Ionic vs Covalent Bond Calculator

Compare two elements on the Pauling scale to see whether the bond between them is ionic, polar covalent or nonpolar covalent.

What it does

Why does table salt melt at 801 degrees C while dry ice sublimates at minus 78? The answer starts with a single number for each element, and that number is electronegativity. It measures how hard an atom pulls on the electrons in a bond. OpenStax Chemistry describes it as the tendency of an atom to attract electrons towards itself, and notes that fluorine, at the top right of the periodic table, is the most electronegative element at 4.0 while the metals on the left are the weakest.

Electronegativity only matters when two atoms are bonded, because then they have to share, or fight over, the same pair of electrons. The bigger the gap between the two atoms, the more lopsided that sharing becomes. Pauling built the scale by comparing the energies needed to break different bonds, and the values are relative rather than absolute, which is why they are quoted without a unit.

The gap between the two atoms then tells you what kind of bond you have. OpenStax puts the rough bands at a difference of about 0.4 or less for a nonpolar covalent bond, between 0.4 and 1.7 for a polar covalent bond, and 1.7 or more for an ionic bond. The chemistry teaching notes at Ohio State give the same three bands, and add the useful reminder that the scale is a guide with many exceptions: a difference of 1.9 in hydrogen fluoride still gives a polar covalent bond, and a difference of 1.0 in manganese iodide still gives an ionic compound.

That last point is worth repeating, because the position of the elements in the periodic table often tells you more than the number. A bond between two nonmetals is usually covalent. A bond between a metal and a nonmetal is usually ionic. The calculator here gives the number and the band, and the note underneath it says plainly that the number is a guide.

The second thing the page tells you is which atom takes the negative charge. In every polar bond the more electronegative atom carries the delta minus and the other carries the delta plus, which is how a molecule like water gets its dipole and how the oxygen in water pulls the hydrogen the way it does. For an ionic bond the same rule applies in the strongest possible form, since the electron moves across completely.

How to use

1. Pick the first element from the list. Common metals, nonmetals and halogens are all there.
2. Pick the second element the same way. You can also choose the same element twice to see a bond with no difference at all.
3. Read the electronegativity difference from the first result.
4. Read the bond type. Nonpolar covalent means the electrons are shared evenly, polar covalent means they are pulled to one side, and ionic means they belong to one atom.
5. Check which atom takes the negative charge, since that is the more electronegative one.
6. Use the bands as a guide and look at the elements involved as well, because a metal and a nonmetal usually form an ionic compound even when the numbers sit low.

Formula

EN

Pauling electronegativity of an element, a relative number with no unit.

Difference

Absolute gap between the two electronegativities.

delta minus

The partial negative charge that sits on the more electronegative atom.

delta plus

The partial positive charge on the less electronegative atom.

The cut-offs of 0.4 and 1.7 are rules of thumb taken from the standard chemistry texts, not sharp physical boundaries. Bonding runs on a continuum from fully covalent to fully ionic, and the same pair of elements can sit in different places depending on the compound. Noble gases are left out because they hold a full shell and rarely share electrons in the way this scale describes.

Worked example

Reference table

Limits

• The scale describes ordinary bonds between two atoms and says nothing about metals, network solids or bonding in molecules with three atoms or more.
• The cut-offs are a teaching guide rather than a physical boundary, and the same pair of elements can bond differently in different compounds.
• Electronegativity depends on oxidation state and on the compound, so a single value per element is an average.
• It does not predict bond strength, melting point or whether a compound will conduct electricity, it only describes how the electrons are shared.

Common mistakes

• Forgetting the absolute value. The difference is a gap, so a negative result from subtracting in the wrong order has to be turned positive.
• Treating 0.4 and 1.7 as exact boundaries. The chemistry texts call them rough approximations with plenty of exceptions.
• Assuming a large difference always means an ionic compound. Manganese iodide has a difference of only 1.0 and is still ionic.
• Ignoring where the elements sit. Two nonmetals bond covalently and a metal with a nonmetal bonds ionically, whatever the arithmetic says.
• Reading electronegativity as electron affinity. They are different quantities, and only electronegativity describes sharing in a bond.

Try it yourself

Quick checks. Answers sit tucked under each question.

Questions

Sources

• OpenStax Chemistry 2e, 7.2: covalent bonding, electronegativity and bond type
• Ohio State University: the chemical bond and the electronegativity difference bands
• University of Calgary chemistry text: electronegativity and bond type

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