How to Predict Whether a Bond Is Ionic or Covalent from Electronegativity

Ammonium chloride, \(\ce{NH4Cl}\), is an ionic compound. Yet the nitrogen and hydrogen atoms inside each ammonium ion are joined by covalent bonds. So how can the same substance contain both types of bonding?

The answer is that you need to distinguish between a bond between two atoms and the attractions holding a whole compound together.

Predict before reading on: if the electronegativity difference between nitrogen and hydrogen is relatively small, does that mean the entire compound \(\ce{NH4Cl}\) is covalent?

No. The N-H bonds are polar covalent, while the attraction between \(\ce{NH4+}\) and \(\mathrm{Cl}^{-}\) ions is ionic.

This distinction is central to predicting bond type from electronegativity. Calculate the difference for the atoms actually bonded, then use valence electrons, charges, structural representations, and physical evidence to check your prediction.

01Predict a bond using electronegativity difference

Electronegativity, represented by \(\chi\), describes how strongly an atom attracts the shared electrons in a chemical bond.

If two atoms have similar electronegativities, they attract the bonding electrons by similar amounts. The electrons are shared relatively evenly.

If one atom is considerably more electronegative, it attracts the bonding electrons more strongly. The electron distribution becomes uneven.

The quantity used to compare these attractions is the electronegativity difference:

\[ \boxed{\Delta\chi=|\chi_A-\chi_B|} \]

Where:

• \(\Delta\chi\) is the electronegativity difference.
• \(\chi_A\) is the electronegativity of atom A.
• \(\chi_B\) is the electronegativity of atom B.
• The absolute value signs ensure that the difference is positive.

Electronegativity values are dimensionless, so \(\Delta\chi\) has no units.

What does the difference tell you?

A commonly used HSC-level classification guide is:

These boundaries are guidelines, not universal laws. Different reference tables may use slightly different values.

The important relationship is:

\[ \text{Increasing }\Delta\chi \quad\Longrightarrow\quad \text{Increasing ionic character} \]

A larger difference means a more uneven distribution of bonding electrons.

Which atom becomes partially negative?

In a polar covalent bond, the more electronegative atom attracts the shared electrons more strongly.

It therefore becomes partially negative, represented by \(\delta-\).

The less electronegative atom becomes partially positive, represented by \(\delta+\).

For example, consider hydrogen chloride:

\[ \ce{H^{\delta+}-Cl^{\delta-}} \]

Chlorine attracts the bonding electrons more strongly than hydrogen.

Notice the distinction between partial charges and ionic charges:

• \(\delta+\) and \(\delta-\) represent partial charges in an unevenly shared electron distribution.
• \(+1\), \(-1\), \(+2\), and similar values represent the net charges of ions.

A polar covalent bond does not automatically mean that one atom has completely lost an electron.

02Where did the electronegativity scale come from?

Electronegativity is not something chemists measure using a device that directly reads how strongly an atom pulls on an electron.

The familiar Pauling electronegativity scale was developed using experimental evidence about bond energies.

Linus Pauling compared the energy required to break bonds between identical atoms with the energy required to break bonds between different atoms.

The relevant measurement is called bond dissociation energy, represented by \(D\). It is the energy required to break a specified bond in gas-phase molecules, usually expressed in \(\mathrm{kJ\,mol^{-1}}\).

Consider three bonds:

• A-A, between two identical A atoms.
• B-B, between two identical B atoms.
• A-B, between different atoms.

If A and B attract bonding electrons equally, their heteronuclear A-B bond would be expected to have a bond energy close to the average of the two homonuclear bond energies.

However, experiments showed that many A-B bonds were stronger than this average.

Pauling interpreted the additional bond strength as evidence of an additional ionic contribution to the bonding.

A small illustrative calculation

Suppose experimental measurements give the following hypothetical bond dissociation energies:

First, calculate the average of the two identical-atom bond energies:

\[ D_{\mathrm{average}}=\frac{250+150}{2} \]

\[ D_{\mathrm{average}}=200\ \mathrm{kJ\,mol^{-1}} \]

Now compare this with the A-B bond:

\[ \Delta D=230-200 \]

\[ \boxed{\Delta D=30\ \mathrm{kJ\,mol^{-1}}} \]

The A-B bond has an additional \(30\ \mathrm{kJ\,mol^{-1}}\) of bond strength compared with the average.

