Question

For the indicated atom in each of the following molecules or ions, give the number and arrangement of the electron pairs around that atom. a. \(\mathrm{As}\) in \(\mathrm{AsO}_{4}^{3-}\) b. Se in \(\mathrm{SeO}_{4}^{2-}\) c. \(\mathrm{S}\) in \(\mathrm{H}_{2} \mathrm{~S}\)

Short answer

a. In AsO₄³⁻, As is surrounded by 4 bonding pairs and 12 non-bonding pairs, making it sp³ hybridized with a tetrahedral electron pair arrangement. b. In SeO₄²⁻, Se is surrounded by 4 bonding pairs and 12 non-bonding pairs, making it sp³ hybridized with a tetrahedral electron pair arrangement. c. In H₂S, S is surrounded by 2 bonding pairs and 2 non-bonding pairs, making it sp³ hybridized with a bent or V-shaped electron pair arrangement.

Step-by-step solution

Determine the Valence Electrons

For each of the atoms, we will first find the number of valence electrons they possess. This can be found by looking at their group number in the periodic table. a. As in AsO₄³⁻: As is in Group 15, so it has 5 valence electrons b. Se in SeO₄²⁻: Se is in Group 16, so it has 6 valence electrons c. S in H₂S: S is in Group 16, so it has 6 valence electrons

Calculate TotalValence Electrons

We'll now find the total number of valence electrons in each molecue/ion by taking into account the atoms it is bonded to and the charge. a. AsO₄³⁻: Total valence electrons = (1 x As (5)) + (4 x O (6)) + 3 (charge) = 5 + 24 + 3 = 32 b. SeO₄²⁻: Total valence electrons = (1 x Se (6)) + (4 x O (6)) +2 (charge)= 6 + 24 + 2 = 32 c. H₂S: Total valence electrons = (2 x H (1)) + (1 x S (6) = 2 + 6 = 8

Calculate Number of Bonding and Non-bonding Electron Pairs

We will now find the number of bonding and non-bonding electron pairs around the central atom. a. AsO₄³⁻: Bonding electrons = 4 (As-O bonds) x 2 = 8 Non-bonding electrons = Total valence electrons (32) - Bonding electrons (8) = 24 Non-bonding electron pairs = 24/2 = 12 b. SeO₄²⁻: Bonding electrons = 4 (Se-O bonds) x 2 = 8 Non-bonding electrons = Total valence electrons (32) - Bonding electrons (8) = 24 Non-bonding electron pairs = 24/2 = 12 c. H₂S: Bonding electrons = 2 (S-H bonds) x 2 = 4 Non-bonding electrons = Total valence electrons (8) - bonding electrons (4) = 4 Non-bonding electron pairs = 4/2 = 2

Describe Arrangement of Electron Pairs

With the bonding and non-bonding electron pair counts, we can now provide the arrangement of electron pairs around the central atom. a. AsO₄³⁻: As is surrounded by 4 bonding pairs (As-O bonds) and 12 non-bonding pairs, making it sp³ hybridized. The electron pair arrangement is tetrahedral. b. SeO₄²⁻: Se is surrounded by 4 bonding pairs (Se-O bonds) and 12 non-bonding pairs, making it sp³ hybridized. The electron pair arrangement is tetrahedral. c. H₂S: S is surrounded by 2 bonding pairs (S-H bonds) and 2 non-bonding pairs, making it sp³ hybridized. The electron pair arrangement is bent or V-shaped (similar to the water molecule).

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial to understanding chemical bonding and molecular structure. They are responsible for the chemical properties of the element since they are the ones that participate in forming chemical bonds.

For example, arsenic (As) in \texttt{AsO\(_4^{3-}\)} has 5 valence electrons because it is in Group 15 of the periodic table. Similarly, selenium (Se) has 6 valence electrons being in Group 16, as does sulfur (S). These valence electrons can be involved in forming covalent bonds with other atoms or can exist as lone pairs, not bonded to any other atom.

When analyzing a molecule or ion, it is critical to calculate the total valence electrons by considering those contributed by each atom and the overall charge of the ion. For instance, \texttt{AsO\(_4^{3-}\)} has a total of 32 valence electrons, taking into account the contribution from the oxygen atoms and the negative charge.

By understanding valence electrons, we lay the foundation for further exploration into the world of molecular geometry and chemical reactivity.

Hybridization

Hybridization is a concept in chemistry that describes the mixing of atomic orbitals to form new hybrid orbitals, which then arrange themselves to optimize the spatial distribution around a central atom. It's like mixing different colors of paints to create new shades that are best suited for a particular image.

The number of hybrid orbitals formed is equal to the number of atomic orbitals mixed, and this mixing is determined by the types of bonds and the number of electron pairs around the central atom.

For instance, in the \texttt{AsO\(_4^{3-}\)} or \texttt{SeO\(_4^{2-}\)} ions, the central atoms (As and Se) undergo sp³ hybridization. This indicates that one s orbital mixes with three p orbitals to form four equivalent sp³ hybrid orbitals, which are then arranged in a tetrahedral shape to minimize electron repulsion.

In the case of \texttt{H\(_2\)S}, sulfur is also sp³ hybridized; despite only two bonds, the presence of lone pairs on the sulfur indicates such hybridization, resulting in a bent molecular shape. This concept allows chemists to predict and explain the shapes of molecules and their resulting properties.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. Determining molecular geometry is crucial as it influences a molecule's reactivity, color, magnetism, biological activity, and even its physical state at a given temperature and pressure.

The geometry is determined by the electron pair arrangement, which seeks to minimize the repulsion between electron pairs according to the Valence Shell Electron Pair Repulsion (VSEPR) theory.

For example, in a molecule like \texttt{AsO\(_4^{3-}\)}, the tetrahedral electron pair arrangement around arsenic is due to the four bonding pairs of electrons. The same tetrahedral arrangement is observed in \texttt{SeO\(_4^{2-}\)}.

In contrast, \texttt{H\(_2\)S} has a 'bent' or 'V-shaped' geometry because the sulfur atom has two bonding pairs and two lone pairs. The presence of lone pairs affects the molecular geometry since lone pairs occupy more space around the nucleus than bonding pairs, leading to deviations from the ideal geometry.

Understanding molecular geometry is not just a theoretical exercise; it is essential for predicting the behavior of molecules in a real-world context, guiding scientists and students alike in their exploration of chemical interactions.