Question

(a) Is a weak Bronsted-Lowry base necessarily a weak Lewis base? Explain with an example. (b) Identify the Lewis bases in the following reaction: $$ \mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}^{2+}(a q)+4 \mathrm{CN}^{-}(a q) \rightleftharpoons \mathrm{Cu}(\mathrm{CN})_{4}^{2-}(a q)+4 \mathrm{H}_{2} \mathrm{O}(l) $$ (c) Given that \(K_{c}>1\) for the reaction in part (b), which Lewis base is stronger?

Short answer

A weak Bronsted-Lowry base is not necessarily a weak Lewis base. In the reaction, \mathrm{H}_{2} \mathrm{O} and \mathrm{CN}^{-} are Lewis bases. \mathrm{CN}^{-} is the stronger Lewis base.

Step-by-step solution

Understanding Bronsted-Lowry and Lewis Bases

A Bronsted-Lowry base is a substance that can accept a proton (H+), while a Lewis base is a substance that can donate an electron pair. Both concepts are related but not identical.

Example of Weak Bronsted-Lowry and Lewis Base

Consider ammonia (NH3). It is a weak Bronsted-Lowry base because it does not completely ionize in water. However, it is not necessarily a weak Lewis base because it readily donates its lone pair of electrons to a proton to form NH4+. Thus, a weak Bronsted-Lowry base is not necessarily a weak Lewis base.

Identifying Lewis Bases in the Given Reaction

In the reaction \ \[\mathrm{Cu} \left( \mathrm{H}_{2} \mathrm{O} \right)_{4}^{2+}(aq) + 4 \mathrm{CN}^{-}(aq) \rightleftharpoons \mathrm{Cu} (\mathrm{CN})_{4}^{2-}(aq) + 4 \mathrm{H}_{2} \mathrm{O}(l)\], the Lewis bases are the substances that donate an electron pair. The molecules \mathrm{H}_{2} \mathrm{O} and \mathrm{CN}^{-} are both acting as Lewis bases by donating their electron pairs to the \mathrm{Cu}^{2+} ion.

Determining the Stronger Lewis Base

Given that \( \mathrm{K_c} > 1 \) for the reaction, it means the equilibrium favors the products, indicating that \mathrm{CN}^{-} is a stronger Lewis base relative to \( \mathrm{H}_{2} \mathrm{O} \). The stronger attraction between \mathrm{Cu}^{2+} and \mathrm{CN}^{-} shifts the equilibrium to the right.

Key concepts

Acid-Base Chemistry

Acid-base chemistry explores the interactions between acids and bases. In this context, an acid is a substance that can donate a proton (H+), and a base is a substance that can accept a proton. The Bronsted-Lowry theory expands on this by defining acids and bases based on their ability to donate or accept protons. This theory helps in understanding various reactions where proton transfer occurs.
For example, when ammonia (NH3) reacts with water, it accepts a proton to form ammonium ions (NH4+), showcasing its property as a Bronsted-Lowry base.
Lewis acids and bases offer a more general approach. A Lewis base is defined as a substance that can donate an electron pair, whereas a Lewis acid can accept an electron pair. This broadened scope includes many reactions not covered by the Bronsted-Lowry theory.
By examining both definitions, we can better understand and predict the behavior of substances in different chemical reactions.

Lewis Bases

A Lewis base is any species that donates an electron pair to form a bond. This definition is wider than the Bronsted-Lowry base concept because it includes species that donate electron pairs without necessarily involving protons.
For example, in the provided reaction:
\[\text{Cu}\text{(}\text{H}_{2}\text{O}\text{)}_{4}^{2+}(aq) + 4 \text{CN}^-(aq) \rightleftharpoons \text{Cu}\text{(CN)}_{4}^{2-}(aq) + 4 \text{H}_{2}\text{O}(l)\],
both H2O and CN- act as Lewis bases because they both donate electron pairs to Cu2+.
In some situations, a substance can be a weak Bronsted-Lowry base but a strong Lewis base. For instance, ammonia (NH3) is a weak Bronsted-Lowry base when it doesn't fully ionize in water. However, it acts as a strong Lewis base by readily donating its lone pair of electrons.
This difference is crucial in understanding different chemical reactions and predicting the strength of bases in various contexts.

Chemical Equilibrium

Chemical equilibrium refers to the state where the concentrations of reactants and products in a reversible reaction remain constant over time. This occurs when the forward and reverse reaction rates are equal.
When considering an equilibrium such as:
\[\text{Cu}\text{(}\text{H}_{2}\text{O}\text{)}_{4}^{2+}(aq) + 4 \text{CN}^-(aq) \rightleftharpoons \text{Cu}\text{(CN)}_{4}^{2-}(aq) + 4\text{H}_{2}\text{O}(l)\],
If the equilibrium constant, Kc, is greater than 1, it indicates that the products are favored over the reactants. This means that the interaction between Cu2+ and CN- is stronger, shifting the equilibrium to the right.
In this reaction, since Kc > 1, CN- is a stronger Lewis base than H2O. The stronger Lewis base binds more effectively with the Lewis acid (Cu2+), favoring the formation of the product (Cu(CN)4^2-).
Understanding chemical equilibrium helps in predicting the outcome of reactions based on the relative strengths of acids and bases involved.