covalent bond forms when two atoms share a pair of electrons. More precisely, the bond results from the electrostatic attraction between the shared negatively charged electron density and the positively charged nuclei of both bonded atoms.
In H₂, each hydrogen atom contributes one electron:
H· + ·H → H:H
The shared pair can also be represented by a line, H to H. Each hydrogen then has access to two electrons in its valence shell, corresponding to the stable arrangement of helium.
Covalent bonds usually form between non-metals because neither atom readily loses enough electrons to produce a favorable ionic arrangement. Sharing allows both atoms to achieve a lower-energy electron distribution.
Single, double, and triple covalent bonds
Atoms can share more than one pair of electrons:
- Single bond: one shared pair, as in H to H
- Double bond: two shared pairs, as in O=O
- Triple bond: three shared pairs, as in N≡N
Between the same two elements, increasing bond order generally makes a bond shorter and stronger. This is useful when interpreting bond-length or bond-enthalpy data.
A coordinate covalent bond is formed when both electrons in the shared pair initially come from the same atom. Once formed, it behaves as a covalent bond. The ammonium ion, NH₄⁺, is a familiar example, formed when the lone pair on NH₃ is donated to H⁺.
For a fuller treatment of Lewis structures, multiple bonds, and polarity, consult Covalent Bonds Explained Clearly for IB Chemistry.
Why ionic and covalent substances have different properties
Bond type alone is not enough to predict a property. You must first identify the overall structure and then determine what must be separated or overcome during the physical process.
Melting and boiling points
Ionic compounds commonly have high melting and boiling points because melting requires enough energy to overcome strong electrostatic attractions between oppositely charged ions throughout the lattice. The ions become mobile in the liquid, but they remain ions.
Simple molecular covalent substances often have much lower melting and boiling points. When a substance such as methane melts or boils, its covalent C to H bonds are not broken. Energy overcomes the intermolecular forces between molecules.
This is one of the most important IB distinctions:
- Within a molecule: strong covalent bonds
- Between molecules: comparatively weaker intermolecular forces
Giant covalent substances behave differently. Diamond and silicon dioxide contain extensive networks of strong covalent bonds. Melting them requires many covalent bonds to be overcome, so their melting points are very high.
Therefore, the statement “ionic compounds have higher melting points than covalent compounds” is too broad. It is usually valid when comparing ionic compounds with simple molecular covalent substances, but not with every covalent material.
Electrical conductivity
Electrical conduction requires mobile charged particles.
In a solid ionic lattice, the ions are charged but fixed in position, so the solid does not conduct electricity. When the compound melts, the ions become mobile and can carry charge. An aqueous solution can also conduct if the ionic compound dissolves and produces mobile ions.
Most simple covalent substances do not conduct because their molecules are neutral and they lack mobile ions or delocalized electrons. However, some covalent substances react with water to form ions. Hydrogen chloride is molecular, for example, but its aqueous solution conducts because HCl ionizes in water.
Graphite is another important exception. It has a giant covalent structure, but each carbon contributes electrons that are delocalized across a layer. These mobile electrons allow graphite to conduct electricity parallel to its layers.
| Substance and state | Conductivity | Structural explanation |
|---|---|---|
| Solid NaCl | Does not conduct | Ions are fixed in the lattice |
| Molten NaCl | Conducts | Ions are mobile |
| NaCl(aq) | Conducts | Hydrated ions can move through the solution |
| Pure liquid water | Very poor conductor | Very low concentration of mobile ions |
| Aqueous HCl | Conducts | Molecular HCl forms ions in water |
| Graphite | Conducts | Delocalized electrons move along the layers |
| Diamond | Does not conduct | Electrons are localized in covalent bonds |
Hardness and brittleness
Ionic solids are often hard because the lattice contains strong attractions. They are also brittle. When a force shifts one layer of ions, ions of the same charge may become aligned; their repulsion causes the crystal to split.
Simple molecular solids are often soft because only intermolecular forces hold the molecules together. Giant covalent structures such as diamond can be exceptionally hard because strong covalent bonds extend in three dimensions.
Graphite shows why structure must be described carefully. Strong covalent bonds exist within each layer, but only weaker attractions act between layers, allowing the layers to slide over one another.
Solubility
The rule “ionic substances dissolve in water and covalent substances do not” is unreliable. Solubility depends on whether the new attractions formed between solute and solvent compensate for the attractions disrupted in both substances.
Water is polar and can form ion-dipole attractions with ions, so many ionic compounds dissolve in it. However, some ionic compounds are only slightly soluble because their lattice attractions are too strong relative to hydration effects.
Polar molecular substances may dissolve in water through dipole attractions or hydrogen bonding. Non-polar molecules usually dissolve more readily in non-polar solvents because their intermolecular interactions are more compatible. In an exam, use polarity and intermolecular forces rather than treating bond type as an infallible solubility rule.
