Electronegativity
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You should be comfortable drawing Lewis structures and applying VSEPR theory to predict molecular geometry. Familiarity with basic vector addition is helpful but not strictly required.
Why this matters
Electronegativity explains why chemical bonds are rarely perfectly shared.
In the previous lecture, we saw that atoms generally strive to achieve a stable octet configuration. However, when two different atoms bond, they do not always share the bonding electrons equally. One atom may attract the shared pair more strongly than the other, creating a subtle imbalance in charge distribution. This imbalance is the foundation of polarity, which dictates how molecules interact with one another, how they dissolve in solvents, and how they participate in biological recognition events. Understanding electronegativity allows us to predict the direction of electron flow in reactions and the physical properties of substances, from boiling points to reactivity.
The Concept of Electronegativity
Imagine two people holding a rope in a tug-of-war. If both are equally strong, the rope stays in the middle. If one is stronger, the rope shifts toward them. In a chemical bond, the "rope" is the pair of electrons, and the "strength" is the atom's ability to attract those electrons. This ability is called electronegativity. It is an intrinsic property of an atom, determined by its nuclear charge and the distance of the bonding electrons from the nucleus. Fluorine, for instance, has the highest electronegativity because its small radius and high nuclear charge exert a very strong pull on electrons. Caesium, at the other end of the scale, has a low electronegativity because its valence electrons are far from the nucleus and shielded by inner shells, resulting in a weak attraction. When two atoms with different electronegativities bond, the electron density shifts toward the more electronegative atom, creating a partial negative charge on that atom and a partial positive charge on the other. This shift is not an all-or-nothing process; it is a continuous spectrum from perfectly covalent to fully ionic.
Quantifying the Difference
To move from qualitative descriptions to quantitative predictions, chemists use electronegativity scales, the most common being the Pauling scale. The difference in electronegativity between two bonded atoms, denoted as \(\Delta\chi\), is a direct indicator of bond polarity. A larger \(\Delta\chi\) means a greater unevenness in electron sharing. While there is no sharp boundary between covalent and ionic character, a general rule of thumb suggests that if \(\Delta\chi\) is less than 0.4, the bond is considered nonpolar covalent; between 0.4 and 1.7, it is polar covalent; and above 1.7, it is largely ionic. This scale helps us estimate the degree of charge separation without needing to measure the dipole moment directly. The relationship between the electronegativity difference and the bond energy is also significant; bonds with moderate polarity are often stronger than purely covalent or purely ionic bonds due to the electrostatic attraction between the partial charges.
From Bond Dipoles to Molecular Dipoles
A single bond dipole is a vector quantity, meaning it has both magnitude and direction. The magnitude is determined by the product of the partial charge \(\delta\) and the bond length \(d\), expressed as \(\mu = \delta d\). The direction points from the positive end to the negative end, with a cross at the positive end to indicate the sign. However, the polarity of a molecule is not simply the sum of its bond polarities; it is the vector sum of all bond dipoles. This is where molecular geometry, predicted by VSEPR theory, becomes critical. In a linear molecule like carbon dioxide, the two C=O bond dipoles are equal in magnitude but opposite in direction, so they cancel out completely, resulting in a nonpolar molecule despite having polar bonds. In contrast, in a bent molecule like water, the two O-H bond dipoles do not cancel because of the angle between them, resulting in a net molecular dipole. Thus, symmetry is the key determinant of whether a molecule is polar overall.
Periodic Trends and Predictions
Electronegativity follows predictable trends across the periodic table. It generally increases from left to right across a period because the nuclear charge increases while the atomic radius remains relatively constant, leading to a stronger pull on bonding electrons. It decreases down a group because the atomic radius increases and the shielding effect of inner electrons becomes more significant, reducing the effective nuclear charge felt by the valence electrons. These trends allow us to predict the direction of electron flow in any bond. For example, in a bond between nitrogen and hydrogen, nitrogen is more electronegative, so the electrons are pulled toward nitrogen, giving it a partial negative charge. In a bond between carbon and chlorine, chlorine is more electronegative, so the electrons are pulled toward chlorine. By combining these local predictions with the molecular geometry, we can determine the overall polarity of complex molecules, which is essential for understanding their solubility, boiling points, and chemical reactivity.
The result
Electronegativity and bond polarity
When two bonded atoms differ in electronegativity \(\chi\) (their intrinsic pull on shared bonding electrons), the bonding electron pair is shared unequally, producing partial charges \(\pm\delta\) and a bond dipole moment \(\mu=Qd\) (charge separation \(Q\) times bond length \(d\)). The molecule's overall dipole moment is the vector sum of every bond dipole, weighted by the molecular geometry (VSEPR) — so a molecule can have strongly polar bonds yet be nonpolar overall if those bond dipoles cancel by symmetry.
A common misconception is that a molecule with polar bonds must be polar. This is false. The polarity of a molecule depends on the vector sum of its bond dipoles, which is determined by its geometry. For instance, carbon tetrachloride (CCl4) has four polar C-Cl bonds, but because the molecule is tetrahedral and symmetric, the bond dipoles cancel out perfectly, making the molecule nonpolar. Always consider the 3D shape when assessing molecular polarity, not just the individual bonds. C
Worked example
Problem. Calculate the percent ionic character of a hypothetical bond between two atoms A and B, given that the observed dipole moment \(\mu_{\text{observed}}\) is 1.5 D, the bond length \(d\) is 1.2 Å, and the elementary charge \(e\) is \(1.602 \times 10^{-19}\) C. Note that 1 Debye (D) = \(3.336 \times 10^{-30}\) C·m and 1 Å = \(10^{-10}\) m.
- Convert the observed dipole moment from Debye to Coulomb-meters: \(\mu_{\text{observed}} = 1.5 \times 3.336 \times 10^{-30} = 5.004 \times 10^{-30}\) C·m.
- Convert the bond length from Ångströms to meters: \(d = 1.2 \times 10^{-10}\) m.
- Calculate the theoretical dipole moment for a fully ionic bond (\(\mu_{\text{ionic}} = ed\)): \(\mu_{\text{ionic}} = (1.602 \times 10^{-19} \text{ C}) \times (1.2 \times 10^{-10} \text{ m}) = 1.9224 \times 10^{-29}\) C·m.
- Calculate the percent ionic character using the formula \(\%\,\text{ionic} = \frac{\mu_{\text{observed}}}{\mu_{\text{ionic}}} \times 100\): \(\frac{5.004 \times 10^{-30}}{1.9224 \times 10^{-29}} \times 100 = 26.03\%.\)
Units check. The ratio of two dipole moments (C·m / C·m) is dimensionless, so the result is a percentage, which is correct.
Three quick tests
- Which atom in a C-O bond has the partial negative charge?
Answer
Oxygen, because it is more electronegative than carbon. - Is BF3 polar or nonpolar?
Answer
Nonpolar, because its trigonal planar geometry causes the three B-F bond dipoles to cancel out. - How does electronegativity change as you move down a group in the periodic table?
Answer
It decreases, due to increasing atomic radius and shielding.
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Every lecture and every result in this unit, with prerequisites and threads.
Lecture notes drafted with a local model and checked line by line by the editor; every number in the worked example was recomputed.