This additional stabilisation was the type of evidence Pauling used to construct relative electronegativity differences.

The calculation illustrates the historical reasoning. It does not directly give an electronegativity difference or prove complete electron transfer.

The Pauling scale remains widely used, although other scales have been developed using different atomic properties. The IUPAC Gold Book definition of electronegativity explains its relative nature and its connection to bond dissociation energies.

The important limitation is that electronegativity is a useful model for predicting electron distribution. It does not independently determine every structural feature of a compound.

03Read the chemical structure before classifying the compound

This is where many otherwise sensible answers go wrong.

Electronegativity difference describes the attraction between atoms involved in a particular bond. However, a chemical formula can represent a substance containing several different types of bonding.

You need to identify what the structural representation actually shows.

Covalent bonds and ionic attractions are represented differently

In a Lewis structure, a line between two atoms represents a shared pair of electrons.

For example:

\[ \ce{H-Cl} \]

This represents a covalent bond between hydrogen and chlorine.

An ionic compound is represented using ions and their charges:

\[ \ce{Na+ \quad Cl-} \]

The attraction between these ions is electrostatic. It is not a shared pair of electrons represented by a conventional covalent bond line.

In a solid ionic compound, the ions form an extended lattice rather than separate little molecules.

What happens inside a polyatomic ion?

A polyatomic ion is a group of covalently bonded atoms that carries an overall electrical charge.

Consider sodium hydroxide, \(\ce{NaOH}\).

It contains:

• A sodium ion, \(\mathrm{Na}^{+}\).
• A hydroxide ion, \(\ce{OH-}\).

Inside the hydroxide ion, oxygen and hydrogen are covalently bonded.

Between sodium and hydroxide, there is ionic attraction.

Therefore, sodium hydroxide contains both covalent bonding and ionic bonding.

This is not contradictory. The two descriptions apply to different interactions.

Valence electrons help check your prediction

Electronegativity tells you which atom attracts bonding electrons more strongly. Valence electrons help you work out whether the proposed ions or shared-electron arrangements make chemical sense.

For example:

• Sodium has one valence electron and commonly forms \(\mathrm{Na}^{+}\).
• Magnesium has two valence electrons and commonly forms \(\mathrm{Mg}^{2+}\).
• Oxygen has six valence electrons and commonly forms \(\mathrm{O}^{2-}\).
• Chlorine has seven valence electrons and commonly forms \(\mathrm{Cl}^{-}\).

These common ion charges help you check whether an ionic formula is electrically neutral.

For example, magnesium chloride contains:

\[ \ce{Mg^2+ + 2Cl-} \]

The total charge is:

\[ (+2)+2(-1)=0 \]

This gives the formula \(\ce{MgCl2}\).

Remember that a Lewis structure is a model of electron sharing and formal charges. It is not a literal photograph of where electrons sit.

The important exception: a large difference does not guarantee ionic bonding

Consider boron trifluoride, \(\ce{BF3}\).

Using the approximate Pauling electronegativities:

\[ \chi_{\ce{B}}=2.04 \]

\[ \chi_{\ce{F}}=3.98 \]

The difference is:

\[ \Delta\chi=|3.98-2.04|=1.94 \]

This exceeds the commonly used 1.7 guideline.

A simplistic classification would label the B-F bonds ionic.

However, \(\ce{BF3}\) is a molecular covalent compound. Boron shares electron pairs with three fluorine atoms, producing discrete trigonal-planar molecules.

The B-F bonds are strongly polar covalent, even though their electronegativity difference is large.

This is a useful warning: do not force every compound into an ionic model simply because its electronegativity difference crosses a particular numerical boundary.

04Use physical evidence to check the bonding model

Electronegativity gives you a prediction. Physical properties can provide additional evidence about whether that prediction makes sense.

Electrical conductivity is particularly useful because it tells you whether mobile charged particles are present.

How does an ionic compound conduct electricity?

In a solid ionic compound, positive and negative ions are held in fixed positions within a crystal lattice.

The ions are charged, but they cannot move freely through the solid.

Therefore, solid ionic compounds generally do not conduct electricity.

When the compound melts, the ions become mobile. They can move through the liquid and carry electrical charge.

This allows the molten ionic compound to conduct electricity.

Comparing ionic and covalent substances

These are general patterns, not universal rules.

