Question

Account for each of the following observations: (a) Oxygen is more electronegative than sulfur. (b) Sulfur forms long \(\mathrm{S}_{n}\) chains, but oxygen does not. (c) The \(\mathrm{SO}_{3}\) molecule is trigonal planar, but the \(\mathrm{SO}_{3}{ }^{2-}\) ion is trigonal pyramidal.

Short answer

(a) Oxygen is more electronegative due to higher effective nuclear charge. (b) Sulfur forms chains via atomic size and available d-orbitals. (c) Presence of lone pairs in \\( \\mathrm{SO}_{3}^{2-}\\) shapes it trigonal pyramidal.

Step-by-step solution

Understanding Electronegativity

Electronegativity refers to the ability of an atom to attract shared electrons. Oxygen is more electronegative than sulfur because it is positioned higher on the periodic table, which means it has a greater effective nuclear charge due to fewer electron shells. These characteristics lead to oxygen having a stronger pull on electrons compared to sulfur.

Explaining Chain Formation

Sulfur can form long chains because sulfur atoms readily bond with each other. This is facilitated by its larger atomic size and availability of d-orbitals for bonding, leading to stronger interactions between sulfur atoms. Oxygen atoms are smaller and lack d-orbitals, thus limiting their ability to form extensive chains and instead preferring to form double bonds or small cyclic structures.

Comparing Molecular Geometries

The \(\mathrm{SO}_{3}\) molecule is trigonal planar because it adopts a geometry where sulfur is double-bonded to three oxygen atoms, minimizing repulsion between electron pairs at 120-degree angles. In contrast, the \(\mathrm{SO}_{3}^{2-}\) ion is trigonal pyramidal due to the addition of two extra electrons. This introduces a lone pair on the sulfur atom, resulting in a tetrahedral arrangement of electronic pairs which causes the shape to become pyramidal as opposed to flat.

Key concepts

Electronegativity

Electronegativity is the tendency of an atom to attract shared electrons in a chemical bond. This concept plays a crucial role in understanding the differences between oxygen and sulfur, especially in terms of their chemical behavior.
Oxygen is more electronegative than sulfur. This is primarily because oxygen is located higher up in Group 16 of the periodic table. It means that oxygen has a stronger effective nuclear charge due to having fewer electron shells that shield the pull of the nucleus on valence electrons. As a result, oxygen exerts a stronger attraction on shared electrons when forming bonds.
Why does this matter? A higher electronegativity influences how elements participate in chemical reactions and bond formations. This property allows oxygen to form strong polar bonds by drawing the shared electrons closer to itself. In comparison, sulfur, with a slightly lower electronegativity, forms less polar bonds, providing it a different set of chemical characteristics.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. It helps predict the shape and the type of bonds formed in a molecule.
Let's consider \( \mathrm{SO}_3 \), which manifests as a trigonal planar molecule. Here, sulfur forms double bonds with three oxygen atoms, arranged evenly around it. Each double bond counts as a single electron domain resulting in three domains around sulfur. These domains orient themselves to be as far apart as possible, minimizing electron pair repulsion according to VSEPR theory, resulting in the bond angles of 120 degrees.
Conversely, when we talk about the \( \mathrm{SO}_3^{2-} \) ion, there is an addition of two extra electrons. These electrons become a lone pair on the sulfur atom. Imagine placing four electron domains around sulfur now—three bonding pairs and one lone pair. This lone pair influences the spatial arrangement by pushing the bonded pairs closer together, hence transitioning the shape from trigonal planar to trigonal pyramidal.
This adjustment showcases the impact electron pairs, particularly lone pairs, have on molecular geometry.

Bond Formation

Bond formation is the process by which atoms join together to form molecules or compounds. This involves interactions between the outermost electrons of atoms, balancing attractive and repulsive forces.
Sulfur's ability to form long chains, like \( \mathrm{S}_n \) chains, is notable in its chemistry. Its relatively larger atomic size compared to oxygen enables sulfur to form strong covalent bonds with other sulfur atoms. Moreover, sulfur possesses available d-orbitals which facilitate additional bonding and enable stronger, more extensive interactions between sulfur atoms. This allows for an array of complex structures through single, double, or even weaker multicoordinate bonds.
Oxygen, on the other hand, prefers forming double bonds, often resulting in smaller, more stable rings or linear forms. Its smaller atomic size and lack of available d-orbitals constrain oxygen to shorter and simpler bonding arrangements.
Understanding these differences in bond formation helps explain why elements have diverse structural forms and functionalities, emphasizing the importance of atomic size and orbital availability in determining chemical behavior.