Question

Write the Lewis structure for each ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. \begin{equation}\mathrm{ a.} {PO}_{4}^{3-} \quad \text { b. } \mathrm{CN}^{-} \quad \text { c. } \mathrm{SO}_{3}^{2-} \quad \text { d. } \mathrm{ClO}_{2}^{-}\end{equation}

Short answer

\( {PO}_{4}^{3-} \) has a central P atom with no formal charge and four O atoms (three with a formal charge of -1 and one double-bonded with no formal charge). \( {CN}^{-} \) has a triple bond between C (1+) and N (1-) atoms. \( {SO}_{3}^{2-} \) has a central S atom double-bonded to two O atoms (no charge) and single-bonded to one O atom (2-), with resonance. \( {ClO}_{2}^{-} \) features a central Cl atom single-bonded to two O atoms (one 0 charge, one with 1-), and the Cl has one extra lone pair.

Step-by-step solution

Lewis Structure of \( {PO}_{4}^{3-} \)

First, count the total number of valence electrons: Phosphorus has 5, each Oxygen has 6, and there are 3 extra electrons due to the negative charge, giving a total of 32. Draw a single bond between the P atom and each O atom. Place the remaining electrons to satisfy the octets of the O atoms. The P atom has more than 8 electrons which is permissible because it is in period 3 and can expand its octet. Assign formal charges to each atom.

Lewis Structure of \( {CN}^{-} \)

Count the total number of valence electrons: Carbon has 4, Nitrogen has 5, and there is 1 extra electron due to the negative charge, for a total of 10. Connect the C atom and the N atom with a triple bond and place the remaining electrons to complete the octet for each atom. Assign formal charges to each atom.

Lewis Structure of \( {SO}_{3}^{2-} \)

Total valence electrons are: Sulfur has 6, each Oxygen has 6, and there are 2 extra electrons from the charge for a total of 26. Draw a double bond between the S atom and two O atoms and a single bond to the third O atom. Place the remaining electrons on the O atoms to satisfy their octets and assign formal charges, considering resonance structures where the double bonds can rotate among the O atoms.

Lewis Structure of \( {ClO}_{2}^{-} \)

Count the valence electrons: Chlorine has 7, each Oxygen has 6, and there's 1 extra electron from the charge for a total of 20. Connect Cl atom to each O atom with a single bond and distribute the remaining electrons to fill the octets. Seeing as Cl is a period 3 element, it can have more than 8 electrons. Assign formal charges to each atom.

Key concepts

Resonance Structures

Resonance structures are a representation of a molecule where two or more valid Lewis structures can be drawn for a single molecular formula. These structures differ only in the position of the electrons, not the position of the atoms. Resonance occurs because some molecules have a delocalized electron system, meaning that electrons are not confined to a single bond or atom but are spread over several atoms.

For instance, in the molecule \( {SO}_{3}^{2-} \), resonance structures are essential for describing the molecule’s bonding. The double bonds between sulfur and oxygen atoms can rotate among the oxygen atoms, resulting in multiple Lewis structures that are valid. These structures only depict potential distributions of electrons; the actual molecule is a hybrid of these structures which contribute to the overall stability of the molecule.

Understanding resonance is crucial in grasping the concept of delocalized electrons, which affects properties such as the molecule’s reactivity and stability.

Formal Charge

The formal charge is a theoretical charge on an individual atom within a molecule, calculated by assuming that electrons in chemical bonds are shared equally between atoms. To calculate it, use the formula:

\r\( \text{Formal charge} = (\text{Valence electrons}) - (0.5 \times \text{Bonding electrons}) - (\text{Non-bonding electrons}) \)

The value represents the difference between the number of valence electrons on the isolated atom and the number assigned to it in the molecule. It helps in evaluating the stability of Lewis structures: the most stable structures usually have formal charges closest to zero. Assigning formal charges is part of drawing Lewis structures and is vital for molecules with multiple valid structures, like \( {PO}_{4}^{3-} \) and \( {ClO}_{2}^{-} \) from our exercises.

Valence Electrons

Valence electrons are the electrons in an atom's outermost shell that are available for forming chemical bonds. The number of valence electrons an element has is crucial for determining its chemical properties and how it will bond with other atoms. In Lewis structures, it's necessary to count valence electrons to accurately represent bonding and lone pairs.

Each group in the periodic table has a consistent number of valence electrons; for example, Group 1 elements have one valence electron, while Group 17 elements, like Chlorine, have seven. In the exercises, the valence electron count is the starting point for drawing Lewis structures, as seen with \( {CN}^{-} \) which has 10 valence electrons.

Octet Rule

The octet rule is a chemical rule of thumb that reflects a pattern of atoms forming stable molecules by having eight electrons in their valence shell. This rule explains why atoms form certain chemical compounds: they seek to either gain, lose, or share electrons to achieve a full outer shell, similar to that of noble gases.

Exceptions to this rule occur, especially among molecules with atoms from periods higher than the second in the periodic table, like phosphorus in \( {PO}_{4}^{3-} \), which can have more than eight electrons around it. However, in most cases, following the octet rule will give you a good prediction of the molecule's electron configuration.

Electron Configuration

The electron configuration of an element details how its electrons are distributed among various orbitals. Understanding an atom's electron configuration is essential for predicting its chemical behavior and the nature of its bonds. The electron configuration determines how many valence electrons an atom has and therefore how it might react chemically with other atoms.

Knowing electron configurations also assists in understanding why certain atoms can expand their octet as seen in periods 3 and higher. For instance, sulfur and phosphorus, found in \( {SO}_{3}^{2-} \) and \( {PO}_{4}^{3-} \) respectively, can have more than eight electrons in their valence shell due to available d-orbitals, which are not present in the second period elements.