Identify the stronger base in each of the following pairs.
Explain your reasoning.
(d) HS^- or Br^-

Verified step by step guidance

1. Identify the conjugate acids of each base: HS^- is the conjugate base of H2S, and Br^- is the conjugate base of HBr.
2. Recall that the strength of a base is inversely related to the strength of its conjugate acid. A weaker conjugate acid means a stronger base.
3. Compare the acid strengths: H2S is a weaker acid than HBr, as HBr is a strong acid and H2S is a weak acid.
4. Since H2S is weaker than HBr, its conjugate base (HS^-) is stronger than the conjugate base of HBr (Br^-).
5. Conclude that HS^- is the stronger base compared to Br^- because it is derived from the weaker acid, H2S.

Key Concepts

Acid-Base Theory

Acid-base theory explains the behavior of acids and bases in chemical reactions. According to the Brønsted-Lowry theory, an acid is a proton donor, while a base is a proton acceptor. Understanding this theory is crucial for determining the strength of bases, as it helps identify which species can more readily accept protons.

Conjugate Acid-Base Pairs

Conjugate acid-base pairs consist of two species that differ by the presence of a proton. For example, HS⁻ is the conjugate base of H₂S, while Br⁻ is the conjugate base of HBr. The strength of a base can often be inferred from the strength of its conjugate acid; a stronger conjugate acid corresponds to a weaker base.

Electronegativity and Basicity

Electronegativity refers to the tendency of an atom to attract electrons in a bond. In the context of basicity, a less electronegative atom is generally better at donating electron density to accept protons. Comparing HS⁻ and Br⁻, sulfur is less electronegative than bromine, making HS⁻ a stronger base due to its greater ability to share electron density.