Question

Draw Lewis structures for the following: (a) \(\mathrm{SiH}_{4}\), (b) \(\mathrm{CO}\), (c) \(\mathrm{SF}_{2}\), (d) \(\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{H}\) is bonded to \(\mathrm{O})\), (e) \(\mathrm{ClO}_{2}^{-}\) (f) \(\mathrm{NH}_{2} \mathrm{OH}\).

Short answer

The Lewis structures for the given molecules and ions are as follows: a) \(\mathrm{SiH}_{4}\): ``` H | Si - H | H ``` b) \(\mathrm{CO}\): ``` O ≡ C ``` c) \(\mathrm{SF}_{2}\): ``` F \ S - F / F ``` d) \(\mathrm{H}_{2} \mathrm{SO}_{4}\): ``` O || O-S-O || O-H ``` e) \(\mathrm{ClO}_{2}^{-}\): ``` O \ Cl - O / O ``` f) \(\mathrm{NH}_{2} \mathrm{OH}\): ``` O | N - H | H ```

Step-by-step solution

a) \(\mathrm{SiH}_{4}\)

Step 1: Count total valence electrons Silicon has 4 valence electrons, and each hydrogen atom has 1 valence electron. The total number of valence electrons for the molecule is: \(4 + 4(1) = 8\). Step 2: Arrange atoms Place the silicon atom in the center and arrange the four hydrogen atoms around it. Step 3: Distribute electrons We have 8 electrons to distribute. Start by forming single bonds between the silicon and each hydrogen atom, which will use 8 electrons. The Lewis structure for \(\mathrm{SiH}_{4}\) is: ``` H | Si - H | H ```

b) \(\mathrm{CO}\)

Step 1: Count total valence electrons Carbon has 4 valence electrons, and oxygen has 6 valence electrons. The total number of valence electrons for the molecule is: \(4 + 6 = 10\). Step 2: Arrange atoms Carbon and oxygen will be bonded together. Step 3: Distribute electrons First, form a single bond between carbon and oxygen, using 2 electrons. Then place 6 more electrons around oxygen. We still need to place 2 more electrons, which should be placed as a double bond between carbon and oxygen. The Lewis structure for \(\mathrm{CO}\) is: ``` O ≡ C ```

c) \(\mathrm{SF}_{2}\)

Step 1: Count total valence electrons Sulfur has 6 valence electrons, and each fluorine atom has 7 valence electrons. The total number of valence electrons for the molecule is: \(6 + 2(7) = 20\). Step 2: Arrange atoms Place the sulfur atom in the center and arrange the two fluorine atoms around it. Step 3: Distribute electrons Form single bonds between the sulfur and each fluorine atom, and distribute remaining electrons pairs according to the octet rule. The Lewis structure for \(\mathrm{SF}_{2}\) is: ``` F \ S - F / F ```

d) \(\mathrm{H}_{2} \mathrm{SO}_{4}\)

Step 1: Count total valence electrons Sulfur has 6 valence electrons, oxygen has 6 valence electrons, and hydrogen has 1 valence electron. The total number of valence electrons for the molecule is: \(6 + 2(1) + 4(6) = 32\). Step 2: Arrange atoms Place the sulfur atom in the center, surrounded by the four oxygen atoms. Attach the two hydrogen atoms to two of the oxygen atoms. Step 3: Distribute electrons Form single bonds between the sulfur and each oxygen atom, single bonds between the hydrogen and oxygen atoms, and distribute remaining electrons pairs according to the octet rule. The Lewis structure for \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is: ``` O || O-S-O || O-H ```

e) \(\mathrm{ClO}_{2}^{-}\)

