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Ch. 2 - General Chemistry Translated: Finding the Electrons
Mullins - Organic Chemistry: A Learner Centered Approach 1st Edition
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All textbooksMullins 1st EditionCh. 2 - General Chemistry Translated: Finding the ElectronsProblem 27c
Chapter 1, Problem 27c

Draw the Lewis structure for the following molecules. Be sure to calculate the formal charge of each atom as a way of confirming your structure is correct.
(c) CF4

Verified step by step guidance
1
Step 1: Identify the central atom in the molecule. In CF₄, the central atom is carbon (C). Carbon typically forms four bonds because it has four valence electrons.
Step 2: Determine the total number of valence electrons in the molecule. Carbon has 4 valence electrons, and each fluorine atom has 7 valence electrons. Since there are 4 fluorine atoms, the total number of valence electrons is: 4 (from C) + 4 × 7 (from F) = 32 valence electrons.
Step 3: Arrange the atoms in the structure. Place the carbon atom in the center and surround it with the four fluorine atoms. Connect each fluorine atom to the carbon atom with a single bond. Each single bond represents 2 electrons, so 8 electrons are used for bonding.
Step 4: Distribute the remaining electrons as lone pairs around the fluorine atoms to satisfy their octet rule. Each fluorine atom needs 8 electrons (including the bonding electrons). After placing 6 lone-pair electrons around each fluorine atom, all 32 valence electrons are accounted for.
Step 5: Calculate the formal charge for each atom to confirm the structure is correct. The formal charge formula is: Formal charge = Valence electrons - (Non-bonding electrons + 1/2 Bonding electrons). For carbon, the formal charge is 0 because it has 4 valence electrons, 0 non-bonding electrons, and 8 bonding electrons. For each fluorine atom, the formal charge is also 0 because they have 7 valence electrons, 6 non-bonding electrons, and 2 bonding electrons.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist. They use dots to represent valence electrons and lines to represent bonds between atoms. Understanding how to draw Lewis structures is essential for visualizing molecular geometry and predicting reactivity.
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Formal Charge

Formal charge is a theoretical charge assigned to an atom in a molecule, calculated based on the number of valence electrons, the number of non-bonding electrons, and the number of bonds. It helps in assessing the stability of a molecule and determining the most favorable Lewis structure. A formal charge of zero on all atoms is often a sign of a stable structure.
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Central Atom Identification

The central atom in a molecule is typically the least electronegative atom that can form multiple bonds, often serving as the core around which other atoms are arranged. Identifying the central atom is crucial for accurately constructing the Lewis structure, as it dictates the overall shape and connectivity of the molecule.
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Related Practice
Textbook Question

Calculate the formal charge of the indicated atom in the following molecules or ions.

(e)

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Textbook Question

Draw the Lewis structure for the following molecules. Be sure to calculate the formal charge of each atom as a way of confirming your structure is correct.

(b) HOBr

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Textbook Question

Using only your intuition, rank the following covalent bonds in terms of their polarity (1 = most polar; 6 = least polar). [You can use the periodic trend of electronegativity, but don't use actual numbers.]

(a) C―C

(b) C―O

(c) C―H

(d) C―F

(e) C―Cl

(f) C―S

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Textbook Question

Calculate the formal charge of the indicated atom in the following molecules or ions.

(f)

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Textbook Question

Without looking at Figure 2.20, use your intuition to estimate whether a bond is ionic, polar covalent, or covalent.

(a) Na―Cl

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Textbook Question

Calculate the formal charge of the indicated atom in the following molecules or ions.

(d)

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