Chemical bonds hold atoms together to form everything around us — from the salt on your chips to the oxygen you're breathing right now. Today we'll look at the three main types of bonding and learn a simple trick for predicting which type you'll get, using electronegativity.
Why Do Atoms Bond at All?
Atoms bond because joining together usually gives them a lower, more stable energy state — often by achieving a full outer shell (the octet rule). The way they achieve this depends on how much the bonding atoms "want" electrons, which is measured by electronegativity.
The Three Main Types of Bonding
- Ionic bonding: electrons are transferred from one atom to another, forming oppositely charged ions that attract each other electrostatically. Favoured when one atom has low electronegativity (a metal) and the other has high electronegativity (a non-metal) (Talbot, p. 208).
- Covalent bonding: atoms share electron pairs. Favoured when both atoms have high electronegativity values, typically two non-metals (Talbot, p. 208).
- Metallic bonding: positive metal ions sit in a "sea" of delocalised electrons. Favoured when atoms have low electronegativity values, as in metals and alloys like brass or bronze (Talbot, p. 208).
Using Electronegativity Differences to Predict Bond Type
Electronegativity measures how strongly an atom attracts shared electrons in a bond. By comparing electronegativity values (Pauling scale) of two bonded atoms, we can predict what kind of bond forms (Talbot, p. 134):
- Difference > 1.7 → likely ionic
- Difference between 0.5 and 1.7 → polar covalent
- Difference < 0.5 → non-polar covalent
Important: these are extremes on a continuous scale, not strict boxes. A polar covalent bond is really a covalent bond with some ionic character — the bigger the electronegativity gap, the more "ionic" the shared electron pair becomes (Talbot, p. 134).
Worked Example 1: Cl₂
Electronegativity of Cl = 3.2. Difference = 3.2 − 3.2 = 0. → Non-polar covalent bond (Talbot, p. 135).
Worked Example 2: HI
Electronegativity: H = 2.2, I = 2.7. Difference = 0.5. → Right at the borderline, classified as polar covalent (Talbot, p. 135).
Worked Example 3: LiF
Electronegativity: Li = 1.0, F = 4.0. Difference = 3.0. → Well above 1.7, so this bond is strongly ionic (Talbot, p. 135).
A Deeper Twist: Charge-Shift Bonds
Interestingly, even "purely covalent" bonds like F–F aren't so simple. Chemists Sason Shaik and Philippe Hiberty used computer modelling (building on Linus Pauling's valence bond theory) and found that F₂'s bond energy is only correctly predicted if you mix in some ionic character (F⁺F⁻) alongside the covalent form. They called this new category a charge-shift bond (Talbot, p. 135). This shows that bonding models are useful simplifications, not rigid boxes — real bonds often blend characteristics.
A Note on Lewis Theory
Gilbert Lewis introduced electron-dot diagrams in 1902 as a teaching tool, and they're still brilliant for predicting valency, octets, and simple molecular shapes (especially combined with VSEPR theory later). However, Lewis theory is a pre-quantum, non-mathematical model — it can't explain why bonding happens at a deeper physical level or account for spectroscopic data (Talbot, p. 186). It's a useful first model, not the final word.
Key Takeaways
- Bonding type depends on electronegativity: ionic (Δ > 1.7), polar covalent (0.5–1.7), non-polar covalent (< 0.5).
- Ionic, covalent, and metallic bonding are extremes on a spectrum — polar bonds sit in between.
- Charge-shift bonds (like in F₂) show that even "pure" covalent bonds can have hidden ionic character.
- Lewis electron-dot diagrams are a simple, powerful tool but don't explain the deeper physics of bonding.