Question

The hypochlorite ion, \(\mathrm{ClO}^{-},\) acts as a weak base. (a) Is \(\mathrm{ClO}^{-},\) a stronger or weaker base than hydroxylamine? (b) When \(\mathrm{ClO}^{-}\) 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) ClO^- is a stronger base than hydroxylamine, as its conjugate acid (HClO) has a higher pKa value than hydroxylamine's conjugate acid (NH3OH^+). (b) The oxygen atom in ClO^- is the proton acceptor when it acts as a base, due to its greater electronegativity and presence of lone pairs of electrons. (c) Formal charges rationalize the role of oxygen as the proton acceptor, as oxygen has a negative formal charge, making it more likely to act as the base.

Step-by-step solution

Determine acidity of ClO^- and hydroxylamine

To compare the basicity of ClO^- and hydroxylamine, we must first analyze their acidity. The relative strengths of bases can be determined by looking at their conjugate acids. A weaker conjugate acid means a stronger base. The conjugate acid of ClO^- is hypochlorous acid (HClO), and the pKa value of HClO is 7.53. Hydroxylamine has the formula NH2OH, and its conjugate acid is NH3OH^+. The pKa value for NH3OH^+ is 5.99.

Compare basicity of ClO^- and hydroxylamine

Given the pKa values of the conjugate acids, we can now compare their basicity. A lower pKa value corresponds to a stronger conjugate acid and a weaker conjugate base. Thus, as the pKa of HClO is greater than the pKa of NH3OH^+, this implies that ClO^- is a stronger base than hydroxylamine. Hence, ClO^- is a stronger base than hydroxylamine.

Identify the proton acceptor in ClO^-

To answer which atom in ClO^- acts as the proton acceptor, we must consider the Lewis structures of the molecule. A Lewis base is a compound that can donate a lone pair of electrons to a Lewis acid. The atom that has the most electronegative character and a lone pair of electrons is likely to act as proton acceptor. Oxygen is more electronegative than chlorine. In the ClO^- ion (O-Cl with 3 lone pairs around Oxygen and one negative charge on Oxygen), the oxygen atom has a lone pair of electrons and a negative charge. So, when ClO^- ion acts as a base, the oxygen atom plays the role of proton acceptor.

Rationalize the role of oxygen using formal charges

Formal charges can be used to rationalize the role of oxygen as proton acceptor in the ClO^- ion, as they provide a way of assessing the most stable Lewis structure for the ion. The formal charge of an atom, FC(A), is given by: FC(A) = Valence electrons of A - 0.5*Number of bonding electrons - Number of non-bonding electrons. On calculating formal charges for ClO^- (O-Cl), we get: FC(O) = 6 - 0.5*(4) - 6 = -1 FC(Cl) = 7 - 0.5*(4) - 6 = +1 The formal charges on oxygen indicate that it has a negative charge, and positive formal charges are present on chlorine. This means that the oxygen atom, being more electronegative, will more likely act as the proton acceptor (base) because of its negative formal charge. In conclusion, (a) ClO^- is a stronger base than hydroxylamine (b) Oxygen atom in ClO^- is the proton acceptor when it acts as a base (c) Formal charges can be used to rationalize the answer to part (b), as the negative formal charge on oxygen indicates that it will act as the proton acceptor.

Key concepts

Hypochlorite Ion as a Base

Understanding the role of the hypochlorite ion (\( \text{ClO}^- \) ) in chemical reactions is essential for students of chemistry. As a commonly encountered ion in both laboratory and environmental settings, it serves as an archetype in illustrating basic principles of acid-base chemistry. The hypochlorite ion is a weak base, which means it has a propensity to accept protons (\( \text{H}^+ \) ions) in aqueous solutions.

It's worth noting that the strength of a base can depend on various factors such as the compound's structure, the stability of its conjugate acid, and the specific medium in which the reaction occurs. In the context of hydroxylamine, we see that the hypochlorite ion is a stronger base due to its higher associated pKa value. This hints at an intrinsic stability of the hypochlorous acid (\( \text{HClO} \) ) compared to the conjugate acid of hydroxylamine (\( \text{NH}_3\text{OH}^+ \) ), emphasizing the utility of pKa as a comparative measure for basicity.

Lewis Structure and Proton Acceptance

Delving into the molecular world, the Lewis structure is a straightforward yet vital tool for visualizing molecules and predicting behavior. It essentially maps out how electrons are shared or paired up around atoms in a molecule. Now, when it comes to proton acceptance, the Lewis structure of the hypochlorite ion reveals that the oxygen atom is the site for this action.

What's compelling about oxygen's role here is its higher electronegativity in comparison to chlorine, granting it a greater affinity for electrons — including the potential to acquire a proton. A hallmark of oxygen's aptitude for proton acceptance is its ownership of a formal negative charge, as depicted in the Lewis structure, serving as a telltale beacon in the ion's reactivity profile.

Comparative pKa Analysis

The concept of pKa is fundamental to understanding acid-base chemistry. The pKa value indicates the acid dissociation constant, providing insight into the strength of an acid — and conversely the strength of its conjugate base. More explicitly, a lower pKa value signifies a stronger acid and therefore a weaker conjugate base, and vice versa.

In comparing the pKa values between hypochlorous acid and the conjugate acid of hydroxylamine, it becomes evident that hypochlorous acid, with its higher pKa, forms a stronger base upon deprotonation. This comparative pKa analysis is invaluable for predicting the outcome of acid-base reactions and for rationalizing the observed reactivity of related compounds.

Formal Charge Calculation

Grasping the concept of formal charge in molecular structures is a critical piece in the puzzle of chemical reactivity. It's calculated by taking the number of valence electrons of an atom, subtracting half the number of bonding electrons, and then subtracting the number of non-bonding electrons. This mathematical construct affords chemists a tool to gauge the most likely arrangement of electrons in a molecule — pivotal in understanding the molecule's behavior.

For the hypochlorite ion, the oxygen atom carries a negative formal charge, forecasting its likelihood to act as a Lewis base by accepting a proton. Such predictions grounded in formal charge analysis prove indispensable for explaining the nuances of reactive encounters in the molecular realm.