Electronegativity and the ionic-covalent continuum
Electronegativity is an atom’s ability to attract a shared pair of electrons in a bond. When two identical atoms bond, the electrons are shared equally and the bond is non-polar covalent. When different atoms attract the shared pair unequally, the result is a polar covalent bond with partial charges, δ⁺ and δ⁻.
As the electronegativity difference increases, ionic character generally increases. However, no universal numerical boundary converts a covalent bond instantly into an ionic one. Threshold values found in textbooks are useful guidelines, not fundamental laws.
This is why the IB course uses models and treats bonding as a continuum. A substance can show substantial ionic and covalent character. Small, highly charged cations can distort the electron cloud of nearby anions, increasing covalent character in an otherwise ionic model.
The RevisionDojo notes on bonding as a continuum introduce the van Arkel-Ketelaar bonding triangle used to compare ionic, covalent, and metallic character. The related position in the bonding triangle resource is useful when a question provides electronegativity data.
For polarity questions, remember that bond polarity and molecular polarity are not identical. A molecule may contain polar covalent bonds but have no overall dipole if its geometry causes the bond dipoles to cancel. Carbon dioxide is the standard example: each C=O bond is polar, but linear CO₂ is non-polar overall.
How to identify ionic and covalent bonding in IB questions
Begin with the elements and particles present, but treat this as an initial classification rather than proof.
- Identify the element types. A metal combined with a non-metal is usually modeled as ionic; non-metals bonded together are usually covalent.
- Look for ions or charges. A lattice of cations and anions indicates ionic bonding.
- Determine the structure. Decide whether the substance is ionic, simple molecular, or giant covalent.
- Use property data. Melting point, conductivity, solubility, and physical state can support or challenge the proposed model.
- Account for exceptions. Graphite, aqueous molecular acids, and compounds with mixed bonding character require more careful reasoning.
Compounds containing polyatomic ions involve both bond types. In ammonium nitrate, NH₄NO₃, covalent bonds hold atoms together within NH₄⁺ and NO₃⁻, while ionic attractions act between the oppositely charged polyatomic ions. Calling the entire substance simply “covalent” because it contains only non-metals would miss its ionic lattice.
How to write an exam-quality explanation
IB questions often use command terms such as explain, compare, or deduce. An explanation should provide a causal sequence rather than a list of memorized properties.
For example, to explain why solid sodium chloride has a high melting point:
Sodium chloride has a giant ionic lattice containing oppositely charged ions. Strong electrostatic attractions act between the ions throughout the lattice. A large amount of energy is required to overcome these attractions, so sodium chloride has a high melting point.
To explain why methane has a low boiling point:
Methane is a simple molecular substance. Although the covalent bonds within each molecule are strong, only weak London dispersion forces act between its molecules. Little energy is required to overcome these intermolecular forces, so methane has a low boiling point.
Avoid these common errors:
- saying ionic bonding is merely “the transfer of electrons”
- referring to NaCl molecules instead of a lattice or formula units
- claiming that covalent bonds break when a molecular substance boils
- saying solid ionic compounds conduct because they contain charged particles
- assuming every covalent substance has a low melting point
- treating an electronegativity cutoff as an absolute boundary
- confusing a polar bond with a polar molecule
After reviewing the concepts, use the Structure 2 bonding questionbank to practise explanations based on unfamiliar evidence. The Structure 2 flashcards are more appropriate for definitions and rapid recall. Jojo AI can then help identify whether a weak answer is missing terminology, structural reasoning, or a complete cause-and-effect chain.
Conclusion
The central ionic vs covalent bond distinction is that ionic bonding attracts oppositely charged ions, whereas covalent bonding attracts two nuclei to shared electron density. Ionic substances normally form extended lattices; covalent substances may form small molecules or giant networks, leading to a much wider range of properties.
For IB Chemistry, do not stop after naming the bond. Identify the structure, state what particles or electrons are mobile, distinguish bonds from intermolecular forces, and explain the energy required to overcome the relevant attractions. RevisionDojo’s Structure notes, Flashcards, Questionbank, and Jojo AI are most useful when combined in that order: establish the model, recall the terminology, and then practise applying it to exam evidence.
Sources and referenced URLs
- IB Diploma Programme Chemistry subject brief, first assessment 2025
- IB Chemistry curriculum updates
- OpenStax Chemistry 2e: Ionic Bonding
- OpenStax Chemistry 2e: Ionic and Molecular Compounds
- OpenStax Chemistry 2e: Chemical Bonding Summary
- RevisionDojo: IB Chemistry Structure Explained
- RevisionDojo: Covalent Bonds Explained Clearly
- RevisionDojo: Bonding as a Continuum Notes
- RevisionDojo: Position in the Bonding Triangle
- RevisionDojo: Structure 2 Bonding Questionbank
- RevisionDojo: Structure 2 Bonding Flashcards
