Problem 8 - Chemistry

Question

Identify the correct electron-dot structure(s) for phosphate, \(\mathrm{PO}_{4}{\underline{\phantom{xx}}}^{3-}\). Explain what is wrong with the incorrect ones. Is it reasonable to have more than one correct electrondot structure? Explain.

Step-by-Step Solution

Answer

Phosphate has more than one correct electron-dot structure due to resonance.

Step 1: Determine Total Valence Electrons

First, calculate the total number of valence electrons for the molecule. Phosphorus (P) has 5 valence electrons, oxygen (O) has 6 valence electrons each, and there are four oxygens. Since the ion has a 3- charge, add 3 extra electrons: \(5 + (6 \times 4) + 3 = 32\) valence electrons.

Step 2: Draw the Basic Skeleton Structure

Draw phosphorus in the center with four oxygens surrounding it because phosphorus is less electronegative than oxygen. Connect each oxygen to phosphorus with a single bond.

Step 3: Distribute Remaining Electrons

After the single bonds, distribute the remaining 24 electrons (since each bond counts as 2 electrons: \(32 - (4 \times 2) = 24\)) to complete the octets of the oxygen atoms. Each oxygen needs 6 more electrons to complete its octet.

Step 4: Check and Assign Formal Charges

Calculate the formal charges for each atom in the molecule. With all single bonds, phosphorus will have a formal charge of +1, each oxygen a formal charge of -1, and the overall charge is -3, matching the ion's charge. Consider making double bonds to reduce formal charges: move one pair of electrons from two oxygens to form P=O bonds.

Step 5: Evaluate Multiple Structures

Draw the alternate structure where two oxygens have double bonds with phosphorus, meaning two oxygens have formal charges of 0, phosphorus has formal charge 0, and the other two oxygens have a formal charge of -1 each, maintaining the correct ion charge of -3. This is a resonance structure.

Step 6: Determine Correct Structures and Explain

Both structures where phosphorus forms single bonds with all oxygens and where it forms two double bonds with two oxygens are valid; they are resonance structures. They balance the overall charge while reducing formal charge on individual atoms. Other structures that move beyond this equivalency may violate the preferred charge distribution or octet rule.

Key Concepts

Valence Electrons

Understanding valence electrons is crucial in determining the electron-dot structure of a molecule. They are the outermost electrons involved in chemical bonding. For phosphate, \(\mathrm{PO}_{4}^{3-}\), we start by calculating the total number of valence electrons available:

• Phosphorus (P) has 5 valence electrons.
• Each oxygen (O) has 6 valence electrons, and there are four oxygens.
• Add 3 extra electrons due to the 3- charge on the ion.

This gives us:\[5 + (6 \times 4) + 3 = 32 \text{ valence electrons}.\]These electrons determine how atoms will connect and share electrons in bonds. Knowing the total count helps us draw the basic skeletal structure, ensuring each atom achieves a stable electron configuration.

Formal Charge

Formal charge helps in assessing the best electron-dot structure by minimizing charges on atoms, contributing to overall stability. The formula to calculate formal charge for an atom is:\[\text{Formal Charge} = \text{Valence Electrons} - \text{Non-bonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons}.\]For phosphate, initially:

• Phosphorus has a formal charge of +1.
• Each oxygen with a single bond has a formal charge of -1.

The molecule's overall charge aligns with the ion's 3- charge.Adjusting for better structures, you can move electrons from two oxygens to form double bonds with phosphorus:

• Phosphorus then has a formal charge of 0.
• Two oxygens with double bonds also have a charge of 0.
• The remaining two oxygens stay at -1.

This balancing act results in a more stable configuration, distributing charges more effectively.

Resonance Structure

Resonance structures propose multiple configurations for a molecule that can't be represented by a single electron-dot structure alone. For phosphate, two main structures exist:

• One with all single P-O bonds.
• Another where two of the P-O bonds are double bonds.

The true form of the molecule is a hybrid, reflecting all valid structures. Both maintain the correct overall charge of -3, but they offer different distributions of formal charges. Resonance helps in depicting electron delocalization which is more realistic. Such structures can stabilize molecules by blending characteristics, making them energetically favorable. Acknowledging resonance recognizes the dynamic nature of electron distribution, key in predicting molecular behavior.