Problem 37 - Chemistry and Chemical Reactivity

Question

The chemistry of the nitrite ion and HNO \(_{2}:\) (a) Two resonance structures are possible for \(\mathrm{NO}_{2}^{-}\). Draw these structures, and then find the formal charge on each atom in each resonance structure. (b) If an \(\mathrm{H}^{+}\) ion is attached to \(\mathrm{NO}_{2}^{-}\) (to form the acid \(\left.\mathrm{HNO}_{2}\right),\) it attaches to the \(\mathrm{O}\) atom and not the \(\mathrm{N}\) atom. Explain why you would predict this structure. (c) Two resonance structures are possible for HNO,. Draw these structures, and then find the formal charge on each atom in each resonance structure. Is either of these structures strongly preferred over the other?

Step-by-Step Solution

Answer

NO₂⁻ connects H⁺ to O due to higher electronegativity. Neither HNO₂ structure is strongly preferred.

Step 1: Drawing the Resonance Structures for NO2^-

There are two possible resonance structures for the nitrite ion (NO₂⁻). In each structure, a nitrogen atom is bonded to two oxygen atoms. The nitrogen bears an unshared pair of electrons. In the first structure, one oxygen forms a double bond with nitrogen, and the other oxygen has a single bond. In the second structure, the double and single bonds are swapped between the two oxygen atoms.

Step 2: Calculating Formal Charges for NO2^-

The formal charge is calculated using the formula: \( \text{Formal charge} = \text{Valence electrons} - (\text{Lone pair electrons} + \frac{1}{2} \text{Bonded electrons}) \). Each nitrogen in the resonance structures has a formal charge of 0. Each oxygen that is double-bonded to nitrogen also has a formal charge of 0, while the single-bonded oxygen has a formal charge of -1.

Step 3: Attaching H+ Ion to NO2^-

The hydrogen ion (H⁺) attaches preferentially to an oxygen atom rather than a nitrogen atom because oxygen is more electronegative and carries a partial negative charge, which makes it more receptive to the positively charged hydrogen ion.

Step 4: Drawing Resonance Structures for HNO2

For HNO₂, two resonance structures are possible. In both structures, hydrogen is attached to one of the oxygen atoms. In the first resonance structure, the nitrogen is double-bonded to one oxygen and single-bonded to the other (which also hosts the hydrogen). In the second structure, the double bond is with the oxygen that has the hydrogen.

Step 5: Calculating Formal Charges for HNO2

Using the formal charge formula from Step 2, we find that the formal charges are equivalent in both resonance structures. The nitrogen atom in each structure typically has a formal charge of +1, the oxygen involved in a double bond has a formal charge of 0, and the other bonded oxygen has a formal charge of 0 when adjusted for hydrogen attachment.

Step 6: Assessing Preferability of Resonance Structures for HNO2

Neither of the resonance structures for HNO₂ is strongly preferred over the other because they both have similar formal charges and stability. This means the actual structure of HNO₂ is a hybrid of both resonance forms.

Key Concepts

Formal Charge

Understanding formal charge is crucial for grasping resonance structures. Formal charge helps determine the most likely or stable structure of a molecule. To calculate the formal charge for an atom, use the formula: \[ \text{Formal charge} = \text{Valence electrons} - (\text{Lone pair electrons} + \frac{1}{2} \text{Bonded electrons}) \] This formula helps identify any charge on an atom in a molecule. It's like bookkeeping electrons!For instance, in a nitrite ion (NO₂⁻), we count the lone pairs and the number of bonds each atom participates in. The goal is to make sure that the sum of formal charges in a molecule matches the overall charge of the molecule itself. - A nitrogen atom in a typical resonance structure might exhibit a formal charge of 0. - Meanwhile, oxygen might show differing formal charges based on whether it forms single or double bonds.

Nitrite Ion

The nitrite ion, represented as \( \text{NO}_2^- \), is an interesting anion due to its resonance. It consists of a nitrogen atom centrally bonded to two oxygen atoms both at different lengths. This difference arises because of resonance.Resonance occurs when two or more structures with different electron arrangements exist for the same molecule. These structures are not real but a way for scientists to depict electron distribution.In nitrite, each valid resonance structure involves: - A nitrogen atom bonded to one oxygen via a double bond. - The second oxygen attached through a single bond. - The resonance structures swap these bonds to show the electron flow.Because of its resonance, NO₂⁻ exhibits properties that are a blend of all its resonance forms.

Chemical Bonding

Chemical bonding is what holds atoms together. It can occur in different forms: covalent, ionic, and metallic bonds are the main types. In resonance structures, especially in the nitrite ion, we primarily see covalent bonds. These bonds occur when atoms share electrons to achieve a more stable electronic arrangement known as the octet rule. What makes resonance unique is how these covalent bonds, particularly within molecules like the nitrite ion, can "flip" or "alternate" without actually changing the molecule's physical structure. - Double bonds are usually stronger and shorter than single bonds. - In resonance, you might find these bonds swapping places, but the bond length is a mix between those of a single and double bond. Understanding these bonds helps students predict molecular structure and reactivity.

Electronegativity

Electronegativity is a crucial concept for predicting how molecules behave. It describes an atom's ability to attract shared electrons in a bond. When looking at nitrite or other molecules with resonance, electronegativity guides where electrons are likely to "settle" in the molecule: - Oxygen, being more electronegative than nitrogen, tends to attract more electron density. - This property is why, in the attachment of H⁺ ions to NO₂⁻, the proton prefers to bond to an oxygen atom. Simple because oxygen retains greater negative character due to its electronegativity. In chemistry, understanding electronegativity helps predict molecular shapes, bond strengths, and even the angles between bonds in complex molecules.