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A weak carbonic acid of 0.01m, 20.00ml, ka = 4.6 x 10^ -7 is with 0.005m potassium Hydroxide(KOH) Determine ph; a) after adding 10.00ml b)after adding 20.00ml c) at equivalence point
Accepted Answer
Here's the breakdown of the pH calculations for the titration of a weak carbonic acid (H₂CO₃) with potassium hydroxide (KOH):
1. Initial pH (before any KOH added):
We need to consider the equilibrium of the weak acid dissociation:
H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)
Use the ICE table and the Ka value to calculate the initial [H⁺] and then determine pH.
2. pH after adding 10.00 mL of KOH:
Calculate the moles of H₂CO₃ and KOH added.
The reaction consumes H₂CO₃ and produces HCO₃⁻. Calculate the moles of each remaining.
Calculate the new concentrations of H₂CO₃ and HCO₃⁻ after the volume change.
Use the Henderson-Hasselbalch equation to find the pH.
3. pH after adding 20.00 mL of KOH:
Repeat the steps above, but remember that you've added more KOH. You might be at the half-equivalence point where pH = pKa.
4. pH at the equivalence point:
At the equivalence point, all the H₂CO₃ has reacted with KOH to form HCO₃⁻.
The HCO₃⁻ will now act as a weak base and undergo hydrolysis.
Calculate the concentration of HCO₃⁻ at the equivalence point.
Use the Kb expression for HCO₃⁻ (related to Ka of H₂CO₃) to find [OH⁻]. Then calculate pH.
Important Notes:
Remember to consider the volume changes throughout the titration.
Use the correct equilibrium expressions and Ka/Kb values for the relevant species.
Be careful with the stoichiometry of the reaction.
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