Chemical Bonding — Basic Concepts: Draw Lewis Structures and Calculate Formal Charge

Question

Iodine forms a series of fluorides. Draw Lewis structures for each of the 4 compounds listed below and determine the formal charge of the iodine atom in each molecule:

  1. IF
  2. IF3
  3. IF5
  4. IF7

 

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Lewis Structures:

  1. Two large letters are connected with a line. The letter "I" is on the left of the line and the letter "F" is on the right of the line. Both letters are surrounded by small filled circles, or dots, grouped in pairs of two. On the letter "I", one pair of circles is on the top, one pair is to the left and one pair is on the bottom. On the letter "F", one pair of circles is on the top, one pair is to the right and one pair is on the bottom.
  2. The letter "I" is centred in the image. To the right of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled ciricle's, or dots. One pair is to the left, one is on the top and one is on the bottom. To the right of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled circles, or dots. One pair is to the right, one is on the top and one is on the bottom. On the bottom of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled circles, or dots. One pair is to the left, one is to the right and one is on the bottom. On the top left of the letter "I" is a pair of small, filled circles, or dots. On the top right of the letter "I" is a pair of small, filled circles, or dots.
  3. The letter "I" is centred in the image and is connected to five lines. Each line is also connected to a letter "F". At the bottom of the letter "I" is a pair of small, filled circles, or dots. Each letter "F"is surrounded by three pairs of small, filled circles, or dots, adding to a total of six small circles around every "F"
  4. The letter "I" is centred in the image and is connected to seven lines. Each line is also connected to a letter "F". Each letter "F" is surrounded by three pairs of small, filled circles, or dots, adding to a total of six small circles around every "F"

Formal Charges on I: 

  1. IF — formal charge on I = 0
  2. IF3 — formal charge on I = 0
  3. IF5 — formal charge on I = 0
  4. IF7 — formal charge on I = 0

Refer to:

Strategy Map

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Check out the strategy map.

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Table 1: Strategy Map
Strategy Map Steps 
1. Count valence electrons in the molecule.

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Valence electrons are the outer shell electrons.
2. Draw the Lewis structures.

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To recall the step for drawing a Lewis structure, refer to Section 4.4: Lewis Symbols and Structures (1).

Show/Hide Steps
  1. Pick your centre atom (if applicable) by selecting the atom that has the lowest electronegativity.
  2. Arrange and connect your atoms. Start with a single bond represented by a single line between connected atoms.
  3. Add lone pairs on terminal atoms until you have used them all.
    Show/Hide Hint

    If you have left over electrons and your centre atom’s octet is full, check and see if that centre atom can have an expanded octet.

  4. Check that you followed the octet rule and that the total number of electrons you used (either as a dot/lone pair or in a line) add up to the total amount you counted in the beginning.
    Show/Hide Hint

    A full octet typically refers to having 8 electrons in an atom’s outermost shell, with some exceptions. An expanded octet is for elements that are in groups 3 and higher.
3. Calculate the formal charge.

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To recall the steps for calculating formal charge, refer to Section 4.5: Formal Charges and Resonance (2).

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  1. Identify the number of valence electrons the atom(s) of interest have.
  2. Using your drawn Lewis structure, identify the number of valence electrons assigned to the atom(s) of interest.
    Show/Hide Hint

    In a Lewis structure, the assigned electrons are its surrounding lone pairs (1 electron per dot and 2 electrons per pair) as well as half of the electrons from each line. (This is because the lines represent shared electrons, so each atom is assigned one from that pair.)

  3. Subtract your assigned electrons from the original valence electrons for each atom of interest. Charges can be positive or negative; 0 means there is no formal charge for that atom.

Solution

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a. IF

Lewis Structure:

(1) Iodine + (1) Fluorine

7ve– + 7ve– = 14e–

Two large letters are connected with a line. The letter "I" is on the left of the line and the letter "F" is on the right of the line. Both letters are surrounded by small filled circles, or dots, grouped in pairs of two. On the letter "I", one pair of circles is on the top, one pair is to the left and one pair is on the bottom. On the letter "F", one pair of circles is on the top, one pair is to the right and one pair is on the bottom.

IF — Formal Changes
Element Iodine
Valence Electrons 7
Assigned Electrons 7
Formal Charges 0

b. IF3

Lewis Structure:

(1) Iodine + (3) Fluorine

7ve– + (3)7ve– = 28e–

The letter "I" is centred in the image. To the right of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled ciricle's, or dots. One pair is to the left, one is on the top and one is on the bottom. To the right of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled circles, or dots. One pair is to the right, one is on the top and one is on the bottom. On the bottom of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled circles, or dots. One pair is to the left, one is to the right and one is on the bottom. On the top left of the letter "I" is a pair of small, filled circles, or dots. On the top right of the letter "I" is a pair of small, filled circles, or dots.

