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# equilibrium constant and pH during titration

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HOBr(aq) ----> H+(aq) + OBr-(aq) Ka = 2.3 x 10-9

1. Hydrobromous acid, HOBr, is a weak acid that disassociates in water, as represented by the equation above.

(a) Calculate the value of [H+] in a HOBr solution that has pH of 4.95.

(b) Write the equilibrium constant expression for the ionization of the HOBr in water, then calculate the concentration of the HOBR(aq) in an HOBr solution that has [H+] equal to 1.8 x 10-5 M.

(c) A solution of Ba(OH)2 is titrated into a solution of HOBr.
(i) Calculate the volume of 0.115 M Ba(OH)2 (aq) needed to reach the equivalence point when titrated into a 65.0 mL sample of 0.146 M HOBr(aq)
(ii) Indicate whether the pH at the equivalence point is less than 7, equal to 7 or greater than 7. Explain.

(d) Calculate the number of moles of NaOBr(s) that would have to be added to 125 mL of 0.160 M HOBr to produce a buffer solution with [H+] = 5.00 x 10-9 M. Assume that volume change is negligible.

(e) HOBr is a weaker acid than HBrO3. Account for this fact in terms of molecular structure.

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#### Solution Preview

ap1
HOBr(aq) ↔ H+(aq) + OBr-(aq) Ka = 2.3 x 10-9

1. Hydrobromous acid, HOBr, is a weak acid that disassociates in water, as represented by the equation above.

(a) Calculate the value of [H+] in a HOBr solution that has pH of 4.95.

(b) Write the equilibrium constant expression for the ionization of the HOBr in water, then calculate the concentration of the HOBR(aq) in an HOBr solution that has [H+] equal to 1.8 x 10-5 M.
From the chemical ...

#### Solution Summary

It caclulates the pH of the solution during the titration process. The solution is detailed and well presented.

\$2.19

## Titration of Strong and Weak Acids. Determination and assessment of equivalence points. Calculation of molarity. Estimation of acid/base equilibrium constant, pKa

Procedure 1
Add 25 mL HCl (of unknown concentration) to the flask. Add 2 drops of phenolphthalein to the flask. Fill a burette with 50 mL of 1M NaOH solution. Record this initial volume (50mL). Record the pH of the solution in the flask. (0.22). Continue to add NaOH in 1 mL increments. Record the pH for each milliliter added.(0.27; 0.32; 0.37; 0.42; 0.48; 0.54; 0.60; 0.67; 0.75; 0.84; 0.95; 1.09; 1.28; 1.59; 7.00; 12.39=purple) The pink color will appear in flask all at once when the endpoint is either reached or crossed. Record the burette volume and pH at which this occurs. (34mL; 12.39)
Continue to add 5 more increments of 1 mL and record the pH at each point. (12.68; 12.84; 12.95; 13.05; 13.12)

Take a new flask and add 25 mL of HCl and 2 drops of phenolphthalein to the flask. Refill the burette to 50 mL NaOH. Based on the results of the previous titration, add enough NaOH solution - all at once - to get to 1 mL BEFORE the endpoint. (15mL/7.00 pH). Add NaOH DROP-WISE into the flask. Record the pH and volume until the endpoint is reached and several drops after it is reached. (at first 0.05mL it turned pink with a pH 11.10; after I added 4 more drops of 0.05 mL each with the following pH: 11.4; 11.57; 11.7; 11.79).

Assignment (Procedure 1)
1. Find the equivalence point from the graph. 15mL
2. At what point did your solution turn pink? 15.05 mL with 11.10pH
3. What is the pH at the equivalence point? (is the equivalence point the same as when it turned pink?... if yes then is 11.10 pH)
4. What is the equivalence point volume of NaOH? 15.05mL
5. Calculate the molarity of the HCl from the volumes of acid and base at the equivalence point and the molarity of the NaOH.
6. What is the pH at the half-equivalence point?

Procedure 2
In this procedure you will titrate a weak acid, acetic acid (CH3COOH), to see the difference in the titration curve of a weak acid as opposed to a strong acid. Take a clean flask and add 5 mL of acetic acid (of unknown concentration), 20 mL water and 2 drops of phenolphthalein to the flask. Refill the burette with the NaOH solution, and record the initial volume. (50mL) Record the initial pH of the flask solution. (2.78 pH) Add NaOH in 1 mL increments and record the pH for each milliliter added.(4.28; 4.76; 5.23; 8.95=turned pink) The pink color will appear all at once as the endpoint is either reached or crossed. Note the burette volume and pH at which this occurs. (46mL) Continue to add 5 more 1 mL increments of NaOH into the flask. Record the pH after each increment. (12.52; 12.81; 12.97; 13.08; 13.17)

Take a new flask and add 5 mL of acetic acid, 20 mL water, and 2 drops of phenolphthalein to the flask. Refill the burette to 50 mL NaOH. Based on the results of the previous titration, add enough NaOH solution - all at once - to get to 1 mL BEFORE the endpoint. (3mL/5.23pH) Add NaOH DROP-WISE into the flask (in 0.05mL increments). Record the pH and volume until the endpoint is reached and several drops after it is reached.(5.26; 5.29; 5.33; 5.36; 5.39; 5.43; 5.47; 5.51; 5.55; 5.60; 5.65; 5.71; 5.78; 5.85; 5.93; 6.04; 6.17; 6.35; 6.65; 8.95=pink)

Assignment (Procedure 2)

1. At what pH did your solution turn pink? 8.95
2. What is the pH at the equivalence point?
3. What is the equivalence point volume of NaOH?
4. What is the pH at the half equivalence point?
5. Calculate the molarity of the acetic acid from the volumes of acid and base at the equivalence point and the molarity of the NaOH.
6. Calculate the value of Ka of the acetic acid.
7. Look up the accepted value for the Ka of acetic acid on the internet. Calculate the percent error between the experimental and accepted values according to:
% error = |experimental Ka - accepted Ka| / accepted Ka * 100%

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