For example, giant covalent structures can have very high melting points without being ionic. Graphite can also conduct electricity because of its delocalised electrons.

Conductivity alone is therefore not enough to identify bonding.

The important example: hydrogen chloride in water

Hydrogen chloride gas, \(\ce{HCl}\), consists of covalently bonded molecules.

Its electronegativity difference is:

\[ \Delta\chi=|3.16-2.20|=0.96 \]

This indicates polar covalent bonding.

However, when hydrogen chloride dissolves in water, it reacts to form ions:

\[ \ce{HCl + H2O -> H3O+ + Cl-} \]

The resulting solution conducts electricity because it contains mobile ions.

Does this mean the original H-Cl bond was ionic?

No. The H-Cl bond in the gas is polar covalent. The ions are produced when hydrogen chloride reacts with water.

This introduces an important distinction:

• Dissociation separates ions that already existed in an ionic substance.
• Ionisation produces ions from initially neutral molecules through a chemical process.

Both processes can produce electrically conducting solutions.

05Worked examples

06A reliable decision process

When you encounter an unfamiliar substance in an HSC Chemistry question, work through these steps.

1. Identify the actual bonded atoms. Do not assume every element in the formula is directly bonded to every other element.

2. Calculate the electronegativity difference. Use \(\Delta\chi=|\chi_A-\chi_B|\).

3. Predict the bonding character. A small difference suggests covalent bonding, while a large difference suggests increasing ionic character.

4. Check valence electrons and charges. Work out whether the proposed ions or shared-electron arrangement is chemically reasonable.

5. Identify the particles in the substance. Distinguish individual molecules, monatomic ions, polyatomic ions, and extended ionic lattices.

6. Use physical evidence. Consider melting behaviour and electrical conductivity, paying attention to whether the substance is solid, molten, or dissolved.

7. State the limitation. Do not claim that a particular electronegativity difference proves complete electron transfer.

A strong explanation connects the numerical difference to the actual electron distribution and the structure of the substance.

07Questions and solutions

Question 1

Sodium has a Pauling electronegativity of \(0.93\), while chlorine has a value of \(3.16\).

Calculate the electronegativity difference for a sodium-chlorine interaction. Predict the predominant bonding type and identify the likely ions formed.

Solution 1

Sodium chloride is predominantly ionic, consisting of \(\mathrm{Na}^{+}\) and \(\mathrm{Cl}^{-}\) ions.

First, calculate the electronegativity difference:

\[ \Delta\chi=|3.16-0.93| \]

\[ \boxed{\Delta\chi=2.23} \]

This is a large difference, indicating strong ionic character.

Sodium has one valence electron and commonly loses it to form \(\mathrm{Na}^{+}\).

Chlorine has seven valence electrons and gains one electron to form \(\mathrm{Cl}^{-}\).

The resulting ions have opposite charges and attract electrostatically.

Their charges balance:

\[ (+1)+(-1)=0 \]

Therefore, the simplest formula is \(\ce{NaCl}\).

The solid consists of an extended ionic lattice, not individual \(\ce{NaCl}\) molecules.

Question 2

Sodium hydroxide, \(\ce{NaOH}\), contains sodium, oxygen, and hydrogen.

Use the following Pauling electronegativities:

• Sodium: \(0.93\)
• Oxygen: \(3.44\)
• Hydrogen: \(2.20\)

A student claims that both the Na-O and O-H interactions must be ionic because sodium is a metal and oxygen is highly electronegative.

Explain why this representation is incorrect.

Calculate the electronegativity difference for O-H, identify the bonding within the hydroxide ion, and explain how the overall compound remains electrically neutral.

Solution 2

Sodium hydroxide contains ionic attraction between \(\mathrm{Na}^{+}\) and \(\ce{OH-}\), but the O-H bond inside the hydroxide ion is polar covalent.

First, calculate the O-H electronegativity difference:

\[ \Delta\chi=|3.44-2.20| \]

\[ \boxed{\Delta\chi=1.24} \]

This indicates polar covalent bonding.

Oxygen and hydrogen share an electron pair. Oxygen attracts the shared electrons more strongly, giving oxygen a partial negative charge and hydrogen a partial positive charge.

Now consider the hydroxide ion.

Oxygen has six valence electrons, and hydrogen contributes one. The negative charge adds another electron:

\[ 6+1+1=8 \]

These eight electrons form one shared O-H pair and three lone pairs on oxygen.

