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Given the equation representing a reaction at equilibrium: h2s(aq) ch3nh2(aq) hs(aq) ch3nh3 (aq) according to one acid-base theory, the forward reaction is classified as an acid-base reaction because

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Since in the forward reaction, H2S is giving a proton away (being the acid) and then CH3NH2 accepting that proton given (being the base).

Answer: [tex]H_2S[/tex] loses proton and [tex]CH_3NH_2[/tex] accepts proton

Explanation: According to the Bronsted-Lowry conjugate acid-base theory, an acid is defined as a substance which looses donates protons and thus forming conjugate base and a base is defined as a substance which accepts protons and thus forming conjugate acid.

For the given chemical equation:

[tex]H_2S(aq.)+CH_3NH_2(aq.)\rightarrow HS^-(aq.)+CH_3NH_3^+(aq.)[/tex]

Here, [tex]H_2S[/tex] is loosing a proton, thus it is considered as an acid and after losing a proton, it forms [tex]HS^-[/tex] which is a conjugate base.

And, [tex]CH_3NH_2[/tex] is gaining a proton, thus it is considered as a base and after gaining a proton, it forms [tex]CH_3NH_3^+[/tex] which is a conjugate acid.

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