How to Actually Tell Whether a Bond Is Ionic or Covalent
The textbook answer is simple: metals give electrons to nonmetals, nonmetals share with each other. That works in practice maybe three-quarters of the time. The remaining quarter is where people mess up, and it's usually because they're treating electronegativity as a hard line instead of a spectrum. I spent years running elemental analyses and interpreting XRD patterns for industrial clients. One of the recurring problems I ran into was with compounds like aluminum chloride (AlCl3). Electronegativity difference says it should be ionic. Aluminum is a metal, chlorine is a nonmetal, the math checks out. But AlCl3 sublimes at 180°C, dissolves in organic solvents, and behaves like a covalent molecular compound in many reactions. It's not a edge case you find in a textbook footnote — it came up regularly when someone was troubleshooting an unexpected reaction pathway. The workaround I ended up using wasn't fancy. I'd look at the physical properties first: melting point, solubility behavior, and how it reacted in solution. If it dissolved in benzene or hexane instead of water and the solution didn't conduct electricity, that was a strong signal the bonding wasn't purely ionic regardless of what the electronegativity table said. Electronegativity is a useful starting point, but it's not a verdict. Polarizability and Fajans' rules matter just as much, and most people skip over them entirely.
Understanding Ionic And Covalent Compounds Beyond the Basics
Here's what actually happens when you're working with these compounds in a lab or production setting, and the things that trip people up: Covalent compounds involve electron sharing between atoms. The sharing can be equal, which gives you nonpolar covalent bonds like O2 or N2, or unequal, which gives you polar covalent bonds like HCl or water. When the electronegativity difference falls between about 0.4 and 1.7 on the Pauling scale, you're usually in polar covalent territory. Below 0.4, it's essentially nonpolar. Above 1.7, you're trending toward ionic. Those numbers are rough guides, not laws of nature. Ionic compounds form when one atom completely transfers electron(s) to another, creating cations and anions held together by electrostatic forces. Sodium chloride is the classic example, but don't let that simplicity fool you. Ionic solids form crystal lattices, not discrete molecules. That means when we talk about the "formula" of an ionic compound, we're really talking about the simplest whole-number ratio of ions in the lattice, not a molecule you can isolate.
A counter-intuitive point that nobody emphasizes enough: not all ionic compounds conduct electricity in their solid state. They only conduct when molten or dissolved in water because the ions need to be free to move. A block of NaCl sitting on your bench is an electrical insulator. This seems obvious when someone tells you, but I've seen it cause genuine confusion on the floor when someone tested a solid salt sample and assumed the bonding wasn't ionic because the multimeter showed no conductivity. Another thing beginners consistently miss: coordinate covalent bonds. These happen when one atom contributes both electrons to a shared pair. Ammonium ion (NH4+) is the textbook example — nitrogen shares its lone pair with a hydrogen proton. The resulting bond is indistinguishable from any other covalent bond. You can't point to one N-H bond and say "this one is coordinate." Once it forms, all four bonds are identical. This matters when you're predicting molecular geometry or calculating formal charges, and getting it wrong leads to bad Lewis structures that cascade into errors downstream. Let me give you a practical procedure for determining bond type when you're given an unknown compound:
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First, identify the elements and their positions on the periodic table. Metal plus nonmetal usually means ionic. Nonmetal plus nonmetal usually means covalent. That gets you somewhere fast, but it won't get you everywhere. Second, calculate the electronegativity difference using Pauling values. If the difference is greater than 1.7, lean ionic. Less than 0.4, lean covalent nonpolar. Between those numbers, it's polar covalent. Again, useful but not definitive. Third, check the physical properties if you have access to the actual substance. Melting point above 600°C and high solubility in water with electrical conductivity in solution points strongly ionic. Low melting point, solubility in organic solvents, and no conductivity in solution points covalent. This third step is the one that catches the AlCl3 situations I mentioned earlier.
Common pitfalls to avoid: The biggest one is assuming that polyatomic ions make a compound covalent. Na2SO4 contains sulfate, which is held together by covalent bonds, but the compound as a whole is ionic because sodium ions bond ionically to sulfate ions. The compound has both types of bonding simultaneously. This isn't a rare situation — most salts of polyatomic ions work this way, and mixing it up will throw off your predictions for solubility, conductivity, and reaction behavior. Another pitfall is the belief that covalent compounds never dissolve in water. Sugar dissolves in water readily despite being covalent, because it can form hydrogen bonds with water molecules. The "like dissolves like" rule is more nuanced than the simplified version you hear in introductory classes. Polar covalent molecules can be very water-soluble if they have the right functional groups.
When I'm working with network covalent solids like diamond, quartz (SiO2), or silicon carbide, the classification gets even messier. SiO2 is technically a covalent compound, but it forms a three-dimensional network with a melting point around 1710°C — higher than many compounds I'd call unambiguously ionic. Calling it simply "covalent" is true but practically misleading if you're trying to predict how it will behave under thermal stress or in a chemical process. The takeaway I wish more people understood: bonding exists on a continuum, and most real compounds fall somewhere in the middle rather than sitting neatly in one category or the other. The ionic-covalent distinction is a model, not a fundamental law of nature. It's useful, but it has boundaries. When those boundaries matter — and they do more often than introductory courses let on — you need to bring in polarization theory, lattice energy calculations, and actual experimental data rather than relying on a memorized electronegativity cutoff.
