Problem 95 - Chemistry

Question

Compare the strengths of the following pairs of acids: (a) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and \(\mathrm{H}_{2} \mathrm{SeO}_{4}\) (b) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) and \(\mathrm{H}_{3} \mathrm{AsO}_{4}\)

Step-by-Step Solution

Answer

\(\mathrm{H}_{2} \mathrm{SO}_{4}\) is stronger than \(\mathrm{H}_{2} \mathrm{SeO}_{4}\), and \(\mathrm{H}_{3} \mathrm{PO}_{4}\) is stronger than \(\mathrm{H}_{3} \mathrm{AsO}_{4}\).

Step 1: Understanding Acid Strength

Acid strength is primarily determined by the stability of its conjugate base. The more stable the conjugate base, the stronger the acid. Factors affecting stability include electronegativity, size of atoms, and resonance within the molecule.

Step 2: Analyze \\(\mathrm{H}_{2} \mathrm{SO}_{4}\) and \\(\mathrm{H}_{2} \mathrm{SeO}_{4}\\)

Sulfuric acid (\(\mathrm{H}_{2} \mathrm{SO}_{4}\)) and Selenic acid (\(\mathrm{H}_{2} \mathrm{SeO}_{4}\)) both form stable conjugate bases through proton donation. However, sulfur is more electronegative and forms stronger π-bonds (double bonds) with oxygen compared to selenium. This results in more stable sulfate ions compared to selenate ions, making \(\mathrm{H}_{2} \mathrm{SO}_{4}\) a stronger acid than \(\mathrm{H}_{2} \mathrm{SeO}_{4}\).

Step 3: Analyze \\(\mathrm{H}_{3} \mathrm{PO}_{4}\) and \\(\mathrm{H}_{3} \mathrm{AsO}_{4}\\)

Phosphoric acid (\(\mathrm{H}_{3} \mathrm{PO}_{4}\)) and Arsenic acid (\(\mathrm{H}_{3} \mathrm{AsO}_{4}\)) are similar in structure, but phosphorus forms stronger bonds with oxygen due to higher electronegativity and ability to form double bonds more effectively. This makes phosphate ions more stable than arsenate ions, so \(\mathrm{H}_{3} \mathrm{PO}_{4}\) is a stronger acid compared to \(\mathrm{H}_{3} \mathrm{AsO}_{4}\).

Step 4: Conclusion

Comparing each pair, \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is stronger than \(\mathrm{H}_{2} \mathrm{SeO}_{4}\) and \(\mathrm{H}_{3} \mathrm{PO}_{4}\) is stronger than \(\mathrm{H}_{3} \mathrm{AsO}_{4}\). The higher electronegativity and bond strength of sulfur and phosphorus are key factors.

Key Concepts

Electronegativity

Electronegativity is a fundamental chemical property that refers to the ability of an atom to attract electrons towards itself within a chemical bond. It plays a vital role in determining the strength of acids. The more electronegative an atom is, the more strongly it will attract electrons. This increased pull on electrons can lead to a more stable conjugate base after the acid donates a proton. For instance, sulfur is more electronegative than selenium. Because of this, sulfur forms a more stable conjugate base when it gives up a proton, which is why sulfuric acid (\(\mathrm{H}_{2} \mathrm{SO}_{4}\)) is a stronger acid than selenic acid (\(\mathrm{H}_{2} \mathrm{SeO}_{4}\)). Understanding how electronegativity influences acid strength can help predict and compare the strength of similar acids.

Conjugate Base Stability

The stability of a conjugate base is crucial for determining the strength of an acid. When an acid donates a proton, it forms a conjugate base. The more stable this base is, the more likely the acid will donate its proton, hence a stronger acid. Several factors contribute to conjugate base stability:

• Electronegativity: As discussed, atoms with higher electronegativity form more stable bases.
• Atom Size: Larger atoms can better distribute the negative charge over a larger volume, enhancing stability.
• Resonance Structures: If the conjugate base can distribute its negative charge across different resonance structures, it gains additional stability.

When comparing phosphoric acid (\(\mathrm{H}_{3} \mathrm{PO}_{4}\)) and arsenic acid (\(\mathrm{H}_{3} \mathrm{AsO}_{4}\)), phosphoric acid is stronger because its conjugate base (phosphate ion) is more stable due to better electronegativity and resonance distribution. This principle further explains why acids with more stable conjugate bases exhibit increased acid strength.

Phosphoric Acid

Phosphoric acid (\(\mathrm{H}_{3} \mathrm{PO}_{4}\)) is a moderately strong triprotic acid, meaning it can donate three protons in acidic solutions. It is commonly used in various industrial and household applications such as fertilisers, food additives, and rust removal products. The strength of phosphoric acid is attributed to the electronegativity and small size of its central atom, phosphorus, which helps form stable bonds with surrounding oxygen atoms. Additionally, phosphoric acid can form multiple resonance structures, allowing it to spread the negative charge of its conjugate base across different positions. This makes its conjugate base more stable once the acid donates a proton, thereby increasing its overall acidic strength when compared to similar acids like arsenic acid (\(\mathrm{H}_{3} \mathrm{AsO}_{4}\)). Understanding the unique properties of phosphoric acid can help when exploring its reactions and applications.

Sulfuric Acid

Sulfuric acid (\(\mathrm{H}_{2} \mathrm{SO}_{4}\)) is a strong, diprotic acid known for its wide applications in industries like fertilisers, battery manufacturing, and chemical synthesis. Its strength as an acid primarily comes from the high electronegativity and strong bonding capability of sulfur. Being more electronegative than selenium, sulfur can better stabilise the negative charge that results when the acid donates a proton.Moreover, sulfur can form effective double bonds with oxygen, contributing to the stability of the sulfate ion (\(\mathrm{SO}_{4}^{2-}\)), its conjugate base. This makes sulfuric acid significantly stronger than selenic acid (\(\mathrm{H}_{2} \mathrm{SeO}_{4}\)). The ability to predict the behavior of sulfuric acid is essential for safe and effective use in its various roles in industry and chemical reactions.