Molecular polarity describes the uneven distribution of electrical charge across a compound, resulting in one end of the molecule having a slight negative charge and the opposite end having a slight positive charge. Selenium Hydride (\(\text{SeH}_2\)) is definitively a polar molecule. This polarity requires two conditions: the presence of polar chemical bonds and an asymmetrical three-dimensional molecular shape.
The Building Blocks of Polarity
For a molecule to be polar, it must first contain polar bonds. Bond polarity is determined by electronegativity, which measures an atom’s ability to attract shared electrons within a chemical bond. When atoms share electrons unequally, the stronger atom pulls the electrons closer, creating a polar bond and a bond dipole moment. The atom with higher electronegativity gains a partial negative charge (\(\delta^-\)), and the other gains a partial positive charge (\(\delta^+\)).
The electronegativity of Hydrogen (H) is \(2.2\), and Selenium (Se) is \(2.55\). This difference of \(0.35\) ensures that the Selenium-Hydrogen (\(\text{Se-H}\)) bonds are polar covalent, with electron density shifted toward the Selenium atom.
Molecular Shape and Symmetry
The presence of polar bonds alone does not guarantee molecular polarity. The second determining factor is the molecular geometry, or the three-dimensional arrangement of the atoms. If the molecule is highly symmetrical, individual bond dipoles may perfectly cancel each other out, resulting in a net dipole moment of zero.
For example, in linear, tetrahedral, or trigonal planar molecules, the vector sum of the dipoles cancels, making the molecule nonpolar despite having polar bonds. Polarity only occurs when the spatial arrangement is asymmetrical, preventing cancellation. Asymmetrical shapes, such as bent or pyramidal structures, ensure the electrical charges remain unevenly distributed, creating a net molecular dipole moment. This geometry is predicted by the Valence Shell Electron Pair Repulsion (\(\text{VSEPR}\)) theory, which minimizes electron repulsion.
Analyzing Selenium Hydride (\(\text{SeH}_2\))
Applying these principles confirms \(\text{SeH}_2\)‘s polar nature through both its bonding and geometry. The electronegativity difference establishes that each \(\text{Se-H}\) bond carries a dipole moment, with Selenium pulling the shared electrons more strongly. This results in partial negative charges on the central Selenium atom and partial positive charges on the two Hydrogen atoms.
Selenium has six valence electrons, using two for bonding and leaving two non-bonding lone pairs. These four electron domains (two bonding pairs and two lone pairs) repel each other, attempting a tetrahedral arrangement. However, the lone pairs exert a stronger repulsive force, distorting the shape into a \(\text{bent}\) or \(\text{V-shaped}\) molecular geometry.
Because the molecule is bent, the two \(\text{Se-H}\) bond dipoles are directed at an angle and do not perfectly oppose each other. This imperfect cancellation creates a permanent, measurable net dipole moment (approximately \(0.81 \text{ D}\)), classifying \(\text{SeH}_2\) as a polar molecule.
How Polarity Affects Molecular Behavior
The polarity of \(\text{SeH}_2\) dictates how the molecule interacts with its environment. Polar molecules exhibit dipole-dipole forces, which are attractive forces between the positive end of one molecule and the negative end of a neighbor. These forces are stronger than the London dispersion forces found in nonpolar compounds of similar size.
This enhanced intermolecular attraction affects physical properties, primarily the boiling point. \(\text{SeH}_2\) has a higher boiling point than comparable nonpolar molecules because extra energy is required to overcome these attractions. Polarity also governs solubility, following the rule that “like dissolves like.” Since \(\text{SeH}_2\) is polar, it mixes readily with other polar solvents, such as water.
\(\text{SeH}_2\) is structurally analogous to water (\(\text{H}_2\text{O}\)) and hydrogen sulfide (\(\text{H}_2\text{S}\)), all of which are bent and polar. The boiling point of \(\text{SeH}_2\) (around \(-41.25^\circ \text{C}\)) is higher than \(\text{H}_2\text{S}\) due to the greater mass and polarizability of the Selenium atom, which increases London forces. However, its boiling point is significantly lower than water’s because Selenium is not electronegative enough to participate in hydrogen bonding.