The hydroxide ion therefore has an overall charge of \(-1\).

Sodium loses one electron to form \(\mathrm{Na}^{+}\), which is attracted to the negatively charged hydroxide ion.

The overall charge is:

\[ (+1)+(-1)=0 \]

This gives neutral \(\ce{NaOH}\).

The mistake in the student’s explanation is treating the entire formula as though every interaction must have the same bonding type. The O-H interaction is covalent, while the attraction between sodium and hydroxide is ionic.

Question 3

Hydrogen chloride gas, \(\ce{HCl(g)}\), does not conduct electricity.

However, when hydrogen chloride dissolves in water, the resulting solution conducts electricity strongly.

The Pauling electronegativities of hydrogen and chlorine are \(2.20\) and \(3.16\), respectively.

Explain:

• The bonding type within hydrogen chloride molecules.
• Why the gas does not conduct electricity.
• Why its aqueous solution conducts electricity.
• The difference between ionisation and dissociation in this example.

Solution 3

Hydrogen chloride contains polar covalent H-Cl bonds. Its aqueous solution conducts because hydrogen chloride reacts with water to produce mobile ions.

First, calculate the electronegativity difference:

\[ \Delta\chi=|3.16-2.20| \]

\[ \boxed{\Delta\chi=0.96} \]

This indicates polar covalent bonding.

In hydrogen chloride gas, the atoms form neutral molecules. Although the H-Cl bond has an uneven electron distribution, the molecules do not contain freely moving ions.

Therefore, the gas does not conduct electricity.

When hydrogen chloride enters water, it transfers a proton to a water molecule:

\[ \ce{HCl + H2O -> H3O+ + Cl-} \]

This produces hydronium ions and chloride ions.

These charged particles can move through the solution and carry electrical charge.

The distinction is important:

• Ionisation occurs when initially neutral HCl molecules react with water to form ions.
• Dissociation occurs when an ionic substance separates into ions that were already present in its structure.

Therefore, the conductivity of hydrochloric acid does not prove that the original H-Cl bonds were ionic.

It shows that mobile ions are present after the reaction with water.

Question 4

Consider magnesium chloride, \(\ce{MgCl2}\), and boron trifluoride, \(\ce{BF3}\).

Use the following Pauling electronegativities:

You are given the following physical and structural observations:

• Magnesium chloride is a high-melting crystalline solid. It does not conduct electricity as a solid but conducts when molten.
• Boron trifluoride is a gas at room temperature. It consists of discrete trigonal-planar molecules, with each boron atom attached to three fluorine atoms.

A student calculates both electronegativity differences and concludes that both substances must be ionic because their differences exceed 1.7.

Evaluate this conclusion. Use the electronegativity calculations, valence electrons, structural information, and physical evidence in your explanation.

Solution 4

Magnesium chloride is predominantly ionic, whereas boron trifluoride is molecular covalent. Both have large electronegativity differences, showing why a numerical cut-off cannot independently determine the bonding type.

First, calculate the magnesium-chlorine difference:

\[ \Delta\chi_{\ce{Mg-Cl}}=|3.16-1.31| \]

\[ \boxed{\Delta\chi_{\ce{Mg-Cl}}=1.85} \]

Now calculate the boron-fluorine difference:

\[ \Delta\chi_{\ce{B-F}}=|3.98-2.04| \]

\[ \boxed{\Delta\chi_{\ce{B-F}}=1.94} \]

Both differences exceed the commonly used 1.7 guideline.

However, their structures are different.

Magnesium chloride

Magnesium has two valence electrons and commonly forms \(\mathrm{Mg}^{2+}\).

Chlorine gains one electron to form \(\mathrm{Cl}^{-}\).

The resulting formula is:

\[ \ce{Mg^2+ + 2Cl- -> MgCl2} \]

The ions form an extended ionic lattice.

The physical evidence supports this model. The solid does not conduct because its ions are fixed, while the molten compound conducts because its ions are mobile.

Boron trifluoride

Boron has three valence electrons, while fluorine has seven.

Boron forms three shared electron pairs with three fluorine atoms.

The resulting \(\ce{BF3}\) structure consists of discrete covalent molecules.

Each B-F bond is strongly polar because fluorine attracts the bonding electrons more strongly than boron.

However, the presence of polar covalent bonds does not mean that the substance forms a giant ionic lattice.

