Compare Arrhenius, Bronsted-Lowry, and Lewis definitions of acids and bases.

• A. Arrhenius acids donate H+ in water; Arrhenius bases donate OH-; Bronsted-Lowry acids donate H+ and Bronsted-Lowry bases accept H+; Lewis acids accept electron pairs; Lewis bases donate electron pairs.
• B. Arrhenius acids donate H+ in water; Bronsted-Lowry acids donate H+; Lewis acids donate electron pairs.
• C. Arrhenius acids donate OH-; Bronsted-Lowry acids donate H+; Lewis acids accept electrons.
• D. Arrhenius acids donate H+ in water; Arrhenius bases donate OH-; Bronsted-Lowry acids donate H+; Bronsted-Lowry bases accept H+; Lewis acids accept electron pairs; Lewis bases donate electron pairs.

Answer: (A)

Start from how each definition views what makes an acid or a base. Arrhenius sticks to solutions: acids produce H+ (often as H3O+) in water, and bases produce OH- in water. Bronsted-Lowry widens this to proton transfer: an acid donates a proton, a base accepts a proton, regardless of solvent. Lewis goes even broader: acids accept electron pairs, bases donate electron pairs. The statement captures all three ideas correctly: Arrhenius acids donate H+ in water; Arrhenius bases donate OH-; Bronsted-Lowry acids donate H+; Bronsted-Lowry bases accept H+; Lewis acids accept electron pairs; Lewis bases donate electron pairs. It also meshes with real examples—HCl in water is an Arrhenius and Bronsted-Lowry acid; NH3 can act as a Bronsted-Lowry base and a Lewis base by donating a lone pair; BF3 is a classic Lewis acid by accepting an electron pair. Other options misstate or omit aspects, such as restricting Arrhenius to water in more general contexts or switching the roles of Lewis acids and bases.