Answered

Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base:

(a) [tex]\( H_2O \)[/tex]

(b) [tex]\( OH^- \)[/tex]

(c) [tex]\( NH_3 \)[/tex]

(d) [tex]\( CN^- \)[/tex]

(e) [tex]\( S^{2-} \)[/tex]

(f) [tex]\( H_2PO_4^- \)[/tex]



Answer :

Sure! To show that each of the given species can act as a Brønsted-Lowry base, we need to demonstrate that they can accept a proton (H⁺). Here's the detailed step-by-step solution for each species:

### (a) Water ([tex]\( \text{H}_2\text{O} \)[/tex])

Water ([tex]\( \text{H}_2\text{O} \)[/tex]) can accept a proton to form the hydronium ion:

[tex]\[ \text{H}_2\text{O} (l) + \text{H}^+ (aq) \rightarrow \text{H}_3\text{O}^+ (aq) \][/tex]

### (b) Hydroxide ion ([tex]\( \text{OH}^- \)[/tex])

The hydroxide ion ([tex]\( \text{OH}^- \)[/tex]) can accept a proton to form water:

[tex]\[ \text{OH}^- (aq) + \text{H}^+ (aq) \rightarrow \text{H}_2\text{O} (l) \][/tex]

### (c) Ammonia ([tex]\( \text{NH}_3 \)[/tex])

Ammonia ([tex]\( \text{NH}_3 \)[/tex]) can accept a proton to form the ammonium ion:

[tex]\[ \text{NH}_3 (aq) + \text{H}^+ (aq) \rightarrow \text{NH}_4^+ (aq) \][/tex]

### (d) Cyanide ion ([tex]\( \text{CN}^- \)[/tex])

The cyanide ion ([tex]\( \text{CN}^- \)[/tex]) can accept a proton to form hydrogen cyanide:

[tex]\[ \text{CN}^- (aq) + \text{H}^+ (aq) \rightarrow \text{HCN} (aq) \][/tex]

### (e) Sulfide ion ([tex]\( \text{S}^{2-} \)[/tex])

The sulfide ion ([tex]\( \text{S}^{2-} \)[/tex]) can accept a proton to form the hydrogen sulfide ion:

[tex]\[ \text{S}^{2-} (aq) + \text{H}^+ (aq) \rightarrow \text{HS}^- (aq) \][/tex]

### (f) Dihydrogen phosphate ion ([tex]\( \text{H}_2\text{PO}_4^- \)[/tex])

The dihydrogen phosphate ion ([tex]\( \text{H}_2\text{PO}_4^- \)[/tex]) can accept a proton to form phosphoric acid:

[tex]\[ \text{H}_2\text{PO}_4^- (aq) + \text{H}^+ (aq) \rightarrow \text{H}_3\text{PO}_4 (aq) \][/tex]

In each of these reactions, the species acts as a Brønsted-Lowry base by accepting a proton ([tex]\(\text{H}^+\)[/tex]).