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

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Identify the conjugate acids of each base: HS^- is the conjugate base of H2S, and OH^- is the conjugate base of H2O.

Recall that the strength of a base is inversely related to the strength of its conjugate acid.

Compare the acid strengths: H2S is a stronger acid than H2O.

Since H2S is a stronger acid, its conjugate base (HS^-) is weaker than the conjugate base of the weaker acid (OH^-).

Conclude that OH^- is the stronger base compared to HS^-.

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Brønsted-Lowry Acids and Bases

The Brønsted-Lowry theory defines acids as proton donors and bases as proton acceptors. In this context, a stronger base is one that can more readily accept protons. Understanding this concept is crucial for comparing the basicity of different species, such as HS⁻ and OH⁻.

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Bronsted-Lowry Acids and Bases Example

Conjugate Acid-Base Pairs

Conjugate acid-base pairs consist of an acid and its corresponding base that differ by a single proton. For example, HS⁻ is the conjugate base of H₂S, while OH⁻ is the conjugate base of H₂O. Analyzing these pairs helps in determining the relative strength of bases based on their ability to form stable conjugate acids.

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Conjugate Acid-Base Pairs

Electronegativity and Basicity

Electronegativity refers to the tendency of an atom to attract electrons. In the context of basicity, a less electronegative atom is generally better at donating electron density to bond with protons. Comparing the electronegativities of sulfur in HS⁻ and oxygen in OH⁻ can provide insight into which species is a stronger base.

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