Question

Write Lewis structures that obey the octet rule (duet rule for H) for each of the following molecules. Carbon is the central atom in \(\mathrm{CH}_{4}\), nitrogen is the central atom in \(\mathrm{NH}_{3}\), and oxygen is the central atom in \(\mathrm{H}_{2} \mathrm{O}\). a. \(\mathrm{F}_{2}\) \(\mathbf{b} . \mathbf{O}_{2}\) c. CO d. \(\overline{\mathrm{CH}_{4}}\) \(\mathbf{e} . \mathrm{NH}_{3}\) \(\mathbf{f .} \quad \mathbf{H}_{2} \mathbf{O}\) g. IIF

Short answer

Lewis structures: a. F2: F - F b. O2: O = O c. CO: C ≡ O d. CH4: H | C - H | H e. NH3: H | N - H - H f. H2O: H - O - H g. IIF: I - F

Step-by-step solution

a. \(\mathrm{F}_{2}\)

1. Count the valence electrons: Fluorine has 7 valence electrons, so the total for F2 is 14. 2. Create a single bond between the fluorine atoms, accounting for 2 electrons. 3. Distribute the remaining 12 electrons in pairs to the outside atoms (fluorine) to satisfy the octet rule for each atom. Lewis structure for F2: F - F

b. \(\mathrm{O}_{2}\)

1. Count the valence electrons: Oxygen has 6 valence electrons, so the total for O2 is 12. 2. Create a double bond between the oxygen atoms, accounting for 4 electrons. 3. Distribute the remaining 8 electrons in pairs to the outside atoms (oxygen) to satisfy the octet rule for each atom. Lewis structure for O2: O = O

c. CO

1. Count the valence electrons: Carbon has 4 and oxygen has 6, so the total is 10. 2. Create a triple bond between the carbon and oxygen atoms, accounting for 6 electrons. 3. Distribute the remaining 4 electrons in pairs to the outside atom (oxygen) to satisfy the octet rule for each atom. Lewis structure for CO: C ≡ O

d. \(\mathrm{CH}_{4}\)

1. Count the valence electrons: Carbon has 4 and each hydrogen has 1, so the total is 8. 2. Create a single bond between the central atom (carbon) and the hydrogen atoms, accounting for 4 electrons. 3. Distribute the remaining 4 electrons in pairs to the hydrogen atoms to satisfy the duet rule for each hydrogen atom. Lewis structure for CH4: H | C - H | H

e. \(\mathrm{NH}_{3}\)

1. Count the valence electrons: Nitrogen has 5 and each hydrogen has 1, so the total is 8. 2. Create a single bond between the central atom (nitrogen) and the hydrogen atoms, accounting for 3 electrons. 3. Distribute the remaining 5 electrons: 3 as a lone pair on the nitrogen atom, and 2 as a pair to one hydrogen atom to satisfy the duet rule for each hydrogen atom. Lewis structure for NH3: H | N - H - H

f. \(\mathrm{H}_{2}\mathrm{O}\)

1. Count the valence electrons: Oxygen has 6 and each hydrogen has 1, so the total is 8. 2. Create a single bond between the central atom (oxygen) and the hydrogen atoms, accounting for 2 electrons. 3. Distribute the remaining 6 electrons in pairs to the outside hydrogen atoms to satisfy the duet rule for each hydrogen atom. Lewis structure for H2O: H - O - H

g. IIF

1. Count the valence electrons: Iodine has 7 and fluorine has 7, so the total is 14. 2. Create a single bond between the central atom (iodine) and the fluorine atoms, accounting for 2 electrons. 3. Distribute the remaining 12 electrons in pairs to the outside atoms (iodine and fluorine) to satisfy the octet rule for each atom. Lewis structure for IIF: I - F

Key concepts

Octet Rule

The octet rule is a fundamental principle in chemistry that helps us understand how atoms bond to form molecules. The rule states that atoms tend to combine in such a way that they each have eight electrons in their outer shell. This configuration is associated with stability, similar to the noble gases.

- Atoms will share, lose, or gain electrons to achieve a full octet.- Exceptions include hydrogen, which follows the "duet rule" and seeks only two electrons in total.- The octet rule guides us in drawing Lewis structures, illustrating how atoms share electrons.
For example, in the molecule \(\mathrm{F}_2\), each fluorine atom shares one electron to satisfy the octet rule, resulting in a single bond.

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom and play a crucial role in chemical bonding. These electrons are the ones involved in forming bonds between atoms.

- Valence electrons determine how atoms interact and bond.- The number of valence electrons can be figured out by the element's group number in the periodic table.- Knowing the number of valence electrons helps in predicting the types of bonds an atom can form.
In our examples, like \(\mathrm{O}_2\), each oxygen atom has six valence electrons. By understanding this, we can predict that each oxygen needs two more electrons for a full octet, leading to the formation of a double bond.

Chemical Bonding

Chemical bonding is the process that atoms undergo to attain a stable electronic configuration. There are several types of bonds that can form, depending on how the valence electrons are shared or transferred.

- **Ionic Bonds:** Formed when electrons are transferred from one atom to another, leading to the formation of ions.- **Covalent Bonds:** Formed when atoms share pairs of electrons.
In the context of covalent bonding, molecules like \(\mathrm{CH}_4\) feature single bonds where carbon and hydrogen atoms share electrons.

Understanding the nature of these bonds helps us predict the structure and properties of molecules.

Diatomic Molecules

Diatomic molecules consist of two atoms bonded together. This is one of the simplest forms of molecules, often leading to straightforward examples of chemical bonding.

- Some common diatomic molecules include \(\mathrm{H}_2, \mathrm{N}_2, \mathrm{O}_2\), and \(\mathrm{F}_2\).- They can be composed of the same element, like \(\mathrm{F}_2\), or different elements, like hydrogen fluoride (HF).
Diatomic molecules are great examples to understand the basics of chemical bonds and the octet rule, as they often form simple single or double bonds, making it easier to visualize electron sharing. For instance, in \(\mathrm{O}_2\), each oxygen atom shares two electrons to form a double bond, achieving a full octet for both atoms.