Suppose you have two solutions to titrate with NaOH:
(a) 10 mL of 0.1 M HCl, and
(b) 10 mL of 0.1 M CH3COOH.
Would you expect either of the solutions to require different volumes of NaOH in
order to reach an end-point? Explain your answer.
Assuming you titrate both solutions with the same NaOH titrant AND that you are able to titrate to the exact equivalence point, the HCl and the CH3COOH will require the same volume of NaOH BECAUSE:
the equivalence point is reached when moles HCl = moles NaOH and when moles CH3COOH = mols NaOH.
moles HCl = M x L = 0.1 M x 0.01 L = 0.001 moles
moles CH3COOH = 0.1 M x 0.01 L = 0.001 moles
Therefore, both titrations will be the same.
A caveat: I note that you use the term "end-point" and not equivalence point. IF THE INDICATOR YOU USE IS SUCH THAT THE END POINT AND THE EQUIVALENCE POINT IS NOT THE SAME YOU WILL SEE A DIFFERENCE.
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To determine if either of the solutions would require different volumes of NaOH to reach an end-point, we need to consider their chemical properties. In this case, we have two solutions to titrate with NaOH: 10 mL of 0.1 M HCl and 10 mL of 0.1 M CH3COOH.
HCl is a strong acid, while CH3COOH (acetic acid) is a weak acid. When a strong acid reacts with a strong base, the neutralization reaction is complete, and an end-point is reached quickly. However, when a weak acid like acetic acid reacts with a strong base, the reaction is not complete initially, and an end-point is reached slowly.
In the case of HCl, it will react quickly and completely with the NaOH, requiring a relatively small volume of NaOH to reach the end-point. On the other hand, CH3COOH will not react as quickly, and the reaction will be slower. As a result, a larger volume of NaOH will be needed to reach the end-point due to the incomplete neutralization reaction.
Therefore, we would expect the solution of 0.1 M HCl to require a different volume of NaOH compared to the solution of 0.1 M CH3COOH in order to reach the end-point.