Problem
State and explain the electronegativity trend in the periodic table, then apply it to rank elements and predict bond polarity.
Step 1: Meaning of electronegativity
Electronegativity (often on the Pauling scale) measures how strongly an atom attracts shared (bonding) electrons in a chemical bond. A larger electronegativity means a stronger pull on electron density in a covalent bond.
Step 2: Electronegativity trend across a period (left to right)
Across a period, the number of protons increases while added electrons enter the same principal energy level. Shielding does not increase dramatically, so the effective nuclear charge felt by valence electrons increases. The valence shell is pulled closer, atomic radius generally decreases, and the nucleus attracts bonding electrons more strongly.
Conclusion: electronegativity increases from left to right across a period because \(Z_{\text{eff}}\) increases and atomic radius decreases.
Step 3: Electronegativity trend down a group (top to bottom)
Down a group, valence electrons occupy higher principal energy levels and experience greater shielding from inner electrons. The valence shell is farther from the nucleus, so the nucleus attracts bonding electrons less strongly.
Conclusion: electronegativity decreases from top to bottom down a group because shielding increases and atomic radius increases.
Step 4: Representative Pauling electronegativities (useful anchors)
The table below lists widely used Pauling electronegativity values for periods 2 and 3 (noble gases are commonly omitted because they do not typically form standard covalent bonds on this scale).
| Period 2 element | Pauling \(\chi\) | Period 3 element | Pauling \(\chi\) |
|---|---|---|---|
| Li | 0.98 | Na | 0.93 |
| Be | 1.57 | Mg | 1.31 |
| B | 2.04 | Al | 1.61 |
| C | 2.55 | Si | 1.90 |
| N | 3.04 | P | 2.19 |
| O | 3.44 | S | 2.58 |
| F | 3.98 | Cl | 3.16 |
Step 5: Apply the electronegativity trend to ranking
Example A (same period): Rank Na, Mg, Al, Si, P, S, Cl from lowest to highest electronegativity.
Because electronegativity increases left to right across period 3:
\[ \text{Na} < \text{Mg} < \text{Al} < \text{Si} < \text{P} < \text{S} < \text{Cl} \]
Example B (same group): Rank F, Cl, Br, I from highest to lowest electronegativity.
Because electronegativity decreases down group 17:
\[ \text{F} > \text{Cl} > \text{Br} > \text{I} \]
| Halogen | Pauling \(\chi\) | Trend position |
|---|---|---|
| F | 3.98 | highest in the group |
| Cl | 3.16 | below F |
| Br | 2.96 | below Cl |
| I | 2.66 | lowest among these four |
Step 6: Use electronegativity to predict bond polarity
Bond polarity can be estimated using the electronegativity difference:
\[ \Delta \chi = \lvert \chi_A - \chi_B \rvert \]
A larger \(\Delta \chi\) generally means a more polar bond, with electron density shifted toward the more electronegative atom.
Example: Compare H–F and H–Cl (Pauling values: H \(\approx 2.20\), F \(= 3.98\), Cl \(= 3.16\)).
\[ \Delta \chi(\text{H–F}) = |3.98 - 2.20| = 1.78 \]
\[ \Delta \chi(\text{H–Cl}) = |3.16 - 2.20| = 0.96 \]
H–F is more polar than H–Cl because its electronegativity difference is larger.
Final result
The electronegativity trend increases from left to right across a period and decreases from top to bottom down a group because of changes in effective nuclear charge, shielding, and atomic radius; this trend supports reliable ranking of elements and prediction of relative bond polarity using \(\Delta \chi = |\chi_A - \chi_B|\).