Polarity of molecules and chemical bonds

Nature of the chemical bond and electron distribution

Chemical interaction is based on the ability of atoms to share electrons. If identical atoms combine (as in simple substances), electron pairs are distributed between them strictly equally. Such a bond is called nonpolar.

However, in complex substances containing different elements, the situation changes. Atoms with higher values of ionization energy ($E_i$) and electron affinity ($E_{ea}$) attract shared electron clouds more strongly. As a result, an excess of negative charge appears at one end of the bond, while a deficit appears at the other (forming a positive pole). This forms a polar covalent bond, and the particle turns into a dipole.

A striking example is hydrogen fluoride ($HF$). Fluorine, being an extremely electronegative element, strongly pulls hydrogen's electrons toward itself. Despite the high degree of ionicity (about 40–45%), complete separation of the electron does not occur. Effective partial charges ($\delta$) of approximately 0.4–0.5 of an electron charge are formed on the atoms.

Influence of geometry on molecular polarity

The presence of polar bonds within a molecule does not guarantee that the molecule itself will be polar. For polyatomic systems (consisting of three or more atoms), their spatial structure plays a critically important role. The dipole moments of each individual bond are vector quantities, and the overall polarity depends on their addition.

• Symmetrical molecules: If a molecule has a high degree of symmetry, the vectors can cancel each other out, and the net moment will be zero. For example, in carbon dioxide ($CO_2$), the carbon-oxygen bonds are polar, but their linear structure leads to mutual cancellation. Methane ($CH_4$) behaves similarly. Due to its regular tetrahedral shape (bond angle $109.5^\circ$), the vector sum of the four dipoles is zero, so methane is nonpolar.
• Asymmetrical molecules: A violation of symmetry makes a molecule polar. A classic example is water ($H_2O$), which has an angular structure, preventing the dipole moments from balancing out.

Features of the ionic bond

When the electronegativity difference between interacting atoms becomes significant (usually $\Delta\chi_r \ge 2$), the bond develops into an ionic bond. It represents electrostatic attraction between positive and negative ions. Valency in such compounds equals the number of actually donated or accepted electrons.

In nature, there are no bonds with one hundred percent electron transfer. Even with an electronegativity difference $\ge 2$, the actual ionicity does not reach 70%, and the charge on the atom is about 0.63. Such bonds are characteristic of the interaction between metals and typical nonmetals—oxygen or halogens. In this case, large-scale ionic crystal lattices (salts, oxides, hydroxides) are formed rather than discrete molecules.

Donor-acceptor mechanism and dipole moments

A covalent bond can also form via a donor-acceptor mechanism. The formation of the complex ion $[Zn(NH_3)_4]^{2+}$ is a clear example. Nitrogen atoms from ammonia molecules act as donors, providing lone electron pairs. The zinc ion ($Zn^{2+}$, configuration $...3d^{10}4s^04p^0$) acts as the acceptor. It provides vacant orbitals (one $4s$ and three $4p$) that undergo $sp^3$ hybridization, creating the tetrahedral shape of the complex.

Polarity varies greatly even among substances with similar structures. When comparing ammonia ($NH_3$) and phosphine ($PH_3$), a vast difference is noticeable. Nitrogen has a much higher electronegativity than phosphorus. The $N-H$ bonds are polar, and together with the lone electron pair, they create a high overall dipole moment. In phosphine, the electronegativities of the elements are close, so the entire molecule has an extremely low dipole moment.

Class and exam questions

What else to study on this topic

• Ionization energy and electron affinity: physical meaning of the parameters
• Calculation of effective atomic charges in polar compounds
• Orbital hybridization theory using zinc complex compounds as an example
• Influence of lone electron pairs on geometry and polarity
• Crystal lattices: properties of substances with an ionic bond type

Other topics in "Chapter 5. Structure of Matter"

Prepared by WhiteCoat Pro faculty in accordance with the medical school curriculum. For educational purposes only.