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Chapter 8: Problem 50

(a) Draw the dominant Lewis structure for the phosphorus trifluoride molecule, \(\mathrm{PF}_{3}\). (b) Determine the oxidation numbers of the \(\mathrm{P}\) and \(\mathrm{F}\) atoms. \((\mathbf{c})\) Determine the formal charges of the \(\mathrm{P}\) and \(\mathrm{F}\) atoms.

Short Answer

Expert verified
The dominant Lewis structure for $\mathrm{PF}_{3}$ is: P - single bond - F | F - single bond - P - single bond - F With 6 lone pair electrons around each fluorine atom. The oxidation number of P is +3, and the oxidation number of each F is -1. The formal charge for the P atom is -1 and the formal charge for each F atom is 0.

Step by step solution

01

Identify the valence electrons in P and F atoms

First, let's determine the number of valence electrons in phosphorus (P) and fluorine (F) atoms. Phosphorus, located in group 15, has 5 valence electrons while fluorine, located in group 17, has 7 valence electrons.
02

Calculate the total number of valence electrons

Now, we need to calculate the total number of valence electrons in PF3. Total valence electrons = 1 (Number of P atoms) x 5 (Valence electrons of P) + 3 (Number of F atoms) x 7 (Valence electrons of F) Total valence electrons = 5 + 21 = 26
03

Draw the base Lewis structure

Place the phosphorus (P) atom in the center and form single covalent bonds with the three fluorine (F) atoms. Since 3 single bonds are formed, 6 valence electrons are used up, leaving 20 valence electrons for putting on the surrounding fluorine atoms.
04

Distribute the remaining valence electrons

We need to distribute the remaining valence electrons in order to complete the octet for all the atoms. Since each fluorine atom already has 2 valence electrons from the single bond, we need to add 6 more valence electrons to each fluorine atom to complete the octet.
05

Draw the dominant Lewis structure

Now, the Lewis structure for PF3 is as follows: P - single bond - F | F - single bond - P - single bond - F With 6 lone pair electrons around each fluorine atom.
06

Determine the oxidation numbers

To find the oxidation numbers for the P and F atoms, we must follow some rules: - The oxidation number of F is always -1. - The sum of oxidation numbers of all atoms in a molecule is equal to the net charge on the molecule. Let x be the oxidation number of P. x + (-1) × 3 = 0 x - 3 = 0 x = 3 The oxidation number of P is +3, and the oxidation number of each F is -1.
07

Determine the formal charges

To calculate the formal charge of an atom, we will use the formula: Formal charge = Valence electrons (neutral atom) - Non-bonded electrons - (1/2) × Bonding electrons Formal charge of P: = 5 (Valence electrons) - 0 (Non-bonded electrons) - (1/2 × 6) (Bonding electrons) = 5 - 0 - 3 = 2 - 3 = -1 Formal charge of each F: = 7 (Valence electrons) - 6 (Non-bonded electrons) - (1/2 × 2) (bonding electrons) = 7 - 6 - 1 = 0 Thus, the formal charge for P is -1, and the formal charge for each F is 0.

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Most popular questions from this chapter

The \(\mathrm{Ti}^{2+}\) ion is isoelectronic with the Ca atom. (a) Write the electron configurations of \(\mathrm{Ti}^{2+}\) and Ca. (b) Calculate the number of unpaired electrons for Ca and for \(\mathrm{Ti}^{2+}\). (c) What charge would Ti have to be isoelectronic with \(\mathrm{Ca}^{2+}\) ?

In the vapor phase, \(\mathrm{BeCl}_{2}\) exists as a discrete molecule. (a) Draw the Lewis structure of this molecule, using only single bonds. Does this Lewis structure satisfy the octet rule? (b) What other resonance structures are possible that satisfy the octet rule? (c) On the basis of the formal charges, which Lewis structure is expected to be dominant for \(\mathrm{BeCl}_{2}\) ?

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Although \(\mathrm{I}_{3}^{-}\) is a known ion, \(\mathrm{F}_{3}^{-}\) is not. \((\mathbf{a})\) Draw the Lewis structure for \(I_{3}^{-}\) (it is linear, not a triangle). (b) One of your classmates says that \(\mathrm{F}_{3}^{-}\) does not exist because \(\mathrm{F}\) is too electronegative to make bonds with another atom. Give an example that proves your classmate is wrong. (c) Another classmate says \(\mathrm{F}_{3}^{-}\) does not exist because it would violate the octet rule. Is this classmate possibly correct? (d) Yet another classmate says \(\mathrm{F}_{3}^{-}\) does not exist because \(\mathrm{F}\) is too small to make bonds to more than one atom. Is this classmate possibly correct?

(a) Draw the Lewis structure for hydrogen peroxide, $\mathrm{H}_{2} \mathrm{O}_{2}$. (b) What is the weakest bond in hydrogen peroxide? (c) Hydrogen peroxide is sold commercially as an aqueous solution in brown bottles to protect it from light. Calculate the longest wavelength of light that has sufficient energy to break the weakest bond in hydrogen peroxide.
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