Weak/Strong Acid w/ Strong Base Titrations and pH Indicator Selection Help

In summary, the end point pH for a weak acid and strong base titration can be calculated using the Henderson-Hasselbalch equation pH = pKa + log [conjugate / weak acid]. For the given example of 0.30 M acetic acid titrated with 0.15 M sodium hydroxide, the end point pH is 4.75, making thymol blue the most appropriate indicator. For the strong acid and strong base titration of 0.30 M hydrochloric acid with 0.15 M sodium hydroxide, the end point pH is 0.82390874, indicating that none of the indicators from the given table would be appropriate. The correct approach to end point
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kirsten_2009
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Homework Statement



Use the information below to answer the following question(s).

methyl orange: red at pH < 3.1: orange at pH 3.1-4.4: yellow-orange above pH 4.4
litmus: red at pH < 4.5: purple at pH 4.5-8.3: blue above pH 8.3
thymol blue: yellow at pH < 8.0: green at pH 8.0-9.6: blue above pH 9.6
trinitrobenzene: colorless at pH < 12: yellow at pH 12.0-1: orange above pH 14.0

Q#1 Weak Acid with Strong Base Titration: Which of the pH indicators from the table would be most appropriate for the titration of 0.30 M acetic acid (Ka = 1.8 × 10-5) with 0.15 M sodium hydroxide?

a.) Litmus
b.) Trinitrobenzene
c.) Methyl Orange
d.) Thymol Blue
e.) None of These

Answer: D Thymol Blue

Q#2 Strong Acid with Strong Base Titration: Which of the pH indicators from the table would be most appropriate for the titration of 0.30 M hydrochloric acid with 0.15 M sodium hydroxide?

a.) Thymol Blue
b.) Methyl Orange
c.) Litmus
d.) Trinitrobenzene
e.) None of These

Answer: A Thymol Blue

I already have the answers as per the information above, I just don't know how to get that answer and I don't know where I'm going wrong. Any help is much appreciated.

Homework Equations



pH = pKa + log [conjugate/weak acid]

The Attempt at a Solution



Q#1 Weak Acid and Strong Base Titration:

HC2H3C2 + OH- ---> C2H3C2- + H2O
Initial 0.30 M HC2H3C2
add 0.15 M OH-
change -015 M HC2H3C2 -0.15 M OH- +0.15 M C2H3C2
Result (0.30-0.15=0.15 M HC2H3C2) ~0 M OH- 0.15 M C2H3C2

pH = pKa + log [conjugate/weak acid]
pH = -log(ka) + log[conjugate/weak acid]
pH = 4.75 + log [0.15/0.15]
pH = 4.75 + 0
pH = 4.75

If the above calculations are correct, I would have picked C "methyl orange" since it seems to somewhat cover the pH range of what I got (4.75) but obviously something is wrong since the answer is NOT methyl orange but rather thymol blue...?

Q#2 Strong Acid with Strong Base Titration

H30+ + OH- -----> 2 H2O
Initial 0.30 M H30+
add 0.15 M OH-
change -0.15 M H30+ -0.15 M OH-
Result (0.30-0.15=0.15 M H30+) ~0 M OH-

pH = -log[H30]
pH = -log[0.15]
pH = 0.82390874

Again, I would have picked Methyl Orange...what am I doing wrong? Help please... :(
 
Last edited:
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FAQ: Weak/Strong Acid w/ Strong Base Titrations and pH Indicator Selection Help

1. What is a titration and how does it work in acid-base reactions?

A titration is a laboratory technique used to determine the concentration of an unknown solution by reacting it with a known solution. In acid-base reactions, a strong base is slowly added to an acid until the equivalence point is reached, where all of the acid has been neutralized. This allows for the determination of the concentration of the acid.

2. What is the difference between a weak acid and a strong acid in titrations?

In titrations, a weak acid requires more base to reach the equivalence point compared to a strong acid. This is because a weak acid does not completely dissociate in water, meaning not all of the acid molecules contribute to the reaction. Strong acids, on the other hand, fully dissociate in water, making them more reactive and requiring less base to neutralize.

3. What is the purpose of using a pH indicator in acid-base titrations?

The purpose of using a pH indicator is to visually determine the endpoint of the titration. pH indicators are substances that change color depending on the pH of the solution. At the endpoint, the indicator color changes, indicating that all of the acid has been neutralized and the equivalence point has been reached.

4. How do I select the appropriate pH indicator for a weak/strong acid with a strong base titration?

The choice of pH indicator depends on the expected pH range of the titration. For a weak acid with a strong base, the expected pH at the equivalence point will be greater than 7, so a base indicator, such as phenolphthalein, would be suitable. For a strong acid with a strong base, the expected pH at the equivalence point will be close to 7, so a universal indicator or a combination of indicators may be used.

5. What are some potential errors in weak/strong acid with strong base titrations?

One potential error is incomplete mixing of the solutions, which can lead to an inaccurate endpoint determination. Another error is inaccuracy in measuring the volumes of the solutions, which can affect the final concentration calculation. Additionally, the use of the wrong pH indicator or an expired indicator can also result in errors. It is important to carefully follow the procedure and use proper techniques to minimize errors in acid-base titrations.

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