Theories of acids and bases URGENT HELP (1 Viewer)

[csi]

Hi,

Please help me with these questions:

1. For sulfuric acid, H2SO4 (aq), write the equation for its-
a) ionisation in water, as required by the Bronsted-Lowry theory (first stage only)
b) neutralisation by the hydrogen carbonate ion, HCO3- (aq)

2.
a) Write the equation showing the ionisation of ammonium in water.
b) Identify the two conjugate acid-base pair.

Note: where possible, please add a bit of explanation to the answers because I’m super stuck.

Much appreciated!!! Thanks guys

 

[Pikapizza]

1.
a) As Bronsted-Lowry theory states, acids are proton donors and, bases are proton acceptors (here proton can be thought of as H+ since hydrogen itself only has one proton too). So we know sulfuric acid is an acid and is reacting with water, thus the H2SO4 donates a H+ and the water accepts it, to form hydronium ion:
H2SO4 (aq) + H2O (l) —> HSO4- (aq) + H3O+ (aq)

b) IMO a weird question in the way they worded it, but I think here it’s just the HCO3- acts as the Bronsted-Lowry base (please correct me if I am wrong):
H2SO4 (aq) + HCO3- (aq) —> HSO4- (aq) + H2CO3 (aq)

For these 2 above reactions, also note that the reaction arrow is forwards only (this is because sulfuric acid is a strong acid, meaning it fully dissociates in water (in part a) and brings neutralisation to completion (in part b)


2.
a) Refer to Bronsted-Lowry definition again, but here ammonia is a base so it instead accepts the proton from water (water acts as an acid here. Also if you are wondering why water can donate and accept protons, it is because water is amphiprotic, meaning it can act as BL acid or BL base; also note that water self ionises to a small extent too)
NH3 (aq) + H2O (l) <—> NH4+ (aq) + OH- (aq)

b) The rule of thumb for conjugate acid-base pairs is that they ‘differ’ by one H. So looking at the equation above, the conjugate pairs are : NH3 and NH4+ (they differ by one H like so) and, H2O and OH-

For the ammonia reaction, it is a weak base, hence it is a equilibrium arrow because it does not fully dissociate in water (the strength of dissociation of acids and bases are in the Ka/Kb part of Mod 6)

I apologise in advance if I made any mistakes!!

 

[csi]

1.
a) As Bronsted-Lowry theory states, acids are proton donors and, bases are proton acceptors (here proton can be thought of as H+ since hydrogen itself only has one proton too). So we know sulfuric acid is an acid and is reacting with water, thus the H2SO4 donates a H+ and the water accepts it, to form hydronium ion:
H2SO4 (aq) + H2O (l) —> HSO4- (aq) + H3O+ (aq)

b) IMO a weird question in the way they worded it, but I think here it’s just the HCO3- acts as the Bronsted-Lowry base (please correct me if I am wrong):
H2SO4 (aq) + HCO3- (aq) —> HSO4- (aq) + H2CO3 (aq)

For these 2 above reactions, also note that the reaction arrow is forwards only (this is because sulfuric acid is a strong acid, meaning it fully dissociates in water (in part a) and brings neutralisation to completion (in part b)


2.
a) Refer to Bronsted-Lowry definition again, but here ammonia is a base so it instead accepts the proton from water (water acts as an acid here. Also if you are wondering why water can donate and accept protons, it is because water is amphiprotic, meaning it can act as BL acid or BL base; also note that water self ionises to a small extent too)
NH3 (aq) + H2O (l) <—> NH4+ (aq) + OH- (aq)

b) The rule of thumb for conjugate acid-base pairs is that they ‘differ’ by one H. So looking at the equation above, the conjugate pairs are : NH3 and NH4+ (they differ by one H like so) and, H2O and OH-

For the ammonia reaction, it is a weak base, hence it is a equilibrium arrow because it does not fully dissociate in water (the strength of dissociation of acids and bases are in the Ka/Kb part of Mod 6)

I apologise in advance if I made any mistakes!!

That’s super helpful! Many thanks