When two atoms form a covalent bond, they share a pair of electrons. However, the sharing is not always equal. The relative ability of an atom to attract this shared pair is described by its electronegativity.
Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond.
🔑 Key Principle
Electronegativity is determined by three main electrostatic factors: nuclear charge (number of protons), atomic radius (distance from nucleus to shared electrons), and shielding by inner shell electrons.
Electronegativity is measured using the Pauling scale, running from 0 to 4.0. Fluorine (\(\text{F}\)) is the most electronegative element, assigned a value of 4.0. It is followed closely by oxygen (3.5), nitrogen (3.0), and chlorine (3.0).
Factors Affecting Electronegativity
Three main factors determine the electronegativity of an atom: nuclear charge, atomic radius, and shielding. The electrostatic force of attraction between the positive nucleus and the shared bonding pair of electrons is dictate by these variables:
- Nuclear Charge: More protons in the nucleus exert a stronger electrostatic pull on the bonding pair of electrons.
- Atomic Radius: As atomic radius decreases, the shared pair of electrons is closer to the positive nucleus, experiencing a stronger pull.
- Shielding: Inner shells of electrons shield the outer shell from the attractive force of the nucleus. More shielding weakens the pull on the shared bonding pair.
Trends in the Periodic Table
1. Trend Across a Period (Increases)
Electronegativity increases across a period (e.g., from sodium to chlorine). This is because:
- Nuclear Charge Increases: The number of protons increases (from +11 in \(\text{Na}\) to +17 in \(\text{Cl}\)).
- Atomic Radius Decreases: Electrons are pulled in tighter.
- Shielding Remains Constant: The bonding electrons are in the same outer shell, meaning shielding by inner core electrons does not change.
Therefore, the nucleus is better able to attract the bonding pair of electrons in a covalent bond.
2. Trend Down a Group (Decreases)
Electronegativity decreases down a group (e.g., from fluorine to iodine). This is because:
- Atomic Radius Increases: There are more electron shells. The bonding pair of electrons is located further from the positive nucleus.
- Shielding Increases: There are more inner electron shells between the nucleus and the shared pair of bonding electrons.
Although the nuclear charge increases (more protons), the effects of increased distance and shielding far outweigh this increase. Consequently, the positive nucleus has a weaker electrostatic attraction for the shared bonding pair of electrons.
You must learn the definition of electronegativity exactly: "the power of an atom to attract the pair of electrons in a covalent bond". The bolded phrase is essential: do not define it simply as attracting electrons in general, which would be electron affinity. Furthermore, note that noble gases are excluded from electronegativity trends because they do not readily form covalent bonds.
Relationship to Bond Polarity
A polar covalent bond is a covalent bond in which the electron density is unequally shared between the two bonded atoms, resulting in permanent partial positive (\(\delta^+\)) and partial negative (\(\delta^-\)) charges.
Electronegativity differences between covalently bonded atoms determine the distribution of electron density in a bond:
- Non-polar covalent bond: Formed between two atoms of identical or very similar electronegativity (e.g. \(\text{H}-\text{H}\) or \(\text{C}-\text{H}\)). The electrons are shared equally.
- Polar covalent bond: Formed between two atoms of moderate electronegativity difference (e.g. \(\text{H}-\text{Cl}\) or \(\text{O}-\text{H}\)). The more electronegative atom attracts the electron density more strongly, gaining a partial negative charge (\(\delta^-\)), while the other atom becomes partially positive (\(\delta^+\)).
- Ionic bond: Formed when the difference in electronegativity is very large (typically greater than 2.0 on the Pauling scale, e.g. \(\text{Na}-\text{Cl}\)). The more electronegative atom attracts the electrons so strongly that they are completely transferred, creating positive and negative ions.
- Explain why the electronegativity of oxygen is greater than that of sulfur.
- Explain why electronegativity increases from sodium to chlorine across Period 3.
Part 1: Oxygen vs Sulfur
- Oxygen is higher up Group 16 than sulfur. Oxygen has fewer electron shells (2 shells) compared to sulfur (3 shells).
- Therefore, oxygen has a smaller atomic radius. The bonding pair of electrons in a covalent bond is closer to oxygen's nucleus.
- Oxygen also has less shielding from inner shells compared to sulfur.
- Despite sulfur having a greater nuclear charge (+16 protons vs +8 in oxygen), the smaller atomic radius and decreased shielding in oxygen outweigh this, allowing oxygen to attract the bonding pair of electrons more strongly.
Part 2: Across Period 3 (Na to Cl)
- Across Period 3, the nuclear charge increases as the number of protons increases from +11 in sodium to +17 in chlorine.
- The electrons are added to the same outermost shell, so the shielding by inner shell core electrons remains roughly constant.
- The atomic radius decreases as the stronger nuclear charge pulls the electrons in closer.
- Because of the increased nuclear charge and smaller radius, the electrostatic attraction on the shared pair of electrons in a covalent bond increases, leading to higher electronegativity.
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