Question

Write a Lewis formula for each of the following: (a) \(\mathrm{NH}_{4}{ }^{+}\) (c) \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO}\) (e) \(\mathrm{CH}_{3} \mathrm{NO}_{2}\) (b) \(\mathrm{H}_{2} \mathrm{CO}\) (d) \(\mathrm{OH}^{-}\)

Short answer

The lewis structures for the given molecules are represented as following: (a) \(\mathrm{NH}_{4}{ }^{+}\) - N is bonded with H's; (c) \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO}\) - 3 lone pairs on O and 2 C-H bonds and 2 CH3 groups; (e) \(\mathrm{CH}_{3} \mathrm{NO}_{2}\) - C single bonds with H's and N, N has a double bond with O and a single bond with another O. That O has 3 lone pairs. (b) \(\mathrm{H}_{2} \mathrm{CO}\) - C forms a double bond with O(has two lone pairs) and single bonds with Hydrogen; (d) \(\mathrm{OH}^{-}\) - O formas a single bond with H and has 3 lone pairs.

Step-by-step solution

Lewis Structure of \(\mathrm{NH}_{4}{ }^{+}\)

N has 5 valence electrons, and each H has 1 electron(total 4 electrons for 4 H), but since it is cationic with a charge of +1, one electron is lost. Thus, total available electrons are 5+4-1=8. N is central atom, and all H's are connecting it with single bonds using 8 electrons. Hence, all electrons are used up, leading to the Lewis structure N with four H's attached to it with a single bond.

Lewis Structure of \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO}\)

The compound has a total of 24 valence electrons. 3 lone pairs on Oxygen for a total of 6 electrons, double bond O-C for a total 4 electrons, two C-H single bonds with each containing 2 electrons for a total of 4, and then two CH3 groups each using 8 electrons for a total of 16. All 24 electrons are accounted for.

Lewis Structure of \(\mathrm{CH}_{3} \mathrm{NO}_{2}\)

Here, there are 24 valence electrons. A carbon atom makes three single bonds with hydrogen atoms using 6 electrons and with the N atom using two electrons. N makes a double bond O (using four electrons) and a single bond to another O (using 2 electrons), and that O has 3 lone pair(electron pairs) (using 6 electrons). Thus, all 24 valence electrons are used.

Lewis Structure of \(\mathrm{H}_{2} \mathrm{CO}\)

There are 12 valence electrons in this compound. C is the central atom and forms two single bonds with two H atoms using 4 electrons. Carbon also forms a double bond with O, using four electrons, and O has 2 lone pair, using 4 electrons.

Lewis Structure of \(\mathrm{OH}^{-}\)

O has 6 valence electrons, and H has 1 electron. An extra electron is added because of the -1 charge. Therefore, we have 8 valence electrons in total. O forms a single bond with hydrogen using 2 electrons and has 3 lone pairs, using 6 electrons. With this, all electrons are used.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in forming chemical bonds. They are the electrons involved in interactions between atoms to form molecules. The number of valence electrons determines an atom's bonding capacity. For example:

• Nitrogen (\( ext{N}\)) has 5 valence electrons.
• Oxygen (\( ext{O}\)) possesses 6 valence electrons.
• Hydrogen (\( ext{H}\)) has just 1 valence electron.

Understanding the number of valence electrons an atom has helps predict how it will bond with other atoms. In the case of\( ext{NH}_4^+\), the nitrogen atom uses its 5 valence electrons to bond with four hydrogen atoms, but due to the +1 charge, you subtract one electron, ending up with an effective count of 8 electrons forming bonds. This intricate dance involving valence electrons is what allows diverse molecules with different properties to form.

Chemical Bonds

Chemical bonds are the connections between atoms in a molecule, formed by the sharing or transferring of electrons to achieve a full outer shell of electrons. There are multiple types of chemical bonds, but in the context of Lewis structures, we focus primarily on covalent bonds:

• **Single Bonds**: Involve sharing one pair of electrons (2 electrons), such as in the connection between the carbon and hydrogen atoms in\( ext{H}_2 ext{CO}\).
• **Double Bonds**: Involve two pairs of shared electrons (4 electrons), like the C=O bond in the same compound.

In the Turner molecule \( ext{CH}_3 ext{NO}_2\), a variety of bonds exist to accommodate the number of valence electrons each atom has and to fulfill the octet rule (where possible). Chemical bonds are vital for the stability and functionality of molecules.

Lone Pairs

Lone pairs, or non-bonding pairs, are pairs of valence electrons that are not shared with another atom and thus do not participate in bonding. These electrons are represented as dots in Lewis structures and play a significant role in the shape and reactivity of molecules.

• In\( ext{OH}^-\), the oxygen atom has three lone pairs of electrons. These help to complete its octet, along with the two electrons shared with hydrogen in the O-H bond.
• The lone pairs also affect molecular geometry by repelling bonding pairs, causing molecules to adopt shapes that minimize this repulsion.

In \( ext{CH}_3 ext{NO}_2\), for example, lone pairs on the oxygen atom significantly influence the three-dimensional geometry and polarity of the molecule. Lone pairs are essential for predicting the behavior of molecules during reactions, as they can partake in phenomena such as coordinate bonding.