Chemical structures of organic compounds and ions.

Problem Analysis:

The task appears to involve identifying the formal charge on specific atoms in various molecular structures. Formal charge is a concept used in chemistry to determine the distribution of electrons in a molecule or ion. The formal charge on an atom can be calculated using the following formula:

\[
\text{Formal Charge} = \text{Valence Electrons} - \left( \text{Non-bonding Electrons} + \frac{\text{Bonding Electrons}}{2} \right)
\]

Where:
- Valence Electrons: The number of valence electrons for the atom in its elemental form.
- Non-bonding Electrons: The number of lone pair electrons on the atom.
- Bonding Electrons: The total number of electrons involved in covalent bonds.

Solution:

We will calculate the formal charge for each labeled atom in the given structures.

---

#### a.
Structure:
\[
\ce{CH3+}
\]
- Carbon (C) has 4 valence electrons.
- Non-bonding electrons: 0 (no lone pairs).
- Bonding electrons: 6 (3 single bonds, each contributing 2 electrons).
- Formal charge on C:
\[
\text{Formal Charge} = 4 - \left( 0 + \frac{6}{2} \right) = 4 - 3 = +1
\]

#### b.
Structure:
\[
\ce{CH3-}
\]
- Carbon (C) has 4 valence electrons.
- Non-bonding electrons: 2 (one lone pair).
- Bonding electrons: 6 (3 single bonds, each contributing 2 electrons).
- Formal charge on C:
\[
\text{Formal Charge} = 4 - \left( 2 + \frac{6}{2} \right) = 4 - 5 = -1
\]

#### c.
Structure:
\[
\ce{CH3-}
\]
- This is the same as structure b. The formal charge on C is:
\[
\text{Formal Charge} = -1
\]

#### d.
Structure:
\[
\ce{D-O=O}
\]
- Oxygen (O) with a double bond has 6 valence electrons.
- Non-bonding electrons: 4 (two lone pairs).
- Bonding electrons: 4 (one double bond, contributing 4 electrons).
- Formal charge on O:
\[
\text{Formal Charge} = 6 - \left( 4 + \frac{4}{2} \right) = 6 - 6 = 0
\]
- Deuterium (D) has 1 valence electron.
- Non-bonding electrons: 0.
- Bonding electrons: 2 (one single bond).
- Formal charge on D:
\[
\text{Formal Charge} = 1 - \left( 0 + \frac{2}{2} \right) = 1 - 1 = 0
\]

#### e.
Structure:
\[
\ce{CH3-NH2}
\]
- Nitrogen (N) has 5 valence electrons.
- Non-bonding electrons: 2 (one lone pair).
- Bonding electrons: 8 (three single bonds, each contributing 2 electrons).
- Formal charge on N:
\[
\text{Formal Charge} = 5 - \left( 2 + \frac{8}{2} \right) = 5 - 6 = -1
\]

#### f.
Structure:
\[
\ce{CH3-NH2}
\]
- This is the same as structure e. The formal charge on N is:
\[
\text{Formal Charge} = -1
\]

#### g.
Structure:
\[
\ce{CH3-O-CH3}
\]
- Oxygen (O) has 6 valence electrons.
- Non-bonding electrons: 6 (three lone pairs).
- Bonding electrons: 4 (one single bond and one double bond, contributing 4 electrons).
- Formal charge on O:
\[
\text{Formal Charge} = 6 - \left( 6 + \frac{4}{2} \right) = 6 - 8 = -2
\]

#### h.
Structure:
\[
\ce{CH3-P(CH3)3}
\]
- Phosphorus (P) has 5 valence electrons.
- Non-bonding electrons: 0.
- Bonding electrons: 12 (four single bonds, each contributing 2 electrons).
- Formal charge on P:
\[
\text{Formal Charge} = 5 - \left( 0 + \frac{12}{2} \right) = 5 - 6 = -1
\]

#### i.
Structure:
\[
\ce{CH3-S-CH3}
\]
- Sulfur (S) has 6 valence electrons.
- Non-bonding electrons: 6 (three lone pairs).
- Bonding electrons: 4 (one single bond and one double bond, contributing 4 electrons).
- Formal charge on S:
\[
\text{Formal Charge} = 6 - \left( 6 + \frac{4}{2} \right) = 6 - 8 = -2
\]

#### j.
Structure:
\[
\ce{CH3-C#N-H}
\]
- Carbon (C) has 4 valence electrons.
- Non-bonding electrons: 0.
- Bonding electrons: 8 (one triple bond and two single bonds, contributing 8 electrons).
- Formal charge on C:
\[
\text{Formal Charge} = 4 - \left( 0 + \frac{8}{2} \right) = 4 - 4 = 0
\]

#### k.
Structure:
\[
\ce{CH3-NO}
\]
- Nitrogen (N) has 5 valence electrons.
- Non-bonding electrons: 0.
- Bonding electrons: 6 (one double bond and one single bond, contributing 6 electrons).
- Formal charge on N:
\[
\text{Formal Charge} = 5 - \left( 0 + \frac{6}{2} \right) = 5 - 3 = +2
\]

#### l.
Structure:
\[
\ce{CH3-CN}
\]
- Carbon (C) has 4 valence electrons.
- Non-bonding electrons: 0.
- Bonding electrons: 8 (one triple bond and one single bond, contributing 8 electrons).
- Formal charge on C:
\[
\text{Formal Charge} = 4 - \left( 0 + \frac{8}{2} \right) = 4 - 4 = 0
\]

#### m.
Structure:
\[
\ce{CH3-NO2}
\]
- Nitrogen (N) has 5 valence electrons.
- Non-bonding electrons: 0.
- Bonding electrons: 6 (one double bond and one single bond, contributing 6 electrons).
- Formal charge on N:
\[
\text{Formal Charge} = 5 - \left( 0 + \frac{6}{2} \right) = 5 - 3 = +2
\]

#### n.
Structure:
\[
\ce{H2C=N=N}
\]
- Nitrogen (N) has 5 valence electrons.
- Non-bonding electrons: 0.
- Bonding electrons: 6 (one double bond and one single bond, contributing 6 electrons).
- Formal charge on N:
\[
\text{Formal Charge} = 5 - \left( 0 + \frac{6}{2} \right) = 5 - 3 = +2
\]

#### o.
Structure:
\[
\ce{H2N-NH2}
\]
- Nitrogen (N) has 5 valence electrons.
- Non-bonding electrons: 2 (one lone pair).
- Bonding electrons: 6 (two single bonds and one double bond, contributing 6 electrons).
- Formal charge on N:
\[
\text{Formal Charge} = 5 - \left( 2 + \frac{6}{2} \right) = 5 - 5 = 0
\]

---

Final Answer:

\[
\boxed{
\begin{array}{ll}
\text{a. } +1 & \text{b. } -1 & \text{c. } -1 \\
\text{d. } 0 & \text{e. } -1 & \text{f. } -1 \\
\text{g. } -2 & \text{h. } -1 & \text{i. } -2 \\
\text{j. } 0 & \text{k. } +2 & \text{l. } 0 \\
\text{m. } +2 & \text{n. } +2 & \text{o. } 0 \\
\end{array}
}
\]