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. \(\mathrm{C}_{2} \mathrm{H}_{6}\) b. \(\mathrm{NF}_{3}\) c. \(\mathrm{C}_{4} \mathrm{H}_{10}\) d. \(\mathrm{SiCl}_{4}\)

Short answer

The Lewis structures for the given molecules are: a. Ethane \(\mathrm{C}_{2} \mathrm{H}_{6}\): \[\mathrm{H}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{H}\] \[|\hspace{10px}|\] b. Nitrogen Trifluoride \(\mathrm{NF}_{3}\): \[..\] \[\mathrm{F}\text{—}\mathrm{N}\text{—}\mathrm{F}\] \[..\hspace{12px}..\hspace{12px}..\] c. Butane \(\mathrm{C}_{4} \mathrm{H}_{10}\): \[\mathrm{H}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{H}\] \[|\hspace{10px}|\hspace{10px}|\hspace{10px}|\] d. Silicon Tetrachloride \(\mathrm{SiCl}_{4}\): \[\mathrm{Cl}\text{—}\mathrm{Si}\text{—}\mathrm{Cl}\] \[|\hspace{18px}|\]

Step-by-step solution

a. Lewis Structure for \(\mathrm{C}_{2} \mathrm{H}_{6}\) (Ethane)#

1. Counting Valence Electrons: Carbon has 4 valence electrons, and hydrogen has 1 valence electron. There are 2 carbon atoms and 6 hydrogen atoms, giving: \[2 \times 4 + 6 \times 1 = 14\text{ valence electrons.}\] 2. Central Atom: Carbon (C) is the central atom in this molecule. 3. Arrange Atoms: Place the carbon atoms in the center and the hydrogen atoms around them: \[\mathrm{H}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{H}\] 4. Distribute Electrons: Each carbon atom needs four bonds, and each hydrogen atom needs one bond. Assign the remaining 10 electrons (5 electron pairs) as single bonds between carbon and hydrogen atoms: \[\mathrm{H}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{H}\] \[|\hspace{10px}|\] The Lewis structure for \(\mathrm{C}_{2} \mathrm{H}_{6}\) (ethane) is: \[\mathrm{H}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{H}\] \[|\hspace{10px}|\]

b. Lewis Structure for \(\mathrm{NF}_{3}\) (Nitrogen Trifluoride)#

1. Counting Valence Electrons: Nitrogen has 5 valence electrons, and fluorine has 7 valence electrons. There is 1 nitrogen atom and 3 fluorine atoms, giving: \[ 5 + 3 \times 7 = 26\text{ valence electrons.}\] 2. Central Atom: Nitrogen (N) is the central atom in this molecule. 3. Arrange Atoms: Place the nitrogen atom in the center and the fluorine atoms around it: \[\mathrm{F}\text{—}\mathrm{N}\text{—}\mathrm{F}\] \[|\hspace{18px}|\] 4. Distribute Electrons: Assign two electrons (one electron pair) for bonding between the N and each F atom, and add the remaining electrons as lone pairs on N and F atoms. The Lewis structure for \(\mathrm{NF}_{3}\) (nitrogen trifluoride) is: \[..\] \[\mathrm{F}\text{—}\mathrm{N}\text{—}\mathrm{F}\] \[..\hspace{12px}..\hspace{12px}..\]

c. Lewis Structure for \(\mathrm{C}_{4} \mathrm{H}_{10}\) (Butane)#

1. Counting Valence Electrons: Carbon has 4 valence electrons, and hydrogen has 1 valence electron. There are 4 carbon atoms and 10 hydrogen atoms, giving: \[4 \times 4 + 10 \times 1 = 26\text{ valence electrons.}\] 2. Central Atoms: Carbon (C) atoms are linked sequentially to form the carbon chain. 3. Arrange Atoms: Place the carbon atoms in a chain and the hydrogen atoms around them: \[\mathrm{H}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{H}\] \[|\hspace{10px}|\hspace{10px}|\hspace{10px}|\] 4. Distribute Electrons: Assign two electrons (one electron pair) for bonding between adjacent carbon atoms and between carbon and hydrogen atoms. The Lewis structure for \(\mathrm{C}_{4} \mathrm{H}_{10}\) (butane) is: \[\mathrm{H}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{C}\text{—}\mathrm{H}\] \[|\hspace{10px}|\hspace{10px}|\hspace{10px}|\]

