A2 AQA CHEM4 3.4.3

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  • Created on: 23-04-14 12:45

3.4.3 Acids and Bases

Brǿnsted-Lowry acid-base equilibria in aqueous solution

·         Acidsà proton donors

·         Basesà proton acceptors

·         Acid-base equilibria involve the transfer of protons (i.e. something is acting as the acid and something is acting as the base)

·         E.g. HA  + H2O à H3O+ + A-

o   The water is acting as the B-L base

o   The HA is acting as the B-L acid

o   The H3O+ is the conjugate acid

o   The A- is the conjugate base

·         Water can act as an acid and a base so it is amphoteric

·        Remember H3O+ is what we call an oxonium ion

Definition and determination of pH

·         pH= -log[H+]

·         [H+]= 10-pH

·         When an acid is diluted

o   New [H+] = original [H+] x (volume of acid/ total volume of solution)

·         A strong acid dissociates completely in solution

o   E.g. HCl à H+ + Cl-

·         Monoprotic acid= donates one proton  

o   [acid] = [H+]

·         Diprotic acid= donates two protons

o   [acid] x 2 = [H+]

·         Strong bases dissociate completely in solution

o   E.g. NaOH à Na+ + OH-

o   [base] = [OH-]

The ionic product of water Kw

·         Water is weakly dissociated

o   i.e. H2O <--> H+ + OH-

·         At a constant temperature [H+]=[OH-] and this means water is always neutral at a constant temperature

·         Kw= [H+][OH-]

·         At 25oC, Kw is 1 x 10-14 mol2dm-6

·         We can rearrange Kw to get [H+] as the subject and sub in our [OH-] to find out the [H+] and use this to find the pH of a base

·         Since [OH-]=[H+] we know that Kw= [H+]2 so Kw = [H+] and again we can use this to find pH

·         The dissociation of water is endothermic so the value of Kw increases and temperature increases and since the position of the equilibrium shifts right and the conc of H+ increases, the…

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