Question

The hypochlorite ion, \(\mathrm{ClO}^{-},\) acts as a weak base. (a) Is \(\mathrm{ClO}^{-}\) a stronger or weaker base than hydroxylamine? (b) When ClO \(^{-}\) acts as a base, which atom, Cl or \(\mathrm{O}\), acts as the proton acceptor? (c) Can you use formal charges to rationalize your answer to part (b)?

Short answer

(a) \(\mathrm{ClO}^{-}\) is a weaker base compared to hydroxylamine, as its conjugate acid (\(\mathrm{HClO}\)) is a stronger acid. (b) The oxygen atom in \(\mathrm{ClO}^{-}\) acts as the proton acceptor. (c) The formal charges support our choice of the oxygen atom as the proton acceptor in \(\mathrm{ClO}^{-}\), as the oxygen atom has a negative formal charge (-1) and the chlorine atom has a neutral formal charge (0).

Step-by-step solution

(a) Comparing basicity of ClO and hydroxylamine

To compare the basicity of hypochlorite ion (\(\mathrm{ClO}^{-}\)) and hydroxylamine, we should look at their conjugate acids (\(\mathrm{HClO}\) and \(\mathrm{NH_3OH^{+}}\)) and compare their acidity. The acidity can be compared by looking at the stability of the conjugate base. A more stable conjugate base means a stronger acid, and thus a weaker base. In general, more electronegative atoms can stabilize a negative charge better. For hydroxylamine, the conjugate base is \(\mathrm{NH_2OH}^{-}\) where the negative charge is on the nitrogen atom. In the case of hypochlorite ion, the conjugate base is \(\mathrm{HClO}\), and when losing a proton, it becomes \(\mathrm{ClO}^{-}\) with a negative charge on the oxygen atom. Oxygen is more electronegative than nitrogen. Therefore, \(\mathrm{ClO}^{-}\) is a weaker base compared to hydroxylamine, as its conjugate acid (\(\mathrm{HClO}\)) is a stronger acid.

(b) Identifying the proton acceptor atom

To identify which atom (Cl or O) in \(\mathrm{ClO}^{-}\) acts as the proton acceptor, we must consider the Lewis structure and resonance structures. The oxygen atom has one lone pair of electrons making it the most likely candidate for accepting a proton. Moreover, the oxygen atom is more electronegative than the chlorine atom, which means it is more likely to accept a proton as a Lewis base. So, the oxygen atom in \(\mathrm{ClO}^{-}\) acts as the proton acceptor.

(c) Rationalizing using formal charges

To rationalize our choice of the oxygen atom as the proton acceptor using formal charges, we can consider the formal charges of both atoms in \(\mathrm{ClO}^{-}\). The formal charge of an atom is calculated based on the valence electrons, the number of electrons in lone pairs, and half the number of shared electrons in bonds. The formal charge of oxygen in \(\mathrm{ClO}^{-}\) is: Formal charge (O) = Valence electrons (O) - Lone pair electrons (O) - (1/2) × Shared electrons (O) Formal charge (O) = 6 - 4 - (1/2) × 2 = -1 The formal charge of chlorine in \(\mathrm{ClO}^{-}\) is: Formal charge (Cl) = Valence electrons (Cl) - Lone pair electrons (Cl) - (1/2) × Shared electrons (Cl) Formal charge (Cl) = 7 - 6 - (1/2) × 2 = 0 Since the oxygen atom has a negative formal charge, it makes more sense for the oxygen atom to act as the proton acceptor rather than the chlorine atom, which has a neutral formal charge. Thus, the formal charges support our choice of the oxygen atom as the proton acceptor in \(\mathrm{ClO}^{-}\).

Key concepts

Hypochlorite Ion

The hypochlorite ion, represented as \( \mathrm{ClO}^{-} \), is an important chemical species often found in household bleach and disinfectants. It is composed of a chlorine atom bonded to an oxygen atom with an overall negative charge. The negative charge is primarily localized on the oxygen due to its higher electronegativity compared to chlorine. This characteristic makes \( \mathrm{ClO}^{-} \) a weak base in chemical reactions, meaning it can accept protons (H\(^+\)) but only to a limited extent.

The hypochlorite ion is notable for its oxidizing properties. It plays a crucial role in sanitation and cleaning processes due to its ability to break down and decompose organic and inorganic substances. Furthermore, \( \mathrm{ClO}^{-} \) can act as a disinfectant by destroying bacteria and viruses through chemical interactions that disrupt their normal function.

Lewis Structures

Lewis structures are a simple way to represent molecules and ions, showing how atoms are bonded and the distribution of electrons. For the hypochlorite ion \( \mathrm{ClO}^{-} \), we start by counting the total number of valence electrons available from the chlorine and oxygen atoms, then accounting for the additional electron due to the negative charge.

- Chlorine has 7 valence electrons.
- Oxygen has 6 valence electrons.
- The negative charge adds 1 extra electron, leading to a total of 14 valence electrons.

In the Lewis structure of \( \mathrm{ClO}^{-} \), the chlorine and oxygen atoms share a bond. Chlorine and oxygen each complete their octets with lone pairs of electrons. The negative charge is shown on the oxygen, which is more stable with an extra electron compared to chlorine.

Drawing accurate Lewis structures allows us to predict the behavior of molecules in chemical reactions, as it demonstrates possible interactions between lone pairs and other atoms or ions.

Formal Charge

Formal charge is a useful concept in evaluating the correct Lewis structure and predicting how atoms will behave in a reaction. To calculate the formal charge for an atom in a molecule, use the formula:

\[ \text{Formal Charge} = \text{Valence Electrons} - \text{Non-bonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons} \]

Applying this to the hypochlorite ion \( \mathrm{ClO}^{-} \):
- For oxygen: It has 6 valence electrons, shares one with chlorine, and has 4 as lone pairs. The formal charge is calculated as 6 - 4 - 1 = -1.
- For chlorine: It has 7 valence electrons, shares one with oxygen, and has 6 as lone pairs. The formal charge is 7 - 6 - 1 = 0.

The distribution of formal charges helps justify why in \( \mathrm{ClO}^{-} \), oxygen acts as the proton acceptor in reactions. The negative charge on oxygen makes it more likely to bond with positively charged protons, fitting its role as a Lewis base.

Acid-Base Chemistry

Acid-base chemistry is crucial in understanding chemical reactions involving the transfer of protons (H\(^+\)). An acid donates a proton, while a base accepts a proton. The hypochlorite ion \( \mathrm{ClO}^{-} \) acts as a base, which means it can accept a proton.

In many acid-base reactions, it's important to compare the strength of acids and bases. The strength is determined by the stability of their conjugate pairs. For instance, comparing \( \mathrm{ClO}^{-} \) with hydroxylamine \( \mathrm{NH_2OH} \), one checks the stability of \( \mathrm{HClO} \) and \( \mathrm{NH_3OH^{+}} \). More electronegative atoms stabilize negative charges better, so electrolyte differences (oxygen's higher electronegativity than nitrogen) affect stability, influencing basicity.

Understanding acid-base behavior, including which atom in a base is more likely to accept a proton, allows chemists to predict material's interactions in various environments, impacting practical applications such as cleaning processes and the formulation of disinfecting agents.