Question

Draw Lewis structures for the following: \((\mathbf{a}) \mathrm{CH}_{2} \mathrm{Cl}_{2},(\mathbf{b}) \mathrm{ClCN},\) (c) \(\mathrm{AsF}_{5}\) (d) \(\mathrm{CH}_{2} \mathrm{O}\) (C the central atom), (e) \(\mathrm{OF}_{2}\) (f) \(\mathrm{NO}_{3}^{-}\).

Short answer

The Lewis structures for the given compounds are as follows: - a) CH2Cl2: C bonded to 2 hydrogen atoms and 2 chlorine atoms arranged tetrahedrally, forming 4 single bonds. - b) ClCN: Cl connected to C with a single bond, C connected to N with a triple bond. - c) AsF5: As bonded to 5 fluorine atoms arranged in a trigonal bipyramidal configuration, forming 5 single bonds. - d) CH2O: C bonded to 2 hydrogen atoms and 1 oxygen atom arranged trigonally planar, forming 2 single bonds and 1 double bond with oxygen. - e) OF2: O bonded to 2 fluorine atoms in a linear configuration, forming 2 single bonds. - f) NO3-: N bonded to 3 oxygen atoms in a trigonal planar configuration, forming 1 single bond with 2 oxygen atoms and 1 double bond with the remaining oxygen atom.

Step-by-step solution

Identify the central atom

For each of the given compounds, the central atom should be the least electronegative element, except in the case of hydrogen (since hydrogen can only form one bond). The central atoms for each compound are as follows: - a) CH2Cl2: Carbon (C) - b) ClCN: Carbon (C) - c) AsF5: Arsenic (As) - d) CH2O: Carbon (C) - e) OF2: Oxygen (O) - f) NO3-: Nitrogen (N)

Find the total number of valence electrons

Now, we will find the total number of valence electrons for each compound. This can be done by adding up the number of valence electrons for each element and, in the case of ions, adding or subtracting the respective number of electrons. - a) CH2Cl2: 4 (C) + 2(1) (H) + 2(7) (Cl) = 20 valence electrons - b) ClCN: 7 (Cl) + 4(C) + 5(N) = 16 valence electrons - c) AsF5: 5 (As) + 5(7) (F) = 40 valence electrons - d) CH2O: 4 (C) + 2(1) (H) + 6 (O) = 12 valence electrons - e) OF2: 6 (O) + 2(7) (F) = 20 valence electrons - f) NO3-: 5 (N) + 3(6) (O) + 1 (extra electron from the negative charge) = 24 valence electrons

Arrange the electrons around the atoms

Now, for each compound, we will arrange the electrons around the atoms in a way that satisfies the octet rule, keeping in mind that hydrogen only needs two electrons to have a full shell. - a) CH2Cl2: 1. Start with Carbon in the center. 2. Place two Hydrogen atoms and two Chlorine atoms around the Carbon in a tetrahedral configuration. 3. Connect each surrounding atom to the Carbon with single bonds. This uses 8 of the 20 valence electrons. 4. Fill the octets of the surrounding atoms with the remaining 12 valence electrons (6 pairs). - b) ClCN: 1. Start with Carbon in the center. 2. Place Chlorine and Nitrogen atoms on either side of the Carbon. 3. Connect the Chlorine and the Nitrogen to the Carbon with single bonds. This uses 4 of the 16 valence electrons. 4. Fill the octet of Chlorine with 6 more valence electrons (3 pairs). 5. Place the remaining 6 valence electrons (3 pairs) on the Nitrogen atom to form a triple bond between the Nitrogen and Carbon. - c) AsF5: 1. Place Arsenic in the center. 2. Surround it with five Fluorine atoms. 3. Connect each Fluorine to the Arsenic with single bonds. This uses 10 of the 40 valence electrons. 4. Fill the octets of the surrounding Fluorine atoms with the remaining 30 valence electrons (15 pairs). - d) CH2O: 1. Start with Carbon in the center. 2. Place two Hydrogen atoms and one Oxygen atom around the Carbon in a trigonal planar configuration. 3. Connect each surrounding atom to the Carbon with single bonds. This uses 6 of the 12 valence electrons. 4. Place the remaining 6 valence electrons (3 pairs) around the Oxygen atom, forming a double bond between Carbon and Oxygen. - e) OF2: 1. Place Oxygen in the center. 2. Place two Fluorine atoms around the Oxygen in a linear configuration. 3. Connect each Fluorine to the Oxygen with single bonds. This uses 4 of the 20 valence electrons. 4. Fill the octets of the surrounding atoms with the remaining 16 valence electrons (8 pairs). - f) NO3-: 1. Start with Nitrogen in the center. 2. Place three Oxygen atoms around the Nitrogen in a trigonal planar configuration. 3. Connect each Oxygen to the Nitrogen with single bonds. This uses 6 of the 24 valence electrons. 4. Fill the octets of two surrounding Oxygen atoms with 12 more valence electrons (6 pairs). 5. Place the remaining 6 valence electrons (3 pairs) around one Oxygen atom, forming a double bond between that Oxygen and Nitrogen. In conclusion, we have drawn the Lewis structures for all six compounds by identifying the central atom, finding the total number of valence electrons, and arranging the electrons around the atoms in a way that satisfies the octet rule.

