Question

Draw Lewis structures for the following ions: (a) \(\mathrm{NO}_{2}^{+}\), (b) \(\mathrm{SCN}^{-},\) (c) \(\mathrm{S}_{2}^{2-}\) (d) \(\mathrm{ClF}_{2}^{+}\). Show formal charges.

Short answer

NO2+ has a positive charge on N, SCN- has a negative charge on N, S2^2- has negative charges on both S atoms, ClF2+ has a positive charge on Cl.

Step-by-step solution

Determine Total Number of Valence Electrons

For each ion, calculate the total number of valence electrons, adding for negative charges or subtracting for positive charges. - (a) \(\text{NO}_2^+\): Nitrogen has 5, each oxygen has 6. Remove 1 electron for the positive charge. Total = 5 + 2(6) - 1 = 16 electrons.- (b) \(\text{SCN}^-\): Sulfur has 6, carbon has 4, nitrogen has 5, add 1 electron for the negative charge. Total = 6 + 4 + 5 + 1 = 16 electrons.- (c) \(\text{S}_2^{2-}\): Each sulfur has 6, add 2 electrons for 2- charge. Total = 2(6) + 2 = 14 electrons.- (d) \(\text{ClF}_2^+\): Chlorine has 7, each fluorine has 7, remove 1 electron for the positive charge. Total = 7 + 2(7) - 1 = 20 electrons.

Draw Skeleton Structures

Decide on connectivity for each ion, usually placing the least electronegative atom in the center.- (a) \(\text{NO}_2^+\) is \(\text{O-N-O}\), with N in the center.- (b) \(\text{SCN}^-\) is \(\text{S-C-N}\), with C in the center.- (c) \(\text{S}_2^{2-}\) is \(\text{S-S}\).- (d) \(\text{ClF}_2^+\) is \(\text{F-Cl-F}\), with Cl in the center.

Distribute Electrons

Start by placing electrons to fulfill the octet rule, starting with terminal atoms.- (a) \(\text{NO}_2^+\): Distribute 8 electrons to each O (4 shared with N), leaving N with none initially.- (b) \(\text{SCN}^-\): Allocate electrons between S and N, ensuring full octets.- (c) \(\text{S}_2^{2-}\): Place 6 electrons on each S and 2 between them.- (d) \(\text{ClF}_2^+\): Place 6 electrons on each F and distribute remaining 4 around Cl.

Create Multiple Bonds If Necessary

Convert lone pairs to double (or triple) bonds to ensure the central atom fulfills the octet.- (a) \(\text{NO}_2^+\): Convert one lone pair from each O to a double bond with N.- (b) \(\text{SCN}^-\): Convert lone pairs to a triple bond between C and N and a double bond between C and S.- (c) \(\text{S}_2^{2-}\) and (d) \(\text{ClF}_2^+\) need no additional bonds.

Calculate Formal Charges

Calculate formal charges using the formula: Formal Charge = (Valence electrons) - (Non-bonding electrons) - (1/2 Bonding electrons).- (a) \(\text{NO}_2^+\): Both Oxygens have a formal charge of 0, and Nitrogen has a charge of +1.- (b) \(\text{SCN}^-\): Sulfur has a charge of 0, Carbon has a charge of 0, and Nitrogen has a charge of -1.- (c) \(\text{S}_2^{2-}\): Each sulfur atom has a charge of -1.- (d) \(\text{ClF}_2^+\): Each fluorine has a charge of 0, and Chlorine has a charge of +1.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding. They are the electrons involved in forming bonds with other atoms. Normally, the number of valence electrons for elements is determined by their group number in the periodic table. For instance, nitrogen which belongs to group 15 has five valence electrons. Understanding the number of valence electrons is the first step in drawing Lewis structures.

When determining the total number of valence electrons in a molecule or ion, it's important to consider the effect of any charge. A positive charge indicates a loss of electrons, while a negative charge suggests a gain. For example, in the ion ext{NO}_2^+ , you should subtract one electron from the total because of the positive charge.

• To calculate valence electrons, add the valence electrons for each atom involved.
• Add one electron for each negative charge.
• Subtract one electron for each positive charge.

Formal Charge

Formal charge helps determine the most stable Lewis structure by calculating the charge distribution within a molecule or ion. It helps decide the most likely structure among many possibilities. The formula for formal charge is:

Formal Charge = (Number of valence electrons) - (Number of non-bonding electrons) - (1/2 Number of bonding electrons)

Calculating formal charge is essential because it helps in validating the Lewis structure drawn by checking for the most stable arrangement of electrons. A stable molecule or ion will generally have a formal charge closer to zero. In many cases, formal charges combined in a valid Lewis structure will equal the overall charge of the ion.

When solving Lewis structures:

• All atoms strive for formal charges as close to zero as possible.
• Sum of all formal charges should equal the total charge of the molecule or ion.
• Configurations with negative charges on more electronegative atoms are more stable.

Chemical Bonds

Chemical bonds are forces that hold atoms together in a molecule or ion. They form as atoms share or transfer valence electrons. Lewis structures depict these bonds to show how atoms connect.

There are different types of chemical bonds, such as single, double, and triple bonds. In Lewis structures, a covalent single bond is represented by a single line between atoms, each line indicating a pair of shared electrons. When a molecule or ion requires additional electrons to satisfy the octet rule, atoms share more electrons, forming multiple bonds.

• Single bonds share one pair of electrons.
• Double bonds share two pairs of electrons.
• Triple bonds share three pairs of electrons.

Chemical bonds stabilize the structure by minimizing the repulsion between electrons, fulfilling the octet rule, or achieving lower energy states.

Octet Rule

The octet rule is a key principle that guides the drawing of Lewis structures. It states that atoms tend to form bonds in such a way that each atom has eight electrons in its valence shell, achieving a noble gas configuration.

Atoms like hydrogen are exceptions and follow the duet rule, aiming for just two electrons like helium.

• Central atoms in a molecule frequently adjust their bonding to satisfy their octet.
• Electrons are first placed to satisfy the octets of the more electronegative atoms (usually surrounding atoms) before the central atom.
• If the octet rule is not satisfied by simple sharing, multiple bonds can form.

It's crucial because it allows chemists to predict the structures of more complex molecules and ions. By satisfying the octet rule, molecules achieve maximum stability through optimal electron sharing or transfer via bonding.