Chemical Bonding Misconceptions (Including "Breaking Bonds Releases Energy")

Chemical bonding is taught early and built on constantly, so misconceptions picked up at the start can linger for years — sometimes all the way to university. Many come from simplified models (“atoms want a full shell”), everyday language (“energy stored in bonds”), or diagrams taken too literally. This article tackles ten of the most common and most damaging, explaining what’s wrong and what’s actually true.

Misconception 1: “Breaking bonds releases energy”

This is probably the most widespread bonding misconception of all. It often comes from phrases like “ATP releases energy when its bond breaks” or “food stores energy in its bonds”.

The truth: breaking a bond always requires energy. Forming a bond always releases energy. Think of two magnets stuck together: you have to put in effort to pull them apart; they release energy (as a click and a little heat) when they snap together.

When a reaction releases energy overall — burning methane, respiration, an explosion — it’s because the bonds formed in the products are stronger overall than the bonds broken in the reactants:

ΔH = energy to break bonds − energy released forming bonds

For methane burning, about 2,650 kJ is needed to break the bonds in CH₄ and 2O₂, but about 3,460 kJ is released forming the bonds in CO₂ and 2H₂O. The difference is released as heat (see bond enthalpy calculations).

The same applies to ATP: energy is released overall when ATP is hydrolysed because the products are more stable, not because breaking a phosphate bond gives out energy (see ATP vs ADP).

Misconception 2: “Atoms bond because they want a full outer shell”

The truth: atoms don’t want anything. They bond because the bonded arrangement has lower energy than the separate atoms (see what is a chemical bond?).

The “full shell” or octet rule is a useful pattern for predicting formulas in periods 2 and 3, but it isn’t the cause. There are plenty of exceptions: BF₃ has only six electrons around boron, SF₆ has twelve around sulfur, and transition metals rarely follow it. Worse, the “wanting” language leads to the next misconception.

Misconception 3: “Sodium gives its electron to chlorine because it wants a full shell”

The truth: removing an electron from a sodium atom costs energy (496 kJ mol⁻¹), and adding it to chlorine releases only 349 kJ mol⁻¹. On their own, those two steps are uphill by 147 kJ mol⁻¹. Isolated sodium atoms don’t spontaneously hand electrons to chlorine atoms.

Sodium chloride forms because of the lattice energy released when the ions come together into a crystal — about 787 kJ mol⁻¹, far more than enough to pay for the electron transfer (see Born–Haber cycles and lattice energy).

Misconception 4: “Salt is made of NaCl molecules”

The truth: sodium chloride is a giant ionic lattice. Each Na⁺ is surrounded by six Cl⁻ ions, and each Cl⁻ by six Na⁺ ions. There’s no particular Cl⁻ that “belongs” to a particular Na⁺. The formula NaCl simply gives the ratio of ions (see ionic lattices).

The same applies to giant covalent substances: SiO₂ is a ratio, not a molecule.

Misconception 5: “An ionic bond is a bond between one pair of ions”

The truth: ionic bonding is non-directional. Each ion attracts all the oppositely charged ions around it (and repels same-charge ions further away). The “ionic bond” is really the sum of all these attractions throughout the lattice. That’s why ionic compounds form 3D lattices, not pairs.

Misconception 6: “When water boils, the bonds in water break”

The truth: when a simple molecular substance melts or boils, the covalent bonds inside the molecules stay intact. What’s overcome are the much weaker intermolecular forces between molecules — in water, hydrogen bonds (see intermolecular forces).

Steam is still H₂O molecules. If boiling broke the O–H bonds, steam would be hydrogen and oxygen — an explosive mixture! The energy numbers make the difference clear: about 41 kJ mol⁻¹ to boil water at 100 °C, versus about 926 kJ mol⁻¹ to break both O–H bonds in a mole of water molecules.

Misconception 7: “Covalent bonds are weak because covalent substances have low melting points”

The truth: individual covalent bonds are strong — often stronger than the attraction between a single pair of ions. Simple molecular substances have low melting points because the forces between molecules are weak, not the bonds within them (see properties of covalent compounds).

Proof: diamond and silicon dioxide, where covalent bonds link every atom, have some of the highest melting points known (see giant covalent structures).

Misconception 8: “A hydrogen bond is a bond to a hydrogen atom within a molecule”

The truth: a hydrogen bond is an intermolecular attraction (or sometimes an attraction between distant parts of a large molecule). It forms between a δ+ hydrogen atom already covalently bonded to N, O or F, and a lone pair on another N, O or F atom. The O–H bond inside a water molecule is a covalent bond; the attraction between that H and the O of a neighbouring molecule is the hydrogen bond (see hydrogen bonding).

Hydrogen bonds are about 10–40 kJ mol⁻¹ — around a twentieth of a typical covalent bond.

Misconception 9: “Solid ionic compounds conduct electricity because they contain ions”

The truth: conduction needs mobile charged particles. In a solid ionic lattice, the ions are locked in place, so solid salt is an insulator. When the compound is molten or dissolved, the ions can move and it conducts (see properties of ionic compounds).

A related error: “metals conduct because of their positive ions”. In metals, it’s the delocalised electrons that move; the positive ions stay in the lattice (see how metallic bonding explains conductivity).

Misconception 10: “Bonding is either ionic or covalent — nothing in between”

The truth: bonding is a continuous spectrum. Two identical atoms share electrons equally (pure covalent). As the electronegativity difference grows, sharing becomes increasingly unequal (polar covalent, as in H–Cl and O–H), until electrons are effectively transferred (ionic, as in NaCl and CsF). Even NaCl has a little covalent character, and some metal–non-metal compounds, such as aluminium chloride, are largely covalent (see bond polarity and ionic vs covalent bonds).

Why these misconceptions matter

Each one blocks later understanding:

• Believing bond breaking releases energy makes energetics, Hess’s law and biochemistry confusing.
• The “wanting a full shell” story hides the energy reasons behind bonding and makes transition metals and expanded octets seem mysterious.
• Mixing up bonds and intermolecular forces makes physical properties impossible to explain correctly.

How to check your own thinking

Ask yourself:

1. Where is the energy? Is energy being put in (breaking) or released (forming)?
2. What particles are there? Ions, molecules, atoms in a network, or metal ions with delocalised electrons?
3. What is actually overcome? Strong bonds or weak intermolecular forces?
4. Can the charged particles move?

Key takeaways

• Breaking bonds needs energy; forming bonds releases it. Exothermic reactions form stronger bonds than they break.
• Atoms bond because it lowers energy, not because they “want” full shells.
• Ionic compounds are lattices, not molecules; their formation is driven by lattice energy.
• Boiling simple molecular substances overcomes intermolecular forces, not covalent bonds.
• Bonding is a spectrum from covalent to ionic.

For a full revision checklist, see the structure and bonding revision guide.