Question

Of the following, the one that is a Bronsted-Lowry acid is (a) \(\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+} ;\) (b) \(\left[\mathrm{FeCl}_{4}\right]^{-} ;\) (c) \(\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\) (d) \(\left[\mathrm{Zn}(\mathrm{OH})_{4}\right]^{-}\).

Short answer

Therefore, the compound that acts as a Bronsted-Lowry acid is \(\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\).

Step-by-step solution

Analyze Cu(NH_3)_4^2+

Looking at \(\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\), we can observe that it doesn't have any hydrogen atom that can be easily donated. The hydrogen atoms are bound to the nitrogen atoms of the NH_3 ligands, and not readily available to be donated. Hence this does not act as a Bronsted-Lowry acid.

Analyze [FeCl_4]^-

The compound \(\left[\mathrm{FeCl}_{4}\right]^{-}\) does not have a hydrogen atom at all and so it cannot donate what it doesn't have. Hence, it is not a Bronsted-Lowry acid.

Analyze [Fe(H_2O)_6]^3+

The compound \(\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\) has six water molecules each with two hydrogen atoms. Upon losing a hydrogen ion (that is, a proton), it forms a hydroxide ion (OH^-). Hence, \(\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\) can act as a Bronsted-Lowry acid.

Analyze [Zn(OH)_4]^2-

The compound \(\left[\mathrm{Zn}(\mathrm{OH})_{4}\right]^{-}\) comprises hydroxide ions (OH^-), which typically act as bases as they can gain a proton but can't easily lose one. Hence it cannot act as a Bronsted-Lowry acid because it is not able to donate a proton.

Key concepts

Acid-Base Reactions

Understanding acid-base reactions is fundamental in chemistry, as these reactions are pervasive in both the laboratory and the natural world. The Bronsted-Lowry theory is one of the key models used to explain these reactions, focusing on the transfer of protons (hydrogen ions) between substances.

A Bronsted-Lowry acid is a substance that donates protons, whereas a base is a substance that accepts protons. This model expands the definition of acids and bases to include ions and molecules that might not contain hydroxide or hydrogen ions. In a typical acid-base reaction, the acid donates a proton to the base, resulting in the formation of a conjugate base and conjugate acid, reflecting the reversible nature of these reactions.

For instance, when a substance like hydrochloric acid (HCl) dissolves in water, it donates a proton to a water molecule, forming hydronium (H3O+) and a chloride ion (Cl-). This process is a classic example of an acid-base reaction. It's important to remember that the ability of a molecule or ion to donate a proton is influenced by its structure and the stability of the formed conjugate base.

Coordination Compounds

Coordination compounds play a significant role in many biological and industrial processes. These compounds consist of a central metal ion bonded to surrounding molecules or ions, known as ligands. The coordination number, or the number of ligand bonds to the metal ion, and the nature of the ligands can determine the compound's reactivity, color, and magnetic properties.

In the context of acid-base chemistry, it's important to note that not all coordination compounds can behave as acids. For a coordination compound to be a Bronsted-Lowry acid, it must contain ligands capable of donating a proton. In the exercise provided, compounds such as \(\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\) have ligands (NH3) that are not readily donating protons. In contrast, water ligands in \(\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\) can donate protons, showcasing that the identity and properties of the ligands are crucial for determining if the coordination compound is an acid.

Proton Donation

Proton donation is a key event in acid-base chemistry, and it is the defining characteristic of a Bronsted-Lowry acid. The ability to donate a proton depends on the acid's molecular structure and the stability of the resulting conjugate base.

When analyzing compounds for their potential as Bronsted-Lowry acids, it is critical to locate any hydrogen atoms that may be given up as protons. For instance, \(\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\), referred to in the exercise, can donate a proton due to the presence of labile hydrogen atoms in the water ligands. Upon donating a proton, the water molecule becomes a hydroxide ion (OH-). Contrastingly, if a compound such as \(\left[\mathrm{Zn}(\mathrm{OH})_{4}\right]^{-}\) already contains hydroxide ions, it is predisposed to accept, not donate, protons. Therefore, it behaves as a base rather than an acid.

Always consider the overall stability of the molecule after proton donation: The more stable the resulting conjugate base, the more likely the compound is to act as an acid. This stability can be influenced by various factors, including the presence of electron-withdrawing groups, which can stabilize the negative charge left behind after proton donation.