Step 1: Count total valence electrons Chlorine has 7 valence electrons, each oxygen atom has 6 valence electrons, and there is one extra electron due to the negative charge. The total number of valence electrons for the ion is: \(7 + 2(6) + 1 = 20\). Step 2: Arrange atoms Place the chlorine atom in the center and arrange the two oxygen atoms around it. Step 3: Distribute electrons and account for formal charge Form single bonds between the chlorine and each oxygen atom, and distribute remaining electrons pairs according to the octet rule. Ensure the formal charge is correct by adjusting bonds if necessary. The Lewis structure for \(\mathrm{ClO}_{2}^{-}\) is: ``` O \ Cl - O / O ```

f) \(\mathrm{NH}_{2} \mathrm{OH}\)

Step 1: Count total valence electrons Nitrogen has 5 valence electrons, each hydrogen atom has 1 valence electron, and oxygen has 6 valence electrons. The total number of valence electrons for the molecule is: \(5 + 2(1) + 6 + 1(1) = 14\). Step 2: Arrange atoms Attach the two hydrogen atoms to nitrogen, and attach the oxygen to the nitrogen atom. Attach the final hydrogen atom to oxygen. Step 3: Distribute electrons Form single bonds between nitrogen and each hydrogen atom, single bonds between nitrogen and oxygen, and a single bond between oxygen and hydrogen. Distribute remaining electrons pairs according to the octet rule. The Lewis structure for \(\mathrm{NH}_{2}\mathrm{OH}\) is: ``` O | N - H | H ```

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom. They play a crucial role in determining how elements interact chemically. Understanding valence electrons is fundamental when drawing Lewis structures, as it guides us in determining how atoms bond with each other. For example:

• Silicon, being in group 14 of the periodic table, has 4 valence electrons.
• Hydrogen, found in group 1, has 1 valence electron.
• Oxygen, which belongs to group 16, has 6 valence electrons.
• Chlorine, in group 17, has 7 valence electrons.

To calculate the total number of valence electrons for a molecule or ion, sum up the valence electrons of each atom involved. For ions, account for any additional electrons caused by the ion's charge. This step is critical, as it ensures enough electrons are available to form bonds while respecting each atom's need for a stable electronic arrangement.

Covalent Bonds

Covalent bonds form when atoms share valence electrons. This type of bond allows atoms to achieve a more stable electron configuration. Covalent bonding is prevalent in many nonmetal compounds and can result in single, double, or even triple bonds.

• Single bond: occurs when two atoms share one pair of electrons. This is seen in molecules like \( \mathrm{SiH}_4 \), where silicon forms single bonds with four hydrogen atoms.
• Double bond: happens when two atoms share two pairs of electrons. An example is the molecule \( \mathrm{CO} \), where carbon and oxygen form a double bond to satisfy both atoms' needs for a stable electronic configuration.
• Triple bond: involves sharing three pairs of electrons, not depicted in our examples but commonly seen in molecules like nitrogen gas (\( \mathrm{N}_2 \)).

Covalent bonds result in a balance of attractive and repulsive forces between atoms, creating stable structures. Forming the correct number of covalent bonds when drawing Lewis structures is essential to represent the molecule accurately.

Octet Rule

The octet rule is a guiding principle in chemistry stating that atoms tend to prefer having eight electrons in their valence shell.This rule helps predict how atoms will bond and arrange themselves within a molecule. While not all elements strictly adhere to the octet rule, it provides a useful framework for understanding molecular structures:

• In \( \mathrm{SiH}_4 \), silicon aims to share its electrons to achieve a complete octet, pairing with four hydrogen atoms.
• \( \mathrm{CO} \) and \( \mathrm{ClO}_2^- \) molecules involve forming additional bonds, such as double bonds, to satisfy the octet needs of oxygen and chlorine.
• Exceptions exist: Hydrogen is content with 2 electrons, reaching a duet rather than an octet, as seen in molecules like \( \mathrm{H}_2\ \text{and} \ \mathrm{NH}_2\mathrm{OH} \).

Variance from the octet rule occurs in some cases, where elements accommodate more or fewer than eight electrons. However, for most main-group elements, achieving an octet contributes to a stable and lower energy state.