d. Lewis Structure for \(\mathrm{SiCl}_{4}\) (Silicon Tetrachloride)#

1. Counting Valence Electrons: Silicon has 4 valence electrons, and chlorine has 7 valence electrons. There is 1 silicon atom and 4 chlorine atoms, giving: \( 4 + 4 \times 7 = 32\text{ valence electrons.}\) 2. Central Atom: Silicon (Si) is the central atom in this molecule. 3. Arrange Atoms: Place the silicon atom in the center and the chlorine atoms around it: \[\mathrm{Cl}\text{—}\mathrm{Si}\text{—}\mathrm{Cl}\] \[|\hspace{18px}|\] 4. Distribute Electrons: Assign two electrons (one electron pair) for bonding between the Si and each Cl atom, and add the remaining electrons as lone pairs on Cl atoms. The Lewis structure for \(\mathrm{SiCl}_{4}\) (silicon tetrachloride) is: \[\mathrm{Cl}\text{—}\mathrm{Si}\text{—}\mathrm{Cl}\] \[|\hspace{18px}|\]

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom that are involved in chemical bonding. They are key to understanding how molecules form bonds and reach stability. Each element's atoms have a unique number of valence electrons, depending on which column or group they belong to in the periodic table. For example, carbon atoms, found in group 14 of the periodic table, have four valence electrons. This explains why carbon forms four covalent bonds, typically with other elements that complete its octet.

When drawing a Lewis structure, the first step involves counting all valence electrons in the molecule. For instance, the ethane molecule (\(\mathrm{C}_{2} \mathrm{H}_{6}\)) requires calculating the valence electrons of both carbon and hydrogen. Carbon contributes four electrons per atom, while hydrogen contributes one, leading to a sum of fourteen valence electrons in total for \(\mathrm{C}_{2} \mathrm{H}_{6}\). Identifying and counting these electrons correctly ensures that the molecular structure respects the correct number of electrons, either shared or as lone pairs.

Molecular Structure

Molecular structure pertains to the 3D arrangement of atoms within a molecule. The Lewis structure is a simple representation showing how atoms are bonded, but it also reflects insights into electron distribution which affects molecular shape and properties. At its core, molecular structure involves understanding which atoms are central and which are peripheral. Normally, the atom with the fewest valence electrons or higher tendency to form multiple bonds acts as a central atom.

Taking nitrogen trifluoride (\(\mathrm{NF}_{3}\)) as an example, nitrogen serves as the central atom surrounded by fluorine atoms. This arrangement aligns with nitrogen's ability to form three bonds, matching each fluorine atom. In the Lewis structure for \(\mathrm{NF}_{3}\), the N atom connects to three F atoms, with lone pairs left on nitrogen, influencing the electron geometry to a trigonal pyramidal shape. This structure illustrates the basic 3D shape we expect, affecting properties such as boiling point and polarity.

Bonding Pairs

Bonding pairs refer to electrons shared between atoms, signified by lines in a Lewis structure. When atoms bond, they tend to share enough electrons to fill their valence shells, achieving stability, often forming single, double, or triple covalent bonds. For example, single bonds involve one pair of shared electrons, as seen between carbon and hydrogen in ethane (\(\mathrm{C}_{2} \mathrm{H}_{6}\)), resulting in the simple depiction: \(\mathrm{H}\text{—}\mathrm{C}\).

In contrast, nonbonding pairs, or lone pairs, are unshared and usually sit on the peripheral atoms, impacting molecular geometry. When drafting the Lewis structure for molecules like \(\mathrm{SiCl}_{4}\), these lone pairs and bonding pairs dictate whether an atom forms a tetrahedral shape or maintains linearity. Silicon’s central atom bonds with four chlorine atoms, where each bond is a shared pair of electrons. These shared electrons count as bonding pairs, determining the sturdy tetrahedral structure in the Lewis representation.