Question

Write Lewis structures that obey the octet rule for each of the following molecules. a. \(\mathrm{CCl}_{4}\) c. \(\mathrm{SeCl}_{2}\) b. \(\mathrm{NCl}_{3}\) d. \(\mathrm{ICl}\) In each case, the atom listed first is the central atom.

Short answer

a. In \(\mathrm{CCl}_{4}\), Carbon (C) forms 4 single bonds with 4 Chlorine (Cl) atoms, each Cl has 3 lone pairs, and the Lewis structure appears as: ``` Cl | Cl -- C -- Cl | Cl ``` c. In \(\mathrm{SeCl}_{2}\), Selenium (Se) forms 2 single bonds with 2 Chlorine (Cl) atoms and has a lone pair, each Cl has 3 lone pairs, and the Lewis structure is: ``` Cl | Se -- Cl ¦ ``` b. In \(\mathrm{NCl}_{3}\), Nitrogen (N) forms 3 single bonds with 3 Chlorine (Cl) atoms and has a lone pair, each Cl has 3 lone pairs, and the Lewis structure appears as: ``` Cl | Cl -- N -- Cl ¦ ``` d. In \(\mathrm{ICl}\), Iodine (I) forms a single bond with the Chlorine (Cl) atom, and both I and Cl have 3 lone pairs each, and the Lewis structure appears as: ``` I -- Cl ¦ ¦ ```

Step-by-step solution

Determine the total number of valence electrons for each molecule

To do this, we need to add the valence electrons of each atom in the molecule. Remember that in the periodic table, group 1 has 1 valence electron, group 2 has 2 valence electrons, and so on. a. \(\mathrm{CCl}_{4}\): Carbon (C) has 4 valence electrons and Chlorine (Cl) has 7 valence electrons each. So, the total valence electrons are \(4 + 4(7) = 32\). c. \(\mathrm{SeCl}_{2}\): Selenium (Se) has 6 valence electrons and Chlorine (Cl) has 7 valence electrons each. So, the total valence electrons are \(6 + 2(7) = 20\). b. \(\mathrm{NCl}_{3}\): Nitrogen (N) has 5 valence electrons and Chlorine (Cl) has 7 valence electrons each. So, the total valence electrons are \(5 + 3(7) = 26\). d. \(\mathrm{ICl}\): Iodine (I) has 7 valence electrons and Chlorine (Cl) has 7 valence electrons. So, the total valence electrons are \(7 + 7 = 14\).

Distribute the valence electrons to satisfy the octet rule

We need to distribute the electrons in such a way that each atom is surrounded by eight electrons, either in lone pairs or bond pairs. a. \(\mathrm{CCl}_{4}\): Carbon (C) shares 1 electron with each of the 4 Chlorine (Cl) atoms, forming 4 single bonds. All 4 Chlorine atoms have 3 lone pairs, and Carbon has no lone pairs. c. \(\mathrm{SeCl}_{2}\): Selenium (Se) shares 1 electron with each of the 2 Chlorine (Cl) atoms, forming 2 single bonds. Both Chlorine atoms have 3 lone pairs. However, to complete the octet for Selenium, which has 6 valence electrons, it also needs to have one lone pair. b. \(\mathrm{NCl}_{3}\): Nitrogen (N) shares 1 electron with each of the 3 Chlorine (Cl) atoms, forming 3 single bonds. All 3 Chlorine atoms have 3 lone pairs, and Nitrogen has 1 lone pair. d. \(\mathrm{ICl}\): Iodine (I) shares 1 electron with the Chlorine (Cl) atom, forming a single bond. Both Iodine and Chlorine have 3 lone pairs each.

Draw the Lewis structures for each molecule

Now, we will draw the Lewis structures for each molecule, keeping in mind the number of bond pairs and lone pairs discussed in the previous step. a. \(\mathrm{CCl}_{4}\): ``` Cl | Cl -- C -- Cl | Cl ``` c. \(\mathrm{SeCl}_{2}\): ``` Cl | Se -- Cl ¦ ``` b. \(\mathrm{NCl}_{3}\): ``` Cl | Cl -- N -- Cl ¦ ``` d. \(\mathrm{ICl}\): ``` I -- Cl ¦ ¦ ```

Key concepts

Octet Rule

The octet rule is a fundamental principle in chemistry that helps us understand how atoms bond with each other. Essentially, this rule states that atoms prefer to have eight electrons in their outermost shell, as this electron configuration is associated with stability, much like the noble gases. There are a few exceptions, like hydrogen and helium, which are stable with two electrons due to their smaller size.

For example, in the case of carbon tetrachloride (\( \text{CCl}_4 \)), each chlorine atom gains additional electrons by sharing with the central carbon atom until all atoms have a complete octet. Thus, the octet rule guides us in distributing electrons and forming bonds in a molecule to achieve maximum stability.

To apply the octet rule when drawing Lewis structures:

• First, count the total valence electrons in the compound.
• Distribute the electrons around each atom until each has an octet (or duet for hydrogen).
• Use single or multiple bonds to connect atoms as needed while adhering to the total electron count.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding. These electrons are responsible for the interactions between different atoms when forming molecules. Typically, the number of valence electrons determines how an atom will bond with others.

To identify valence electrons:

• Check the Group number in the periodic table for the element. This will give you the number of valence electrons for most elements (e.g., Group 1 elements have 1 valence electron, Group 17 have 7).
• Sum up these electrons for the entire molecule to know how many electrons must be placed in the Lewis structure.

In molecules like \( \text{NCl}_3\), nitrogen has 5 valence electrons and chlorine has 7. Calculating the total available valence electrons helps determine the structure by ensuring electrons are properly distributed.

Understanding valence electrons is crucial since they impact the shape, reactivity, and properties of the resulting compounds.

Chemical Bonding

Chemical bonding is the process where atoms come together to form a molecule by sharing or transferring electrons. This process results in a stable energy configuration that benefits all the atoms involved.

The main types of bonds are:

• Covalent bonds, where electrons are shared between atoms, like in \( \text{SeCl}_2\). Here, selenium shares electrons with chlorine to satisfy the octet rule.
• Ionic bonds, where electrons are transferred, resulting in charged ions that attract each other.
• Metallic bonds, which occur between metal atoms, sharing free electrons in a "sea."

Covalent bonding often involves sharing pairs of electrons – a single covalent bond involves one pair of shared electrons, whereas double or triple bonds involve two or three pairs of electrons, respectively.

In understanding \( \text{ICl}\), the formation of a single bond between iodine and chlorine involves sharing one electron pair, illustrating how chemical bonds are integral to molecule formation.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. This geometry significantly impacts physical properties and chemical reactivity.

The VSEPR (Valence Shell Electron Pair Repulsion) theory is commonly used to predict molecular shapes based on the idea that electron pairs repel each other and arrange themselves to minimize this repulsion.

Important points about molecular geometry:

• The number of bonding electron pairs and lone electron pairs around an atom determine its geometry.
• Molecular shapes such as linear, bent, tetrahedral, pyramidal, and trigonal planar arise from differing electron pair arrangements.

For instance, \( \text{CCl}_4\) has a tetrahedral shape as the central carbon forms four bonds, arranging themselves symmetrically. Whereas \( \text{SeCl}_2\) is bent due to the presence of lone pairs on selenium affecting the arrangement of atoms.

Understanding molecular geometry is essential for predicting how molecules will interact and react with one another.