Explain electronegativity and predict bonding
Connect electronegativity to atomic attraction, explain periodic trends, and use supplied Pauling differences to predict ionic or covalent bonding.
AS Level; Cambridge International Chemistry 9701, 2025–2027 syllabus. Use the supplied Pauling values when given, show particle and bond reasoning, and distinguish intramolecular bonds from intermolecular forces. Written practice is self-checked
Before you begin
- Read element symbols and simple molecular formulae; use the Periodic Table.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
What does electronegativity describe?
What you will learn
- Define electronegativity
- Explain its trends using nuclear charge, radius and shielding
- Use Pauling differences to predict ionic or covalent bonding
Electronegativity describes attraction within a bond
Electronegativity is the power of an atom to attract electrons to itself. It is a relative scale; use the Pauling values supplied for the atoms in the question rather than treating electronegativity as a directly measured atomic charge.
Across a period, nuclear charge generally increases while added electrons enter the same principal shell. Greater attraction pulls bonding electrons more strongly. Down a group, the bonding electrons are farther from the nucleus and more shielded by inner shells and sub-shells, so the attraction generally decreases.
A difference helps predict the bond type
Compare the electronegativities of the bonded atoms. A larger difference supports ionic bonding; a smaller difference supports covalent bonding. Use the values and context in the question, and do not invent a universal numerical boundary unless the question supplies one.
This syllabus outcome asks candidates to use Pauling differences to predict ionic or covalent bonding. It does not require predicting covalent character within ionic compounds.
Pause and explain
Why does electronegativity generally decrease down a group?
Worked example
Pauling values are Na 0.9, Cl 3.0, and H 2.1. Compare the likely bond types in NaCl and HCl.
Show the worked solution
- Na–Cl difference = 3.0 − 0.9 = 2.1.
- H–Cl difference = 3.0 − 2.1 = 0.9.
- The much larger first difference supports ionic bonding; the smaller second difference supports covalent bonding.
Answer NaCl is predicted to be ionic; HCl is predicted to be covalent from the supplied values.
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Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
Explain the general electronegativity trend across Period 2 using nuclear charge, radius and shielding.
Give me a hint
Follow an outer bonding electron across the period.
Compare my reasoning
- Nuclear charge increases.
- Added electrons enter the same principal shell, so shielding does not rise by the same amount.
- The stronger attraction and smaller radius raise electronegativity overall.
Electronegativity generally increases across Period 2 because nuclear charge rises, shielding changes less, and atomic radius decreases, strengthening attraction for bonding electrons.
Look for these in your work
- Trend stated
- Nuclear charge included
- Shielding and radius linked to attraction
Use supplied Pauling values Li 1.0, F 4.0, C 2.5 and H 2.1 to predict whether LiF and CH₄ are mainly ionic or covalent.
Give me a hint
Find each bonded pair's electronegativity difference and compare the relative sizes.
Compare my reasoning
- Li–F difference is 3.0.
- C–H difference is 0.4.
- A much larger difference supports ionic bonding; the smaller difference supports covalent bonding.
LiF is predicted to be ionic; the C–H bonds in CH₄ are predicted to be covalent.
Look for these in your work
- Both differences calculated
- Ionic prediction linked to larger difference
- Covalent prediction linked to smaller difference
A student says a larger electronegativity difference proves a compound is always ionic and gives each atom a full integer charge. Correct the claim.
Give me a hint
Separate a bonding prediction from an exact charge assignment.
Compare my reasoning
- The Pauling difference helps predict ionic versus covalent bonding.
- It is not itself an atomic charge.
- Use the supplied values and evidence; do not invent a universal cutoff.
The difference is evidence for predicting bond type, not a direct charge measurement. A larger difference supports an ionic prediction, but the value alone does not assign exact integer charges or establish a universal threshold.
Look for these in your work
- Scale purpose explained
- Difference distinguished from charge
- Limit of the inference stated
Common mistakes
- Defining electronegativity as an atom's size
- Explaining a trend by nuclear charge alone
- Treating a Pauling difference as a guaranteed charge
- Applying an unstated universal ionic threshold
What can you explain now?
What does electronegativity describe?
Compare with the explanation
An atom's power to attract electrons to itself
It describes attraction for bonding electrons.
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Make it stick
How do nuclear charge, radius and shielding change the attraction for bonding electrons?
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.