IF3 — Formal Changes
Element Iodine
Valence Electrons 7
Assigned Electrons 7
Formal Charges 0

c. IF5

Lewis Structure:

(1) Iodine + (5) Fluorine

7ve– + (5)7ve– = 42e–

The letter "I" is centred in the image and is connected to five lines. Each line is also connected to a letter "F". At the bottom of the letter "I" is a pair of small, filled circles, or dots. Each letter "F"is surrounded by three pairs of small, filled circles, or dots, adding to a total of six small circles around every "F"

IF5 — Formal Changes
Element Iodine
Valence Electrons 7
Assigned Electrons 7
Formal Charges 0

d. IF7Lewis Structure:

(1) Iodine + (7) Fluorine

7ve– + (7)7ve– = 56e–

The letter "I" is centred in the image and is connected to seven lines. Each line is also connected to a letter "F". Each letter "F" is surrounded by three pairs of small, filled circles, or dots, adding to a total of six small circles around every "F"

IF7 — Formal Changes
Element Iodine
Valence Electrons 7
Assigned Electrons 7
Formal Charges 0

Guided Solution

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The guided solution below will give you the reasoning for each step to get your answer, with reminders and hints.

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Table 2: Guided Solution
Guided Solution Ideas
This question is a Lewis structure application problem. In this problem, you must follow the steps to create 4 different Lewis structures and find their corresponding formal charges on I.

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Refer to:
Note: The Guided Solution detailed below is for question 2 (IF3) only.

To work through creating Lewis structures and calculating formal charges for all species, check out the following interactive video:

If you are using a printed copy, you can scan the QR code with your digital device to go directly to the h5p:

 
Count valence electrons in the molecule by identifying what atoms there are, counting the valence electrons each has, and adding them.

Show/Hide Don’t Forget

To find how many valence electrons an atom has, count across the periodic table (skipping the d-block elements). You can use the group number or electron configuration to find the number of valence electrons in each type of atom.

  • Group 1 = 1
  • Group 2 = 2
  • Group 13 = 3
  • Group 14 = 4
  • Group 15 = 5
  • Group 16 = 6
  • Group 17 = 7
  • Group 18 = 8 (full outer shell)

You cannot have more electrons in a molecule than the amount each of the atoms has combined. For example, since iodine has 7 valence electrons and fluorine has 7 valence electrons, the total number of electrons in the structure of IF will be 14.
Draw the Lewis structure.

Show/Hide Steps
  1. Pick your centre atom (if applicable) by selecting the atom that has the lowest electronegativity.
    Show/Hide Think About This!

    In our case, I has a lower electronegativity, and F can only form 1 bond, so I will be the central atom and will be bonded to each F.

  2. Arrange and connect your atoms. Start by drawing all your atoms and connecting them with 1 line; this line represents a single bond and is your “bare-bones step” showing how the atoms are arranged. Each line represents 2 shared electrons.
  3. Add lone pairs on terminal atoms until you have used them all, or until each atom has a full octet.
  4. In the case that you run out of electrons before all octets are full, make double or triple bonds between sharing atoms.
    Show/Hide Don’t Forget!

    Sometimes, it can be confusing where to add a double bond. If you run out of available electrons and some of your atoms still remain without a full octet, you can add a double bond between 2 or more atoms, as long as it still abides by the octet rule. If you have left over electrons and your centre atom’s octet is full, check and see if that centre atom can have an expanded octet.

  5. In the case that you have remaining electrons, add them to the centre atom until its octet is full.

Check that you followed the octet rule and that the total number of electrons you used (either as a dot/lone pair or in a line) add up to the total amount you counted in the beginning.
Calculate formal charge:

  1. Identify the number of valence electrons the atom(s) of interest have.
  2. Using your drawn Lewis structure, identify the number of valence electrons assigned to the atom(s) of interest.
    Show/Hide Don’t Forget!

    In a Lewis structure, the assigned electrons are its surrounding lone pairs (1 electron per dot and 2 electrons per pair) as well as half of the electrons from each line. (This is because the lines represent shared electrons, so each atom is assigned 1 from that pair.)

  3. Subtract your assigned electrons from the original valence electrons for each atom of interest. Charges can be positive or negative; 0 means there is no formal charge for that atom.
    Show/Hide Think About This!

    The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms.

It can be useful to track valence electrons and calculate formal charges using a table, as shown below, especially if you are working with many atoms.

Show/Hide Table Template
Formal Charges Template
Element [Element Name] [Element Name]
Valence electrons (ve–) – –
Assigned electrons (ae–) – –
Formal Charge ((ve–) − (ae–)) – –
Table 3: Complete Solution
Complete Solution for b. IF3
Count valence electrons:

  1. (1) Iodine + (3) Fluorine
  2. Iodine has 7 valence e–
  3. Each fluorine has 7 valence e–, and we have three F
  4. 7ve– + 3x(7ve–) = 28 total valence electrons
Show/Hide Don’t Forget!

