It's tempting to want to say exactly how much of an electronegativity difference is needed before a bond becomes ionic. That would let us decide whether a bond is ionic or covalent. Putting a hard number on that distinction does not work reliably, though, so we need to be cautious. Many textbook use the number 0.4 as a threshold electronegativity difference for a polar covalent bond. The threshold many textbooks use for an ionic bond is an electronegativity difference of 1.8. Potassium bromide, KBr, has a potassium with electronegativity 0.82 and bromine with electronegativity 2.96. The difference is 2.16, and we definitely think of potassium bromide as an ionic solid containing arrays of bromide ions with potassium ions in between them.

For comparison, in an O-H bond, oxygen has an electronegativity 3.44 and hydrogen has an electronegativity hydrogen has an electronegativity 2.20. The difference is 1.24, which is well above the threshold for a polar covalent bond but still below the threshold for a purely ionic bond, in which oxygen would always have a negative charge and hydrogen would always have a positive charge.

In contrast, a C-H bond has a carbon with electronegativity 2.55, so the difference is only 0.35. Although the carbon is more electronegative than the hydrogen, the difference isn't big enough to make an appreciable dipole.

But there are some compounds in which there are large electronegativity differences but the bonds do not seem to be ionic. Silicon tetrafluoride, SiF₄, is an example. SiF₄ is a gas at room temperature. A gas is usually a molecular compound with covalent bonds. The Si-F bond has a silicon with electronegativity 1.90 and a fluorine with electronegativity 3.98. The difference is 2.08, which seems to clear the threshold for ionic bonds, which is only 1.8. There are lots of other examples, too. Transition metal oxides have a significant amount of covalency in their bonds despite their electronegativity differences. There are examples such as potassium permanganate, KMnO₄, in which MnO₄^- is a molecular ion with covalent Mn-O bonds, but Mn only has electronegativity 1.55, making an electronegativity difference with oxygen of 1.89. That's above the threshold for an ionic bond.

Some sources get around this problem by using a higher threshold for an ionic bond, but that doesn't explain some compounds like sodium iodide, NaI. That's a water-soluble, ionic compound. But sodium has electronegativity 0.93 compared to iodine with electronegativity 2.66. The difference is only 1.73. In this case, that seems to small to be ionic.

The point is that while students want simple tests to make sure they are getting the right answer and instructors want simple tests so their students don't get confused, those tests may not always work in reality. It's still OK to use those kinds of parameters in the classroom, because everyone is still learning and sometimes we need to simplify things a little. Usually the reason these rules sometimes fail is that there are additional factors we haven't considered. One of the factors here is ion stability, whih is something you will learn about if you start looking at structural factors in acid-base chemistry. Other factors can include the spatial and energetic overlap between the atomic orbitals that are coming together to form bonds between the atoms, but those are more sophisticated concepts that are often assessed using computational chemistry.

Problem IM7.5.1.

i) Use the following criteria to predict whether the bond is nonpolar covalent, polar covalent, or ionic.

ii) Compare your prediction with the experimental properties of the compound.

• electronegativity difference below 0.4: nonpolar covalent
• electronegativity difference between 0.4 and 1.8: polar covalent
• electronegativity difference above 1.8: ionic

a) BH₃, a gas at room temperature.

b) KCl, a water-soluble solid at room temperature.

c) BF₃, a gas at room temperature.

d) LiBr, a water-soluble solid at room temperature.

e) CaF₂, a water-insoluble, crystalline solid at room temperature that melts around 1400°C.

f) SnCl₄, a water-insoluble, crystalline solid at room temperature that melts around 240°C.

g) SnF₂, a water-soluble, crystalline solid at room temperature that melts around 200°C and boils around 850°C.

h) InF₃, a crystalline solid at room temperature that melts around 1200°C.

i) InCl₃, a water-soluble, crystalline solid at room temperature that melts around 500°C and boils around 800°C.

This site was written by Chris P. Schaller, Ph.D., College of Saint Benedict / Saint John's University (retired) with contributions from other authors as noted.  It is freely available for educational use.

Structure & Reactivity in Organic, Biological and Inorganic Chemistry by Chris Schaller is licensed under a Creative Commons Attribution-NonCommercial 3.0 Unported License. 

Send corrections to cschaller@csbsju.edu

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