Selenium dioxide, represented by the chemical formula \(\text{SeO}_2\), is a compound consisting of one selenium atom and two oxygen atoms. Molecular polarity is a fundamental property that dictates how a molecule will interact with electric fields and with other molecules, influencing everything from its melting point to its solubility.
Valence Electrons and Bonding Structure
Both selenium (\(\text{Se}\)) and oxygen (\(\text{O}\)) are elements found in Group 16 of the periodic table, meaning each atom contributes six valence electrons. With one selenium atom and two oxygen atoms, the total count for the \(\text{SeO}_2\) molecule is eighteen valence electrons.
The central atom in the structure is selenium, as it is less electronegative than oxygen, and the two oxygen atoms are positioned around it. The central selenium atom creates a double bond with each of the two surrounding oxygen atoms. This bonding arrangement utilizes eight of the eighteen valence electrons, leaving ten remaining electrons to be distributed as lone pairs.
The oxygen atoms each receive two lone pairs to complete their octets, accounting for eight electrons. The remaining two electrons must be placed on the central selenium atom, forming one non-bonding lone pair. This shows the central selenium atom is surrounded by three groups of electrons: two bonding groups (the double bonds to oxygen) and one non-bonding lone pair.
The Influence of Molecular Geometry
The three-dimensional arrangement of atoms in a molecule, known as its molecular geometry, is predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory posits that electron groups—whether they are bonding pairs or lone pairs—repel each other and will arrange themselves as far apart as possible around the central atom. The structure of \(\text{SeO}_2\) has three electron domains around the central selenium atom.
These three electron domains would ideally arrange themselves in a flat, triangular shape called a trigonal planar electron geometry. However, the molecular geometry, which describes only the position of the atoms, is different because one of the domains is a lone pair of electrons. Lone pairs exert a greater repulsive force than bonding pairs, pushing the two oxygen atoms closer together.
This increased repulsion distorts the ideal trigonal planar structure, forcing the molecule into a characteristic bent or V-shape. A fundamental rule in chemistry is that molecules with a highly symmetrical geometry, such as linear or tetrahedral shapes, are often nonpolar because their bond effects cancel out. Because the bent shape of \(\text{SeO}_2\) is inherently asymmetrical, it sets the stage for the molecule to be polar.
Electronegativity and the Overall Dipole Moment
The next step in determining polarity is to analyze the nature of the chemical bonds themselves using the concept of electronegativity. Oxygen and selenium are in the same column of the periodic table, but oxygen is higher up, making it significantly more electronegative than selenium.
On the widely used Pauling scale, the electronegativity value for selenium is approximately \(2.55\), while the value for oxygen is substantially higher. This difference in electron-pulling power means that the electrons in the \(\text{Se-O}\) bonds are not shared equally; the electrons spend more time closer to the oxygen atoms. The unequal sharing creates a partial negative charge (\(\delta^-\)) on the oxygen atoms and a partial positive charge (\(\delta^+\)) on the central selenium atom, establishing a polar bond.
Each polar bond generates a bond dipole moment, which is a vector quantity having both magnitude and direction, pointing toward the more electronegative oxygen atom. For a molecule to be nonpolar, these individual bond dipoles must perfectly cancel each other out due to symmetry. In a symmetrical linear molecule, two equal and opposite bond dipoles would result in a net cancellation.
However, since \(\text{SeO}_2\) has a bent molecular geometry, the two \(\text{Se-O}\) bond dipoles are not pulling in opposite directions. Instead, they are pulling at an angle, and their effects add together to create a non-zero, overall net molecular dipole moment. Experimental evidence confirms this, showing that a single \(\text{SeO}_2\) molecule in the gas phase has a measured dipole moment of \(2.62\) Debye (D). The existence of a net dipole moment confirms the molecule has a separation of charge, leading to the definitive conclusion that \(\text{SeO}_2\) is a polar molecule.
Solubility and Other Properties
The established polarity of selenium dioxide directly influences its physical and chemical behavior, particularly its solubility. A simple rule of thumb in chemistry, often stated as “like dissolves like,” governs solubility, meaning polar substances tend to dissolve polar solvents, and nonpolar substances dissolve nonpolar solvents. Since \(\text{SeO}_2\) is a polar molecule, it exhibits high solubility in polar solvents.
Selenium dioxide is highly soluble in water, a classic polar solvent, where it dissolves to form selenous acid. At \(25^\circ\text{C}\), the solubility of \(\text{SeO}_2\) in water is approximately \(39.5\) grams per \(100\) milliliters. This high solubility is a direct consequence of the strong attractive forces, specifically dipole-dipole interactions, that form between the polar \(\text{SeO}_2\) molecules and the polar water molecules.
The compound is also soluble in other polar organic solvents such as ethanol and acetone. Conversely, \(\text{SeO}_2\) is notably insoluble in nonpolar solvents like hydrocarbons. This difference in solubility provides a macroscopic, observable validation of the theoretical prediction that the selenium dioxide molecule possesses a permanent net dipole moment.