Question

Explain the following observations: (a) \(\mathrm{HCl}\) is a stronger acid than \(\mathrm{H}_{2} \mathrm{~S} ;\) (b) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) is a stronger acid than \(\mathrm{H}_{3} \mathrm{AsO}_{4}\); (c) \(\mathrm{HBrO}_{3}\) is a stronger acid than \(\mathrm{HBrO}_{2}\); (d) \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) is a stronger acid than \(\mathrm{HC}_{2} \mathrm{O}_{4}{ }^{-} ;(\mathrm{e})\) benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right)\) is a stronger acid than phenol \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}\right)\).

Short answer

In summary, HCl is stronger than H2S due to its higher electronegativity, H3PO4 is stronger than H3AsO4 since phosphorus is more electronegative than arsenic, HBrO3 is stronger than HBrO2 because of the additional oxygen atom stabilizing the conjugate base, H2C2O4 is stronger than HC2O4- due to the neutral charge making it easier to lose a proton, and benzoic acid (C6H5COOH) is stronger than phenol (C6H5OH) because of the carboxylic acid group providing resonance stabilization.

Step-by-step solution

(1. Understanding acidity definition)

Acidity refers to the ability of a compound to lose protons in the form of hydrogen ions (H+) in a solution. The more easily a compound loses a proton, the stronger the acid it will be.

(2. Comparing HCl vs H2S)

HCl is a stronger acid than H2S. This is due to the fact that chlorine (Cl) is more electronegative than sulfur (S). This means that chlorine pulls electrons towards itself more than sulfur does. As a result, HCl produces a more stable conjugated base (Cl-) which makes it easier to lose the proton, making HCl a stronger acid.

(3. Comparing H3PO4 vs H3AsO4)

H3PO4 is a stronger acid than H3AsO4. Phosphorus (P) and arsenic (As) are found in the same group of the periodic table. However, since phosphorus is more electronegative than arsenic, it can stabilize the negative charge on the conjugated base more effectively. Thus, H3PO4 is more likely to lose a proton compared to H3AsO4, making it a stronger acid.

(4. Comparing HBrO3 vs HBrO2)

HBrO3 is a stronger acid than HBrO2 because it has more oxygen atoms which help to stabilize the negative charge of the conjugated base. The additional oxygen atom in HBrO3 contributes to electron delocalization of the conjugate base (BrO3-) which lowers the energy of the resulting anion and causes HBrO3 to more readily lose a proton, making it a stronger acid.

(5. Comparing H2C2O4 vs HC2O4-)

H2C2O4 is a stronger acid than HC2O4- because the neutral molecule, H2C2O4, can lose a proton more easily due to its neutral charge. The negatively charged compound, HC2O4-, has a more difficult time losing a proton because doing so would cause it to become even more negatively charged, resulting in a less stable conjugate base.

(6. Comparing C6H5COOH vs C6H5OH)

Benzoic acid (C6H5COOH) is a stronger acid than phenol (C6H5OH) because of the presence of the carboxylic acid group (COOH) in benzoic acid. The carboxylic acid group is more acidic than the hydroxyl group (OH) in phenol due to resonance stabilization and greater electron delocalization in the benzoate ion. When benzoic acid loses a proton, the resulting conjugate base (benzoate ion) is stabilized by resonance structures, making it easier to lose a proton as compared to phenol.

Key concepts

Acid Strength Comparison

Understanding acid strength is essential in chemistry as it helps explain why some substances can donate protons more readily than others. For instance, hydrochloric acid (HCl) is stronger than hydrogen sulfide (H₂S) because of the elements involved. Chlorine in HCl is more electronegative than sulfur in H₂S, enabling HCl to stabilize its conjugate base better and release protons with ease.

Similarly, phosphoric acid (H₃PO₄) is stronger than arsenic acid (H₃AsO₄) because the extra electronegativity of phosphorus aids in the stabilization of the conjugate base, promoting proton donation. This comparison of acid strengths is critical to understanding reactions in organic and inorganic chemistry.

Electronegativity and Acidity

Electronegativity plays a pivotal role in determining the acidity of compounds. A higher electronegative atom attached to a hydrogen atom will attract the bonding electrons closer, thus weakening the H–atom bond. This makes the release of the hydrogen ion (proton) easier. For example, chlorine (Cl) in HCl has a higher electronegativity compared to sulfur (S) in H₂S. This is why HCl is much more acidic. Electronegativity helps us predict and compare the acidities of different compounds within the same family, enhancing our understanding of chemical behavior.

Resonance Stabilization

Resonance stabilization is crucial in understanding the strength of an acid. When an acid donates a proton, it forms a conjugate base. If this base is stabilized by resonance, it will distribute the negative charge over multiple atoms, decreasing the energy of the system. This is seen in benzoic acid (C₆H₅COOH), where the negative charge on the benzoate ion is delocalized over the ring framework, making it more stable and hence a stronger acid compared to phenol (C₆H₅OH), which lacks such extensive charge distribution.

Conjugate Bases

A conjugate base, which is what remains of an acid after it has donated its proton, also serves as a good predictor of acid strength. A stable conjugate base translates to a stronger acid. The stability depends on factors like charge distribution, size of the atom that carries the charge, and the atom's ability to spread this charge as seen in resonance. This principle helps us explain why oxalic acid (H₂C₂O₄) is stronger than its one-proton-lost form, hydrogen oxalate ion (HC₂O₄⁻), since the additional negative charge in the latter makes it less favorable to lose another proton.

Periodic Trends in Acidity

The periodic table showcases trends that can be used to predict acid strength. Moving across a period, electronegativity increases, often leading to stronger acids. Conversely, as you descend a group, acidity typically increases due to the larger size of the atoms, which better accommodate the negative charge on the conjugate base.

This trend explains why phosphoric acid is more acidic than arsenic acid; despite being in the same group, phosphorus is higher up and thus binds protons more tightly. Similarly, perbromic acid (HBrO₃)'s increased acidity over hypobromous acid (HBrO₂) can be attributed to more oxygen atoms aiding in charge delocalization, aligning with periodic trends and electronegativity principles.