Question

1. Identify the conjugate acid/base pairs in each of the following equations:

(a) H₂S + NH₃ ⇔ NH₄⁺ + HS^-

Pair 1: H₂S and

Pair 2: NH₃ and

(b) HSO₄^- + NH₃ ⇔ SO₄^2- + NH₄⁺

Pair 1: HSO₄^- and​

Pair 2: NH₃ and

(c) HBr + CH₃O^- ⇔ Br^- + CH₃OH

Pair 1: HBr and  

Pair 2: CH₃O^- and  

(d) HNO₃ + H₂O → NO₃^- + H₃O⁺

Pair 1: HNO₃ and​

Pair 2: H₂O and

Answer 1

First, let us define Bronsted Lowry acid/base:

Bronsted Lowry acid: any species that will donate H+ (protons) in solution, and makes pH lower (i.e HCl)

Bronsted Lowry base: any species that will accept H+ (protons) in solution, and makes pH higher (NH3 will accept H+ to form NH4+)

Typically, acid/bases are shown in the left (reactants)

when we write the products:

Bronsted Lowery conjugate base = the base formed when the B.L. acid donates its H+ proton ( i.e. HCl -> Cl-

Bronsted Lowery conjugate acid = the acid formed when the B.L. base accept its H+ proton ( i.e. NH4+ has accept H+ proton)

Note that, typically conjugate bases/acids are shown in the right (product) side

So, from your reaction:

(a) H2S + NH3 ⇔ NH4+ + HS-

Pair 1: H2S and HS- (base)

Pair 2: NH3 and NH4+ (acid)

(b) HSO4- + NH3 ⇔ SO42- + NH4+

Pair 1: HSO4- and​ SO4-2

Pair 2: NH3 and NH4+

(c) HBr + CH3O- ⇔ Br- + CH3OH

Pair 1: HBr and Br-

Pair 2: CH3O- and CH3OH

(d) HNO3 + H2O → NO3- + H3O+

Pair 1: HNO3 and​ NO3-

Pair 2: H2O and H3O+