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. \(P\) in \(\mathrm{PO}_{4}^{3-}\) b. \(\mathrm{S}\) in \(\mathrm{SO}_{4}^{2-}\) c. \(\operatorname{Sin} \mathrm{H}_{2} \mathrm{S}\)

Short answer

a. In \(PO_4^{3-}\), there are 4 electron pairs around \(P\), arranged in a tetrahedral geometry. b. In \(SO_4^{2-}\), there are 4 electron pairs around \(S\), arranged in a tetrahedral geometry. c. In \(H_2S\), there are 4 electron pairs around \(S\), arranged in a tetrahedral geometry, but due to the presence of lone pairs, the molecular geometry is bent or V-shaped.

Step-by-step solution

Draw Lewis Structures

For each given molecule/ion, determine the total number of valence electrons and draw the corresponding Lewis structure. a. \(P\) in \(PO_4^{3-}\) Total valence electrons = 5 (from \(P\)) + 6*4 (from 4 \(O\)) + 3 (extra electrons due to negative charge) = 32. b. \(S\) in \(SO_4^{2-}\) Total valence electrons = 6 (from \(S\)) + 6*4 (from 4 \(O\)) + 2 (extra electrons due to negative charge) = 32. c. \(S\) in \(H_2S\) Total valence electrons = 6 (from \(S\)) + 1*2 (from 2 \(H\)) = 8.

Determine Number and Arrangement of Electron Pairs

Now that we have the Lewis structures, let's determine the number and arrangement of electron pairs around the central atom. a. \(P\) in \(PO_4^{3-}\) In the Lewis structure of \(PO_4^{3-}\), \(P\) is the central atom bonded to four \(O\) atoms. It has no lone pairs just bonded electron pairs, so there are 4 electron pairs around \(P\). These electron pairs are arranged in a tetrahedral geometry. b. \(S\) in \(SO_4^{2-}\) In the Lewis structure of \(SO_4^{2-}\), \(S\) is the central atom bonded to four \(O\) atoms with 2 double bonds and 2 single bonds. It has no lone pairs just bonded electron pairs, so there are 4 electron pairs around \(S\). These electron pairs are also arranged in a tetrahedral geometry. c. \(S\) in \(H_2S\) In the Lewis structure of \(H_2S\), \(S\) is the central atom bonded to two \(H\) atoms and has 2 lone pairs of electrons. In this case, there are 4 electron pairs around \(S\) (including the lone pairs), and the electron pairs are arranged in a tetrahedral geometry, but due to the presence of lone pairs, the molecular geometry will be bent or V-shaped.

Key concepts

Valence Electrons

Valence electrons play a crucial role in chemistry. These are the outermost electrons in an atom and are involved in chemical bonding. The number of valence electrons determines how an atom will interact with others. To find the total number of valence electrons in a molecule:

• Add the valence electrons from each atom. For instance, phosphorus (P) in \(PO_4^{3-}\) has 5 valence electrons.
• Account for electrons added or removed due to the molecule's overall charge. A negative charge means extra electrons. In \(PO_4^{3-}\), three extra electrons are added.

Calculating valence electrons accurately is essential. It helps predict how atoms will bond. Getting this wrong can lead to incorrect molecular structures.

Tetrahedral Geometry

Tetrahedral geometry is a common structure in molecular bonding. It occurs when four pairs of electrons surround a central atom. A perfect example is the arrangement in \(PO_4^{3-}\) and \(SO_4^{2-}\).In tetrahedral geometry, the four electron pairs are spread out evenly around the central atom, creating a shape with four corners. Think of this as a three-dimensional shape, almost like a pyramid with a triangular base:

• All the bond angles are approximately \(109.5^\circ\). This provides a stable configuration.
• The symmetrical shape allows for the minimization of electron pair repulsions according to VSEPR theory.

It's essential to understand this aspect, especially when determining how a molecule will behave or react with others.

Lone Pairs

Lone pairs are pairs of valence electrons not shared between atoms in a molecule. These electrons reside on the central atom but do not participate directly in bonding. Let's take \(H_2S\) as an example:Sulfur, the central atom, has two lone pairs. While the electron pairs around sulfur are in a tetrahedral arrangement, the presence of lone pairs affects the molecule's overall shape.

• Lone pairs occupy more space than bonding pairs. This results in a bent or V-shaped molecular geometry for \(H_2S\).
• The presence of lone pairs can influence polarities and reactivity. They can act as sites for further chemical reactions.

Understanding lone pairs can be critical in predicting and explaining molecular properties and their interactions.