Key concepts

Valence Electrons

Valence electrons are crucial for understanding chemical bonding and reactivity. These are the electrons present in the outermost electron shell of an atom. They are important because they are involved in forming bonds with other atoms.
For example, carbon has four valence electrons, hydrogen has one, chlorine has seven, nitrogen has five, and oxygen has six. Knowing the number of valence electrons helps in predicting how atoms combine to form compounds.
To calculate the total number of valence electrons in a molecule, simply add up the valence electrons from all the atoms involved. In the case of ions, like \( ext{NO}_3^-\), you must also account for the charge by adding or removing electrons. For example, \( ext{NO}_3^-\) has 24 valence electrons because we add an extra electron to account for the negative charge.

Octet Rule

The octet rule is a guideline used in chemistry to predict the bonding behavior of atoms. It suggests that atoms are more stable when they have eight valence electrons, resembling the electron configuration of a noble gas. However, there are exceptions such as hydrogen, which only needs two electrons to achieve stability.
When drawing Lewis structures, the goal is to arrange the electrons so that as many atoms as possible have an octet. This may involve forming single, double, or even triple bonds. Some elements, like phosphorus and sulfur, can have expanded octets and hold more than eight electrons.
In compounds like \( ext{AsF}_5\), arsenic has an expanded octet with ten valence electrons when it forms five single bonds with fluorine atoms. This shows how some elements can exceed the simple octet rule.

Covalent Bonding

Covalent bonding occurs when atoms share pairs of valence electrons to achieve stability. This type of bond typically forms between nonmetals, which have similar electronegativities.
In Lewis structures, covalent bonds are depicted as lines connecting atoms. Each line represents a pair of shared electrons. For example, in \( ext{CH}_2 ext{O} \), carbon forms single bonds with hydrogen by sharing one pair of electrons with each hydrogen atom. To satisfy the octet rule for oxygen, a double bond is formed between carbon and oxygen, sharing two pairs of electrons.
Understanding covalent bonding is essential as it explains the structures and strengths of molecular compounds. These bonds can vary in strength from weak (van der Waals forces) to very strong (covalent network bonds). A solid understanding of electron sharing and bond strength helps in predicting molecular behavior and reactivity.

Chemical Compounds

Chemical compounds are substances formed by the chemical combination of two or more different elements. These can be categorized as ionic or covalent, based on the type of bonding between the atoms.
For example, in a covalent compound like \( ext{CH}_2 ext{Cl}_2 \), atoms share electrons to form stable configurations. The specific arrangement of these atoms is depicted through Lewis structures, which visually represent the molecular structure and bonding.
The properties of chemical compounds depend not only on the types of atoms present but also on their arrangement and the type of bonds formed. For instance, \( ext{AsF}_5 \) is a chemical compound where a central arsenic atom forms single bonds with five fluorine atoms, resulting in a specific molecular geometry that affects its reactivity and interaction with other substances.
By analyzing Lewis structures, we can predict molecular geometries, possible chemical reactions, and the physical properties of compounds, aiding in the understanding and application of chemical principles in real-world scenarios.