Question

Write a Lewis structure for each of the following simple molecules. Show all bonding valence electron pairs as lines and all nonbonding valence electron pairs as dots. a. \(\mathrm{H}_{2}\) b. HCl c. \(\mathrm{CF}_{4}\) d. \(\mathrm{C}_{2} \mathrm{F}_{6}\)

Short answer

The Lewis structures for the given simple molecules are: a. \(\mathrm{H}_{2}\): H-H b. HCl: H-Cl with 3 electron pairs around Cl c. \(\mathrm{CF}_{4}\): Carbon in the center, surrounded by four fluorine atoms, each single-bonded to the carbon, and each with 3 electron pairs around it d. \(\mathrm{C}_{2} \mathrm{F}_{6}\): Two carbon atoms in the center connected by a single bond, each carbon atom is surrounded by three fluorine atoms, each single-bonded to the carbon, and each with 3 electron pairs around it

Step-by-step solution

Determine the number of valence electrons for each element

First, you need to find the total number of valence electrons for each molecule. To do this, look up each element's group number on the periodic table and use that as the number of valence electrons. a. For \(\mathrm{H}_{2}\), each hydrogen atom belongs to group 1, so it has 1 valence electron, and there are 2 hydrogen atoms in total. b. For HCl, hydrogen has 1 valence electron, and chlorine belongs to group 7, so it has 7 valence electrons. c. For \(\mathrm{CF}_{4}\), carbon belongs to group 4, so it has 4 valence electrons, and fluorine belongs to group 7, so it has 7 valence electrons. There are 4 fluorine atoms in total. d. For \(\mathrm{C}_{2} \mathrm{F}_{6}\), there are 2 carbon atoms and 6 fluorine atoms in total.

Determine the central atom and arrange the other atoms around it

The central atom is usually the one with the lowest electronegativity, which is not a terminal atom. a. For \(\mathrm{H}_{2}\), there is no central atom since there are only two hydrogen atoms. b. For HCl, there is no central atom since there are only two atoms: hydrogen and chlorine. c. For \(\mathrm{CF}_{4}\), the central atom is carbon, and the four fluorine atoms will surround it. d. For \(\mathrm{C}_{2} \mathrm{F}_{6}\), there are two central carbon atoms connected to each other, and the six fluorine atoms will surround the carbon atoms.

Place bonding pairs of electrons between the atoms

Create single, double, or triple bonds between atoms by placing bonding pairs of electrons as lines between them. a. For \(H_2\), create a single bond between the two hydrogen atoms. b. For HCl, create a single bond between the hydrogen and chlorine atoms. c. For \(\mathrm{CF}_{4}\), create single bonds between the carbon atom and the four fluorine atoms. d. For \(\mathrm{C}_{2} \mathrm{F}_{6}\), create a single bond between the two carbon atoms, and single bonds between each carbon atom and the three fluorine atoms attached to it.

Add nonbonding electron pairs

Now distribute the remaining valence electrons as nonbonding electron pairs, represented by dots, around each atom. Keep in mind the octet rule which states that most atoms tend to hold 8 electrons in their outermost shell. a. For \(H_2\), both hydrogen atoms share their 1 valence electron, so there are no nonbonding electron pairs. b. For HCl, the chlorine atom has six nonbonding electrons, since it shares one electron with hydrogen. c. For \(CF_4\), each fluorine atom has six nonbonding electrons, as they share one electron with carbon. d. For \(C_2 F_6\), each fluorine atom has six nonbonding electrons, and each carbon atom has no nonbonding electron pairs since they share two electrons.

Draw the final Lewis structures

With all the bonding and nonbonding electron pairs in place, draw the Lewis structures for each molecule: a. \(\mathrm{H}_{2}\): H-H b. HCl: H-Cl with 3 electron pairs around Cl c. \(CF_4\): Carbon in the center, surrounded by four fluorine atoms, each single-bonded to the carbon, and each with 3 electron pairs around it d. \(C_2 F_6\): Two carbon atoms in the center connected by a single bond, each carbon atom is surrounded by three fluorine atoms, each single-bonded to the carbon, and each with 3 electron pairs around it Now you have drawn the Lewis structures for each of the given simple molecules!

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a key role in bonding and chemical reactions. These electrons are found in the highest energy level (or shell), and they are the ones involved in forming chemical bonds. Identifying valence electrons is the first step when drawing Lewis structures since they determine how atoms connect in a molecule.
To find the number of valence electrons, look at the element’s group number on the periodic table. For example, hydrogen, which is in group 1, has one valence electron. Chlorine, in group 17, has seven. Understanding these basics helps in predicting how atoms bond to form stable molecules.
Valence electrons are crucial because they allow for the calculation of total available electrons that can be used to form bonds, helping determine the structure and properties of molecules.

Bonding Pairs

Bonding pairs of electrons are electron pairs shared between atoms to form chemical bonds. These shared pairs are represented by lines in Lewis structures. For instance, in the molecule \( ext{H}_2 \), each hydrogen contributes one valence electron to form a single bond. Thus, two electrons (one from each hydrogen atom) create a bonding pair.
In terms of representation, if you see a line in a Lewis structure, that's a bonding pair. These pairs are fundamental to understanding how molecular structures are formed. The number of bonds an atom can form is determined by the number of its unpaired valence electrons.
Bonding pairs are the key to building stable molecular structures. They significantly affect the molecular geometry and determine how molecules interact chemically and physically.

Nonbonding Electron Pairs

Nonbonding electron pairs, also known as lone pairs, are valence electrons not involved in chemical bonds. These are indicated by dots in Lewis structures and are located around the central atom or atoms. In the \( ext{HCl} \) molecule, for example, chlorine has three pairs of nonbonding electrons.
Lone pairs play an important role in determining the shape and polarity of a molecule. They can influence the angle between bonded atoms, affecting the molecule's overall geometry. While they don't form bonds, they occupy space around the atom, impacting electron distribution and repulsion within the molecule.
This concept helps explain why molecules have different shapes and properties, such as how water's bent shape leads to its polar nature, giving it distinct chemical behaviors.

Octet Rule

The octet rule is a principle that many atoms follow to achieve stability when forming molecules. It states that atoms tend to bond in such a way that they have eight electrons in their valence shell, resembling the electron configuration of a noble gas. This rule guides how atoms will bond and share electrons.
For example, carbon, which has four valence electrons, will form bonds to share electrons and achieve a total of eight in its outer shell. In the molecule \( ext{CF}_4 \), carbon shares electrons with four fluorine atoms to fulfill this requirement.
While the octet rule is applicable to many elements, there are exceptions, such as elements in period 3 and beyond, which can have more than eight valence electrons due to available d orbitals. Understanding the octet rule is essential for predicting the formation and structure of most compounds.