(amount of Iodine atoms)x7ve– + (amount of Fluorine atoms)x7ve– = total valence electrons
Draw Lewis Structure:

  1. Connect I and F by a line representing a single bond. This has used 6 e–, and we have 22 left.
  2. Add lone pairs to each F atom. This adds 3 lone pairs to each F to complete their octet, for a total of 18 more electrons used.
  3. We have 4 electrons left. Add them as 2 lone pairs to the central I. This has used the remaining electrons.
Show/Hide Think About This!

Now, each F atom has a complete octet, so the octet rule is satisfied. We can have an expanded octet since it is in period 5.

Show/Hide Complete Lewis Structure

The letter "I" is centred in the image. To the right of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled ciricle's, or dots. One pair is to the left, one is on the top and one is on the bottom. To the right of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled circles, or dots. One pair is to the right, one is on the top and one is on the bottom. On the bottom of the "I" is a line which Is connected to the letter "F". The letter "F" is surrounded by three pairs of small, filled circles, or dots. One pair is to the left, one is to the right and one is on the bottom. On the top left of the letter "I" is a pair of small, filled circles, or dots. On the top right of the letter "I" is a pair of small, filled circles, or dots.
Formal Charge:

Show/Hide Think About This!

Looking back at the Lewis structure assign 1 electron to I from the single bond, and count up all the electrons in the lone pairs on I.

IF3 — Formal Change
Element Iodine
Valence electrons 7
Assigned electrons 7
Formal Charge(7 – 7) 0

Answer: The formal charge on I in IF3 is 0.

Check Your Work

Summary of what we would expect based on the related chemistry theory.

Show/Hide Watch Out!

The formal charge on a neutral molecule should always be 0. IF3 is neutral, so the 0 formal charge is correct.

Does your answer make chemical sense?

Show/Hide Answer

Lewis structures are a visual representation of the electrons from an atom’s outermost electron shell; they are used to illustrate when these electrions are shared in covalent bonds or present as lone pairs.

Formal charges occur when the number of electrons an atom has is different from the amount it had prior to the reaction. A formal charge can be either positive or negative. It is positive if the atom has less assigned than it originally had and negative if more are assigned. This makes sense as electrons carry a negative charge.

Formal charges can also be used to help decide which of multiple structures is preferred.

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Refer to Section 4.5.2: Using Formal Charges to Predict Molecular Structure (2).

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All figures by the authors (Brewer S. and Blackstock L.) are free to use under a CC0 license.

References

1. OpenStax. 4.4: Lewis Symbols and Structures. In CHEM 1500: Chemical Bonding  and Organic Chemistry. LibreTexts, 2023. https://chem.libretexts.org/Courses/Thompson_Rivers_University/CHEM_1500%3A_Chemical_Bonding_and_Organic_Chemistry/04%3A_Chemical_Bonding_I-_Basic_Concepts/4.04%3A_Lewis_Symbols_and_Structures.

2. OpenStax. 4.5: Formal Charges and Resonance. In CHEM 1500: Chemical Bonding  and Organic Chemistry. LibreTexts, 2023. https://chem.libretexts.org/Courses/Thompson_Rivers_University/CHEM_1500%3A_Chemical_Bonding_and_Organic_Chemistry/04%3A_Chemical_Bonding_I-_Basic_Concepts/4.05%3A_Formal_Charges_and_Resonance.

3. Blackstock, L.; Brewer, S.; Jensen, A. PASS Chemistry Book CHEM 1500; LibreTexts, 2023. https://chem.libretexts.org/Courses/Thompson_Rivers_University/PASS_Chemistry_Book_CHEM_1500.

4. Blackstock, L.; Brewer, S.; Jensen, A. 4.2: Question 4.E.60 PASS – Draw Lewis Structure and Calculate Formal Charge. In PASS Chemistry Book CHEM 1500. LibreTexts, 2023. https://chem.libretexts.org/Courses/Thompson_Rivers_University/PASS_Chemistry_Book_CHEM_1500/04%3A_Chemical_Bonding_I_-_Basic_Concepts/4.02%3A_Question_4.E.60_PASS_-_draw_Lewis_Structure_and_calculate_formal_charge.

5. OpenStax. 7.E: Chemical Bonding and Molecular Geometry (Exercises). In Chemistry 1e (OpenSTAX). LibreTexts, 2023. https://chem.libretexts.org/Bookshelves/General_Chemistry/Chemistry_1e_(OpenSTAX)/07%3A_Chemical_Bonding_and_Molecular_Geometry/7.E%3A_Chemical_Bonding_and_Molecular_Geometry_(Exercises).

6. Flowers, P.; Robinson, W. R.; Langley, R.; Theopold, K. Ch. 6 Exercises. In Chemistry 2e; OpenStax, 2019. https://openstax.org/books/chemistry-2e/pages/7-exercises.

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