Question

Some chemists believe that satisfaction of the octet rule should be the top criterion for choosing the dominant Lewis structure of a molecule or ion. Other chemists believe that achieving the best formal charges should be the top criterion. Consider the dihydrogen phosphate ion, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-},\) in which the \(\mathrm{H}\) atoms are bonded to \(\mathrm{O}\) atoms. (a) What would be the predicted dominant Lewis structure if satisfying the octet rule is the top criterion? (b) What would it be if achieving the best formal charges is the top criterion? (c) Is there another Lewis structure you can draw that satisfies neither of these criteria?

Short answer

The predicted dominant Lewis structure for \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) is Structure A, as it satisfies both the (a) octet rule criterion by having complete octets for all atoms, and (b) achieves the best formal charges, with the lowest number of formal charges closest to zero.

Step-by-step solution

(Step 1: Calculate the total number of valence electrons)

: To build the Lewis structures, we first need to find the total number of valence electrons in the dihydrogen phosphate ion. The valence electrons for each atom are: - H: 1 valence electron (each), total = 1 * 2 = 2 - P: 5 valence electrons - O: 6 valence electrons (each), total = 6 * 4 = 24 Additionally, since the molecule has a -1 charge, we need to add one extra electron, giving a total of 2 + 5 + 24 + 1 = 32 valence electrons.

(Step 2: Draw the possible Lewis structures)

: Now let's build the Lewis structures for the dihydrogen phosphate ion. We can start with a central phosphorus atom bonded to four oxygen atoms, with two of the oxygen atoms also bonded to a hydrogen atom. Next, complete the octets for all atoms using the remaining valence electrons: 1. Structure A: Place the hydrogen atoms on different oxygen atoms (not double-bonded to the phosphorus) 2. Structure B: Place the hydrogen atoms on the same oxygen atom (not double-bonded to the phosphorus) 3. Structure C: Place one hydrogen atom on an oxygen atom double-bonded to the phosphorus (resulting in an oxygen atom with a positive formal charge)

(Step 3: Compare the structures based on the octet rule)

: To decide which Lewis structure satisfies the octet rule, compare the structures to see if they satisfy the octet rule for all atoms in each structure: Structure A and B both satisfy the octet rule for all atoms in the structure. Structure C, however, does not satisfy the octet rule for the oxygen atom with a positive formal charge. Thus, based on the octet rule criterion, the dominant structure for \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) is either Structure A or B.

(Step 4: Compare the structures based on formal charges)

: Now let's compare the structures based on the formal charges for each molecule: 1. Structure A: Formal charges = 0 (P), -1 (O with single bond to H), -1 (O with another single bond to H), 0 (double bonded O's) 2. Structure B: Formal charges = 0 (P), -2 (O with two single bonds to H), 0 (double bonded O's) 3. Structure C: Formal charges = 0 (P), +1 (O with one single bond to H and one double bond to P), -1 (O with single bond to H), 0 (double bonded O) Based on formal charge criterion, Structure A has the best formal charges as it has the lowest number of formal charges closest to zero, which is the more stable configuration. #Conclusion#: The predicted dominant Lewis structure for \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) based on the (a) octet rule criterion is Structure A or B, while (b) based on the best formal charges criterion, it is Structure A. As Structure A is the only structure that satisfies both criteria, it is the most likely dominant Lewis structure for the dihydrogen phosphate ion.

Key concepts

Octet Rule

The octet rule is a cornerstone of chemical bonding theory. It is based on the premise that atoms tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas and, consequently, a stable structure.

When chemists predict the dominant Lewis structure of a molecule or ion, satisfying the octet rule is often a priority. An atom that does not possess a complete octet is typically more reactive and less stable. For example, in the case of the dihydrogen phosphate ion (\(\mathrm{H}_2\mathrm{PO}_4^{-}\)), the preferred Lewis structures would be those where the phosphorus and oxygen atoms all satisfy the octet rule, meaning eight electrons are surrounding these atoms.

The exercise asked to identify a dominant structure according to the octet rule. Both Structures A and B satisfy this rule for all atoms, making them potential dominant structures. However, another structure, Structure C, does not satisfy the octet rule for an oxygen atom, as it does not have eight electrons around it due to a double bond with phosphorus and a bond to a hydrogen atom, which leads to a formal charge.

This highlights the importance of the octet rule as a guiding principle in determining the shape and stability of a molecule's Lewis structure.

Formal Charges

Formal charge is a concept used to estimate the distribution of electric charge within a molecule. The formal charge on an atom is the theoretical charge it would have if all atoms in the molecule had the same electronegativity; it is calculated based on the number of valence electrons on the free atom, the number of non-bonding electrons on the atom in the molecule, and the number of bonds the atom participates in within the molecule.

To compute the formal charge, the formula used is:

\[\text{Formal Charge} = (\text{Valence Electrons}) - \frac{1}{2}(\text{Bonding Electrons}) - (\text{Non-bonding Electrons})\]

When determining the dominant Lewis structure for \(\mathrm{H}_2\mathrm{PO}_4^{-}\), considering the best formal charges can lead to a different conclusion than when applying the octet rule criteria. For instance, according to formal charge considerations, Structure A is more favorable, as it has the charges spread more evenly across the molecule and avoids highly charged atoms.

The comparison based on formal charges is valuable because it helps to identify structures with charges closest to zero. Such structures are generally more stable, as they have a distribution of charge that is closer to the actual distribution in the molecule. The formal charge analysis is particularly important for structures where the octet rule doesn't apply, such as in molecules with an odd number of electrons or expanded valence shells.

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom. They play a critical role in chemical bonding and reactions because they are the electrons involved in bond formation, either as shared or transferred electrons. When drawing Lewis structures, the total number of valence electrons in all of the atoms gives an indication of the bonding possibilities.

In the exercise with the dihydrogen phosphate ion (\(\mathrm{H}_2\mathrm{PO}_4^{-}\)), calculating the number of valence electrons is the starting point for predicting potential structures. By knowing that hydrogen has 1 valence electron, phosphorus has 5, and oxygen has 6, we can determine that the \(\mathrm{H}_2\mathrm{PO}_4^{-}\) ion has 32 valence electrons in total to be placed in the Lewis structure.

Placing these valence electrons in a way that satisfies both the octet rule and assigns formal charges close to zero is a thoughtful way to approach the construction of a Lewis structure. It emphasizes the balancing act between fulfilling the requirement for a full valence shell (octet rule) and ensuring individual atoms within a molecule have the lowest possible formal charges for stability and reactivity—a reflection of the true nature of chemical bonding.

Understanding the behavior and interaction of these valence electrons is crucial not just for homework exercises but also for grasping fundamental chemistry concepts such as bonding, molecular geometry, and reactivity.