Hydrogen sulfide (\(\text{H}_2\text{S}\)) is a molecular compound consisting of one sulfur atom bonded to two hydrogen atoms. Analyzing the physical properties of this molecule requires a detailed examination of the forces that act between individual \(\text{H}_2\text{S}\) units. These forces arise from an uneven distribution of electrical charge within the molecule, which is a consequence of its atomic composition and three-dimensional shape.
Understanding Intermolecular Forces and Polarity
Molecules attract one another through weak electrical attractions known as intermolecular forces (IMFs). These attractions exist between separate molecules, distinguishing them from the stronger covalent or ionic bonds that hold atoms within a single molecule. Dipole-dipole forces represent one specific type of IMF, which acts only between molecules that are electrically polar. A polar molecule possesses a permanent, asymmetrical distribution of electrical charge, creating a positive end and a negative end.
This charge separation is described by a net dipole moment. The positive end of one polar molecule is consistently attracted to the negative end of an adjacent polar molecule, resulting in the dipole-dipole force.
Molecular polarity depends on two factors: the polarity of the individual bonds within the molecule and the overall geometry of the molecule. Even if the individual bonds are polar, the molecule may still be non-polar if its structure is perfectly symmetrical, causing the bond polarities to cancel each other out.
How Polar Bonds Form in \(\text{H}_2\text{S}\)
The polarity of a chemical bond is determined by the difference in electronegativity between the two atoms involved. Electronegativity is an atom’s inherent power to attract the shared pair of electrons toward itself. The greater the difference in this attracting power, the more unequally the electrons are shared, leading to a polar covalent bond.
Sulfur (S) and Hydrogen (H) possess measurably different electronegativity values. Sulfur has a value of approximately \(2.58\), while hydrogen is slightly less electronegative at about \(2.20\). Since sulfur has the higher value, it exerts a stronger pull on the shared electrons in the \(\text{H-S}\) bonds.
This unequal sharing creates a separation of charge within each individual bond, resulting in a bond dipole moment. The sulfur atom acquires a partial negative charge (\(\delta-\)), and the hydrogen atoms acquire partial positive charges (\(\delta+\)). Each of the two \(\text{H-S}\) bonds is considered a polar covalent bond.
The Critical Influence of Molecular Geometry
The existence of polar bonds does not guarantee that the entire molecule will be polar because the effect of the individual bond dipoles must be considered vectorially. If the molecule’s shape is highly symmetrical, the opposing bond dipoles can effectively cancel each other out, resulting in a net dipole moment of zero. This is a common feature in molecules like carbon dioxide (\(\text{CO}_2\)), which has polar bonds but is non-polar overall due to its linear geometry.
The geometry of the \(\text{H}_2\text{S}\) molecule is predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory. The central sulfur atom in \(\text{H}_2\text{S}\) is surrounded by four electron pairs: two are bonding pairs with the hydrogen atoms, and two are non-bonding lone pairs.
These four electron pairs arrange themselves in a three-dimensional tetrahedral electron-pair geometry. The molecular geometry, which only considers the positions of the atoms, is distorted from this ideal shape. The two lone pairs exert a greater repulsive force on the bonding pairs, pushing the two hydrogen atoms closer together.
This repulsion results in a “bent” or “V-shaped” molecular geometry for \(\text{H}_2\text{S}\), with an \(\text{H-S-H}\) bond angle of approximately \(92^\circ\). This bent shape is inherently asymmetrical, meaning the two individual bond dipole moments are not directed in opposite, canceling directions. The \(\text{H}_2\text{S}\) molecule possesses a permanent, non-zero net dipole moment.
Determining the Net Dipole and Existing Forces
The combination of polar \(\text{H-S}\) bonds and the molecule’s asymmetrical bent geometry confirms that \(\text{H}_2\text{S}\) is a polar molecule. The experimentally determined dipole moment of hydrogen sulfide is approximately \(0.97\) Debye, which confirms the molecule’s polarity. This permanent net dipole moment means that hydrogen sulfide molecules are attracted to one another through dipole-dipole interactions.
The positive region of one \(\text{H}_2\text{S}\) molecule will align and attract the negative region of a neighboring molecule. This specific attraction is the dipole-dipole force at work, making it the strongest intermolecular force present in hydrogen sulfide.
The molecule also exhibits London Dispersion Forces (LDF), which are temporary attractions caused by random electron movement. LDF is present in all molecules, but for \(\text{H}_2\text{S}\), the permanent dipole-dipole force is significantly stronger than the transient LDF. \(\text{H}_2\text{S}\) does not exhibit hydrogen bonding because sulfur is not one of the highly electronegative atoms (oxygen, nitrogen, or fluorine) required to form this strong type of dipole-dipole interaction.