Question

Draw Lewis structures for the following ions: (a) \(\mathrm{O}_{2}^{2-}\), (b) \(\mathrm{C}_{2}^{2-}\) (c) \(\mathrm{NO}^{+}\) (d) \(\mathrm{NH}_{4}^{+}\). Show formal charges.

Short answer

Lewis structures are: (a) \( :O-O^{2-}: \), (b) \( :C\equiv C^{2-}: \), (c) \( :N\equiv O^{+}: \), (d) \( \mathrm{NH}_{4}^{+} \) with +1 on N.

Step-by-step solution

Lewis Structure of Peroxide Ion \( \mathrm{O}_{2}^{2-} \)

1. Calculate total valence electrons: Oxygen has 6 valence electrons, so \(2 \times 6 = 12\) electrons. Add 2 for the -2 charge, giving 14 electrons.2. Place the two oxygen atoms side by side.3. Draw a single bond between the oxygen atoms, accounting for 2 electrons. Distribute the remaining 12 electrons to complete the octets.4. Each oxygen receives 3 lone pairs.5. Calculate formal charges: \( \text{Formal charge} = \text{valence electrons} - \text{lone pair electrons} - \frac{1}{2}\text{bonding electrons} \). - For each oxygen: \(6 - 6 - 1 = -1\).6. Result: \[ \begin{array}{ c } : \cdot \cdot \ddot O - O^{2-} \cdot \cdot : \ \end{array} \]

Lewis Structure of \( \mathrm{C}_{2}^{2-} \)

1. Calculate total valence electrons: Carbon has 4 valence electrons, so \(2 \times 4 = 8\) electrons. Add 2 for the -2 charge, giving 10 electrons.2. Place the two carbon atoms side by side.3. Draw a triple bond between the carbon atoms, accounting for 6 electrons. Provide each carbon with one lone pair.4. Calculate formal charges: \( \text{Formal charge} = \text{valence electrons} - \text{lone pair electrons} - \frac{1}{2}\text{bonding electrons} \). - For each carbon: \(4 - 2 - 3 = -1\).5. Result: \[ \begin{array}{ c } : \cdot \cdot C \equiv C^{2-} \cdot \cdot : \end{array} \]

Lewis Structure of Nitrosonium Ion \( \mathrm{NO}^{+} \)

1. Calculate total valence electrons: Nitrogen has 5 valence electrons and oxygen has 6, so \(5 + 6 = 11\). Subtract 1 for the +1 charge, resulting in 10 electrons.2. Place nitrogen and oxygen side by side, nitrogen can be the central atom.3. Draw a triple bond between nitrogen and oxygen using 6 electrons. Assign remaining electrons as lone pairs to fill octets.4. For formal charges: - Nitrogen \(\text{Formal charge} = 5 - 2 - 3 = 0\). - Oxygen \(\text{Formal charge} = 6 - 4 - 3 = +1\).5. Result:\[ \begin{array}{ c } :N \equiv O^{+}: \cdot \cdot \end{array} \]

Lewis Structure of Ammonium Ion \( \mathrm{NH}_{4}^{+} \)

1. Calculate total valence electrons: Nitrogen has 5 valence electrons, hydrogen has 1 each, so \(5 + 4 \times 1 = 9\). Subtract 1 for the +1 charge, leading to 8 electrons.2. Place nitrogen in the center, surround with 4 hydrogens.3. Draw a single bond from nitrogen to each hydrogen, using 8 electrons.4. All hydrogens have 2 electrons each (full), and nitrogen has 4 single bonds.5. Calculate formal charges: - Nitrogen \(\text{Formal charge} = 5 - 0 - 4 = +1\). - Hydrogens are neutral.6. Result: \[ \begin{array}{ c } \begin{array}{ c } H ewline N^{+} ewline H \end{array} \begin{array}{ c } H ewline H ewline \end{array}\end{array} \]

Key concepts

Formal Charges

The concept of formal charge is essential for understanding molecular structures and their stability. To find the formal charge on an atom within a molecule or ion, you can use the formula: \[ \text{Formal charge} = \text{valence electrons} - \text{lone pair electrons} - \frac{1}{2} \times \text{bonding electrons} \]This formula accounts for both the electrons owned by an atom in lone pairs and its shared electrons in chemical bonds. Formal charges allow chemists to determine the most likely structure for a molecule or ion, helping to predict reactivity and properties. In Lewis structures, atoms ideally have formal charges as close to zero as possible, as this indicates stability. If there are charges, they should align with the atom's electronegativity: more electronegative atoms are more likely to bear negative charges.

Valence Electrons

Valence electrons play a crucial role in determining how atoms bond with one another. They are the outermost electrons in an atom that participate in chemical bonding. You can determine the number of valence electrons by looking at the group number of the element in the periodic table. For example:

• Oxygen, in Group 16, has 6 valence electrons.
• Carbon, in Group 14, has 4 valence electrons.
• Nitrogen, in Group 15, has 5 valence electrons.
• Hydrogen, in Group 1, has 1 valence electron.

Knowing how many valence electrons are available helps you in drawing Lewis structures, as you need to distribute these electrons to create bonds and achieve full outer shells (or octets) for each atom in a molecule.

Chemical Bonds

Chemical bonds form when atoms share or transfer valence electrons to achieve stable electronic configurations, often resembling the noble gases. There are several types of chemical bonds:

• **Covalent bonds**, where atoms share electron pairs. They can be single (one pair shared), double (two pairs shared), or triple bonds (three pairs shared).
• **Ionic bonds**, where electrons are transferred from one atom to another, forming ions that attract each other.

In Lewis structures, these bonds are represented by lines between atoms. For example, in the peroxide ion \(\mathrm{O}_{2}^{2-}\), a single bond connects the two oxygen atoms. Conversely, in the \(\mathrm{C}_{2}^{2-}\) ion, the carbon atoms are linked by a triple bond, signifying the sharing of three electron pairs.

Ions

Ions are atoms or molecules that have lost or gained electrons to achieve a full electron shell. They exist as cations (positively charged) or anions (negatively charged). In the context of Lewis structures and chemical bonding:

• Cations, like \(\mathrm{NO}^{+}\) and \(\mathrm{NH}_{4}^{+}\), have fewer electrons than protons, resulting in a positive charge. When drawing them, ensure you subtract electrons for the positive charge.
• Anions, such as \(\mathrm{O}_{2}^{2-}\) and \(\mathrm{C}_{2}^{2-}\), have more electrons than protons, resulting in a negative charge. Add extra electrons to account for the additional charge.

This gained or lost electron property significantly affects the reactivity and interaction of ions with other species in chemical reactions.