Question

Draw Lewis structures for the following molecules: (a) $\mathrm{ICl},$ (b) ${\mathrm{PH}}_3,$ (c) ${\mathrm{P}}_4$ (each $\mathrm{P}$ is bonded to three other $\mathrm{P}$ atoms), (d) ${\mathrm{H}}_2\mathrm{S},$ (e) ${\mathrm{N}}_2{\mathrm{H}}_4$ (f) ${\mathrm{HClO}}_3$

Short answer

Draw each Lewis structure according to valence electrons and bonding rules: ICl (single bond, lone pairs), PH₃ (lone pair on P), P₄ (P bonded to 3 other P's), H₂S (lone pairs on S), N₂H₄ (single N-N bond, lone pairs on N), HClO₃ (single bonds, possible double bond for an octet).

Step-by-step solution

Understanding Electrons and Bonds

Before we start drawing the Lewis structures, it's important to understand that each atom wants to complete its outer shell, typically aiming for 8 electrons (octet rule), though hydrogen seeks 2. We'll count the total number of valence electrons available for each molecule, as this will guide the drawing. Also, remember that shared pairs of electrons indicate covalent bonds.

Drawing \( \mathrm{ICl} \\)

Iodine and chlorine both have 7 valence electrons. That's a total of 14 valence electrons. Here's how we distribute them: - Place I and Cl next to each other, connected by a single bond (using 2 electrons). - Distribute the remaining 12 electrons to complete the outer shells (octet rule for both I and Cl). Both will have one single bond and three lone pairs.

Drawing \( \mathrm{PH}_3 \\)

Phosphorus has 5 valence electrons, and each hydrogen has 1, totaling up to 8 valence electrons. - Place P in the center. - Assign each of the 3 H atoms around P, forming three P-H bonds (using 6 electrons). - The remaining 2 electrons form a lone pair on the P atom.

Drawing \( \mathrm{P}_4 \\)

Each phosphorous atom has 5 valence electrons, and in ${\mathrm{P}}_4$, each P is bonded to 3 other P atoms. - Arrange P atoms in a tetrahedral shape where each P is connected to the three others in a circle or square.- Form single bonds between each pair of adjacent P atoms. Use the remaining electrons to satisfy octet rules for P atoms, forming lone pairs as needed.

Drawing \( \mathrm{H}_2 \mathrm{S} \\)

Sulfur has 6 valence electrons, and each hydrogen has 1, making a total of 8. - Place S in the center. - Form two single bonds to the H atoms (using 4 electrons). - Place the remaining 4 electrons as two lone pairs on the sulfur atom to complete its octet.

Drawing \( \mathrm{N}_2 \mathrm{H}_4 \\)

Each nitrogen has 5 valence electrons, and each hydrogen has 1. For ${\mathrm{N}}_2{\mathrm{H}}_4$, that's a total of 14.- Connect the 2 N atoms with a single bond.- Attach 2 H atoms to each N atom (forming 4 N-H bonds, using 8 electrons).- Use the remaining 4 electrons to form lone pairs on each nitrogen atom.

Drawing \( \mathrm{HClO}_3 \\)

Chlorine typically forms 5 bonds, oxygen wants 2 electrons more than it has, and hydrogen 1. - Place Cl in the center, connect it with single bonds to 3 O atoms. Use double bonds if needed to satisfy the octet for O. - Attach 1 H to an O, and adjust electrons on the O's to complete the octet for each atom. Assign lone pairs to O's and if Cl has fewer than 5 bonds, use them to form a coordinate bond as needed.

Key concepts

Octet Rule

In chemistry, the octet rule is a principle that reflects the tendency of atoms to have eight electrons in their outer shell to achieve stability. This is akin to the electron configuration of noble gases, which are inherently stable. The octet rule is a guiding principle for drawing Lewis structures, where we depict molecules using dots to represent valence electrons.
Atoms will typically share, lose, or gain electrons to fulfill the octet rule:

• Sharing electrons leads to covalent bonds.
• Losing or gaining electrons results in ionic bonds, but that's for another lesson!

While observing the octet rule, it's important to note exceptions. Hydrogen is a notable one; it seeks only two electrons. Other notable exceptions include molecules with odd numbers of electrons or atoms that prefer to have fewer or more than eight electrons, such as phosphorus or chlorine.

Covalent Bonds

Covalent bonds form when two atoms share pairs of electrons. This happens primarily between non-metal atoms, such as in the cases of iodine and chlorine in a molecule like $\mathrm{ICl}$. The shared electrons crisscross the space around both atoms, holding them together energetically.
Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs:

• Single Bonds: Sharing one pair of electrons, like in hydrogen chloride $\mathrm{HCl}$.
• Double Bonds: Sharing two pairs of electrons, for example in oxygen ${\mathrm{O}}_2$.
• Triple Bonds: Sharing three pairs of electrons, like in nitrogen ${\mathrm{N}}_2$.

In drawing Lewis structures, we show these covalent bonds as lines connecting atoms. One line represents a shared pair of electrons. For example, in ${\mathrm{PH}}_3$, each P-H bond is formed by sharing a pair of electrons between phosphorus and hydrogen.Remember, covalent bonding is crucial for achieving the octet rule for many atoms

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a pivotal role in chemical bonding because they can be transferred or shared between atoms to form bonds. When drawing Lewis structures, these are the electrons we are most interested in.
Each element has a characteristic number of valence electrons:

• Chlorine has seven, needing one more to complete an octet, which is achieved by forming a covalent bond.
• Phosphorus has five, allowing it to form three covalent bonds for a stable structure like in ${\mathrm{P}}_4$.
• Hydrogen, seeking to reach two electrons, will typically form one covalent bond.

Understanding valence electrons is essential because this concept:

• Guides the formation of molecules and how atoms will bond.
• Helps in predicting the shape and reactivity of molecules. For example, ${\mathrm{H}}_2\mathrm{S}$ has two lone pairs on sulfur that affect its molecular shape and properties.

Counting the total number of valence electrons helps us to distribute them appropriately to satisfy the octet rule (or exceptions) while drawing Lewis structures.

Molecular Geometry

Molecular geometry refers to the spatial arrangement of atoms within a molecule. Understanding molecular geometry is crucial as it affects the physical and chemical properties of compounds. The arrangement is dictated by the number of bonds and lone pairs on the central atom.
Some common geometries include:

• Linear: Central atom with two bonded atoms and no lone pairs, like ${\mathrm{CO}}_2$.
• Tetrahedral: Central atom bonded to four other atoms, as with ${\mathrm{CH}}_4$ or ${\mathrm{P}}_4$ structures.
• Trigonal Pyramidal: Central atom with three bonds and one lone pair, seen in ${\mathrm{PH}}_3$.
• Bent: Central atom with two bonds and two lone pairs, demonstrated in ${\mathrm{H}}_2\mathrm{S}$.

The number of lone pairs and bonds influences the molecular geometry as the electrons will repel each other. This results in specific angles and shapes that enhance or impede interactions with other molecules. Understanding these shapes helps predict molecular reactions and interactions in various environments.