Electronegativity and Bonding

2026 Syllabus Objectives

By the end of this topic, you should be able to:

  1. Define electronegativity as the power of an atom to attract electrons to itself
  2. Explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells
  3. State and explain the trends in electronegativity across a period and down a group of the Periodic Table
  4. Use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds

1. What is Electronegativity?

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Electronegativity is the power (or ability) of an atom to attract electrons to itself when it is bonded to another atom in a covalent bond.

Think of it like a "tug of war" between two atoms sharing electrons. The atom with higher electronegativity pulls the shared electrons closer to itself.

The Pauling Scale

Scientists measure electronegativity using the Pauling scale, named after chemist Linus Pauling. This scale gives each element a number that shows how strongly it attracts electrons.

Key points about the Pauling scale:

  • Fluorine (F) is the most electronegative element with a value of 4.0
  • Values range from about 0.7 (for francium) to 4.0 (for fluorine)
  • Noble gases (like helium and neon) are not usually given electronegativity values because they rarely form bonds

Example values:

  • Hydrogen (H): 2.1
  • Carbon (C): 2.5
  • Nitrogen (N): 3.0
  • Oxygen (O): 3.5
  • Fluorine (F): 4.0
  • Chlorine (Cl): 3.2

Why Electronegativity Matters

When two atoms with different electronegativities form a bond, the shared electrons are not equally distributed. The more electronegative atom pulls the electrons closer, creating partial charges in the molecule. This affects the properties of the substance, such as whether it dissolves in water or how it reacts with other chemicals.

2. Factors That Affect Electronegativity

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Three main factors determine how electronegative an atom is:

Factor 1: Nuclear Charge

Nuclear charge refers to the number of protons in the nucleus of an atom.

  • More protons = stronger positive charge in the nucleus
  • A stronger positive charge attracts the negative electrons more strongly
  • Therefore: Higher nuclear charge = Higher electronegativity

Example:

  • Sodium (Na) has 11 protons → electronegativity = 0.9
  • Magnesium (Mg) has 12 protons → electronegativity = 1.2
  • Aluminium (Al) has 13 protons → electronegativity = 1.5

As you move across the period from Na to Al, the number of protons increases, so the nucleus pulls on electrons more strongly.

Factor 2: Atomic Radius

Atomic radius is the distance from the nucleus to the outermost electrons.

  • Smaller atomic radius = electrons are closer to the nucleus
  • Electrons that are closer to the nucleus feel a stronger pull
  • Therefore: Smaller atomic radius = Higher electronegativity

Why this happens: The force of attraction between two charged particles gets weaker as they get further apart. If the outer electrons are far from the nucleus, the nucleus cannot pull on them as strongly.

Example: In a small atom like fluorine, the bonding electrons are very close to the nucleus, so they are strongly attracted. In a large atom like iodine, the bonding electrons are much further from the nucleus, so they are attracted less strongly.

Factor 3: Shielding (or Screening)

Shielding happens when inner electron shells block (or "shield") the outer electrons from feeling the full attractive force of the nucleus.

  • More inner electron shells = more shielding
  • More shielding means outer electrons feel less attraction from the nucleus
  • Therefore: More shielding = Lower electronegativity

How shielding works: Imagine the nucleus as a magnet and the outer electrons as pieces of metal. If you place layers of cardboard (inner electron shells) between the magnet and the metal, the magnetic pull gets weaker. Each electron shell acts like a layer of cardboard.

Example:

  • Sodium (Period 3, Group 1) has fewer electron shells than caesium (Period 6, Group 1)
  • Caesium has more shielding, so its outer electrons feel less attraction from the nucleus
  • Result: Sodium has higher electronegativity than caesium

How These Factors Work Together

All three factors interact:

  • Across a period: Nuclear charge increases, but atomic radius decreases and shielding stays roughly the same → electronegativity increases
  • Down a group: Nuclear charge increases, but atomic radius increases a lot and shielding increases a lot → the effects of radius and shielding outweigh the nuclear charge increase → electronegativity decreases

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