Chapter 8: Problem 50
(a) Draw the dominant Lewis structure for the phosphorus trifluoride molecule, PF \(_{3}\) . (b) Determine the oxidation numbers of the \(P\) and \(F\) atoms. (c) Determine the formal charges of the \(P\) and \(F\) atoms.
Short Answer
Expert verified
The dominant Lewis structure of PF$_3$ is:
F
|
P - F - F
|
F
The oxidation numbers of the P and F atoms are -1 and +1, respectively. The formal charges of both P and F atoms are 0.
Step by step solution
01
Determine the valence electrons for P and F atoms
To draw the Lewis structure of PF3, first determine the number of valence electrons for each atom. Phosphorus (P) is in Group 15 (or Group 5A) and has 5 valence electrons. Fluorine (F) is in Group 17 (Group 7A) and has 7 valence electrons.
02
Connect the atoms and distribute the valence electrons
Connect the P atom to three F atoms with a single bond (two valence electrons) to each F atom. This will use 6 of the 5 valence electrons of P. Leave the remaining 3 F atoms' valence electrons as lone pairs. Each F atom will now have two lone pairs (four valence electrons), and then we need to place the remaining two electrons on the P atom to complete its octet.
03
Complete the Lewis structure
Draw the additional lone pair on the P atom. The final Lewis structure for PF3 is as follows:
F
|
P - F - F
|
F
04
Determine the oxidation numbers of P and F atoms
To calculate the oxidation numbers, use the following rules:
1. The oxidation number of an atom in a free element is 0.
2. For monoatomic ions, the oxidation number is equal to the charge of the ion.
3. The sum of the oxidation numbers in a neutral species is 0.
In this case, since P is more electronegative than F, P will have a -1 oxidation number while F will have a +1 oxidation number.
Oxidation numbers: P = -1, F = +1
05
Determine the formal charges of P and F atoms
To calculate the formal charges, follow the formula:
Formal Charge = (Valence Electrons) - (0.5 x Bonding Electrons) - (Lone-pair Electrons)
For the P atom:
Formal Charge = (5) - (0.5 x 6) - (2) = 0
For the F (Any single) atom:
Formal Charge = (7) - (0.5 x 2) - (6) = 0
Formal charges: P = 0; F = 0
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Oxidation Numbers
Oxidation numbers are a useful concept in understanding chemical bonding and the transfer of electrons between atoms in a molecule. These numbers help in determining how electrons are distributed in compounds. The oxidation number, sometimes called oxidation state, indicates the degree of oxidation of an atom. Here are the fundamental rules used to find oxidation numbers:
- The oxidation number of a free element is always 0. For example, in diatomic oxygen (\(O_2\)) or elemental phosphorus (\(P\)), each atom's oxidation number is 0.
- In a neutral molecule, like \(\text{PF}_3\), the sum of oxidation numbers of all atoms equals 0.
- The oxidation number of fluorine (F) is -1 in compounds because it is highly electronegative.
- Phosphorus (P) is generally less electronegative than fluorine in its compounds and typically has varying oxidation states depending on the compound's nature.
Formal Charges
Formal charges are used to determine the distribution of electrons in a molecule. They are useful for predicting the most stable Lewis structure or resonance form for a molecule. Calculating formal charges involves the following formula:Formal Charge = (Valence Electrons) - (0.5 x Bonding Electrons) - (Lone pair Electrons)This helps identify the charges "felt" by each atom if the electrons were equally shared, making it a simple bookkeeping tool within Lewis structures.
- Starting with phosphorus (\(P\)) in \(\text{PF}_3\): It has 5 valence electrons, is involved in 3 bonds with the fluorine atoms, and has one lone pair. Thus, \( 5 - 0.5 \times 6 - 2 = 0 \), giving a formal charge of 0 for phosphorus.
- For each fluorine (\(F\)) atom: It brings 7 valence electrons and, forming a single bond with phosphorus, involves in 2 shared electrons and 3 lone pairs. Hence, \( 7 - 0.5 \times 2 - 6 = 0 \), leaving each fluorine with a formal charge of 0.
Phosphorus Trifluoride
Phosphorus trifluoride (\(\text{PF}_3\)) is a fascinating compound with a trigonal pyramidal shape, which exhibits unique chemical behaviors due to its structure and electron distribution. This molecule is noteworthy because of how phosphorus unexpectedly takes on a central role in bonding, even though it is less electronegative compared to fluorine.**Structure and Bonding**:- The molecule consists of one phosphorus (P) atom in the center, bonded to three fluorine (F) atoms arranged around it. - Its trigonal pyramidal shape results from the P atom having a lone pair, which reduces symmetry.
The lone pair on phosphorus pushes the fluorines downwards, slightly distorting the angle to less than the ideal 109.5° of a tetrahedron. This contributes to unique physical properties and reactivities.**Properties and Use**:- Phosphorus trifluoride is colorless and toxic, typically used as a precursor in organophosphorus chemistry.- Its stability and reactivity hinge on the strong P-F bond and the molecule's geometry.- The compound finds use in industrial applications, particularly in the synthesis of various organophosphate compounds.Understanding phosphorus trifluoride and its behavior helps chemists grasp the subtleties of molecular geometry, electron sharing, and reactivity in complex chemical systems.
The lone pair on phosphorus pushes the fluorines downwards, slightly distorting the angle to less than the ideal 109.5° of a tetrahedron. This contributes to unique physical properties and reactivities.**Properties and Use**:- Phosphorus trifluoride is colorless and toxic, typically used as a precursor in organophosphorus chemistry.- Its stability and reactivity hinge on the strong P-F bond and the molecule's geometry.- The compound finds use in industrial applications, particularly in the synthesis of various organophosphate compounds.Understanding phosphorus trifluoride and its behavior helps chemists grasp the subtleties of molecular geometry, electron sharing, and reactivity in complex chemical systems.