The gas-phase molecular structure supports a covalent model.

The student’s mistake is treating the 1.7 guideline as an absolute boundary. Electronegativity predicts the extent of unequal electron sharing, but the energy and stability of the complete structure also matter.

A better conclusion is that electronegativity difference provides evidence about bonding character, while valence, structural representations, and physical properties help establish the predominant bonding model.

Question 5

Two substances are investigated in a laboratory.

Substance P: Ammonium chloride, \(\ce{NH4Cl}\).

• It is a crystalline solid.
• It does not conduct electricity as a solid.
• Its aqueous solution conducts electricity.

Substance Q: Hydrogen chloride, \(\ce{HCl}\).

• It is a gas at room temperature.
• It does not conduct electricity as a gas.
• Its aqueous solution also conducts electricity.

A student argues:

“Both substances must be ionic because both aqueous solutions conduct electricity. If the solutions contain ions, the original substances must have contained ionic bonds.”

Use the following information:

Answer the following:

• Calculate the electronegativity differences for N-H and H-Cl.
• Identify the bonding inside the ammonium ion.
• Show how the valence electron count supports the structure of \(\ce{NH4+}\).
• Identify the bonding responsible for holding solid ammonium chloride together.
• Write the equations describing what happens when each substance enters water.
• Explain why the two solutions conduct electricity even though the original substances have different bonding structures.

Solution 5

The student’s conclusion is incorrect. Ammonium chloride is an ionic compound containing covalent N-H bonds, whereas hydrogen chloride is a molecular covalent substance. Both solutions conduct because they contain mobile ions, but those ions are produced in different ways.

1. Calculate the N-H electronegativity difference.

\[ \Delta\chi_{\ce{N-H}}=|3.04-2.20| \]

\[ \boxed{\Delta\chi_{\ce{N-H}}=0.84} \]

This indicates polar covalent bonding.

Nitrogen attracts the shared electrons more strongly than hydrogen, so the N-H bonds are polar covalent.

2. Check the valence electron count of ammonium.

Nitrogen contributes five valence electrons, and four hydrogen atoms contribute four.

The positive charge means one electron must be subtracted:

\[ 5+4-1=8 \]

These eight electrons form four N-H shared electron pairs.

The ammonium ion therefore has four single covalent bonds and no lone pairs on nitrogen.

The overall ion has a charge of \(+1\).

3. Identify the bonding in solid ammonium chloride.

Ammonium chloride contains:

\[ \ce{NH4+} \]

and

\[ \mathrm{Cl}^{-} \]

The N-H bonds inside the ammonium ion are polar covalent.

The attraction between the positively charged ammonium ions and negatively charged chloride ions is ionic.

Therefore, ammonium chloride is an ionic compound containing covalent bonds within its polyatomic ions.

Its solid does not conduct electricity because the ions are held in fixed positions within the crystal lattice.

4. Calculate the H-Cl electronegativity difference.

\[ \Delta\chi_{\ce{H-Cl}}=|3.16-2.20| \]

\[ \boxed{\Delta\chi_{\ce{H-Cl}}=0.96} \]

This indicates polar covalent bonding.

Hydrogen chloride consists of neutral molecules with shared electron pairs.

It is not an ionic lattice.

5. Compare what happens when each substance dissolves in water.

Ammonium chloride separates into ions that already existed in its ionic crystal:

\[ \ce{NH4Cl(s) -> NH4+(aq) + Cl-(aq)} \]

This is dissociation.

Hydrogen chloride reacts with water to produce ions:

\[ \ce{HCl(g) + H2O(l) -> H3O+(aq) + Cl-(aq)} \]

This is ionisation.

The two processes produce solutions containing mobile ions, allowing both solutions to conduct electricity.

However, the starting structures are different:

The crucial constraint is that conductivity identifies the presence of mobile charged particles. It does not tell you whether those particles existed before the substance dissolved.

This is why you must combine electronegativity, valence electrons, structural representations, and physical evidence rather than relying on a single observation.

Once you can identify the bonding correctly, the next step is to investigate how polar bonds combine across an entire molecule. A molecule can contain polar covalent bonds without being polar overall because of its three-dimensional shape.

For that next stage, see Bond Polarity and Electronegativity: HSC Chemistry Guide and Electronegativity and Bond Type: HSC Chemistry Guide.