Sulfamic Acid Titration C12-5-10
Surfamic acid titration c12-5-10 Introduction:
Neutralization Responses Include the Response of an acid and a Bottom to Generate a salt (ionic compound) and Drinking water.
Acid + Base→
In this lab, sulfamic acid (a weak acid which contains one acidic hydrogen) will be used:
H2NSO2OH(aq) + NaOH(aq) → NaOSO2NH2(aq) + H2O(l)
(Net Equation: H+(aq) + OH-(aq) → H2O(l))
Titration is a process of neutralization, it is commonly used to determine the concentration of an acid or base in a solution.
The moles of H+ = moles of OH- at this point (called the equivalence point).
Information about the analyte (i.e. mass) can be calculated at the equivalence point.
The volume of titrant is recorded and the moles of titrant can then be calculated using n = C·V, where n = # of moles, C = concentration in mol/L and V = volume in L.
The end point in a titration is often signaled by the color change of an indicator and occurs just slightly past the equivalence point.
An indicator is a substance (weak acid) that has distinctively different colors in acidic and basic media.
The progress of an acid-base titration is often monitored by plotting the pH of the solution being analyzed as a function of the amount of titrant added (called a titration curve).
Types of Titrations:
1. Strong Acid / Strong Base pH at equivalence point = 7
2. Weak Acid / Strong Base pH at equivalence point >7
3. Strong Acid / Weak Base pH at equivalence point <7
*Note: weak acid / weak base titrations are too complicated and are almost never carried out.
1. To determine the mass of sulfamic acid (H2NSO2OH) in an unknown sample by titrating with a 0.1 mol/L solution of NaOH.
2. To plot a graph of pH as a function of the volume of NaOH added and generate a titration curve.
Standardized NaOH solution (0.1 mol/L)*
Sulfamic acid sample**
Phenolphthalein indicator solution
2-50 mL Burets
100 mL volumetric flask
250 mL Erlenmeyer flask
25 mL graduated cylinder
Digital pH meter
NOTE: Students find it very helpful if the titration technique is demonstrated first.
A great titration simulation experiment can be performed as a demonstration, or as a lab before students perform their own “wet” lab. Credit goes to Tom Greenbowe.
1. Rinse* and fill a buret with standardized 0.1 mol/L NaOH. Open the stopcock briefly to allow any air bubbles to pass through. (Why is this important?) Record the initial volume of NaOH in the buret to the nearest 0.05 mL.
To rinse the buret, add 2-3 mL of the NaOH solution and allow it to run through the buret. Repeat 3 times.
2. Add ~50 mL of distilled water to a 100 mL volumetric flask. Add the sulfamic acid sample to the flask (be sure to get every last crystal!). Swirl the flask to dissolve the sulfamic acid completely. Add distilled water to the 100 mL mark on the volumetric flask, using an eyedropper near the end for accuracy. (Why is accuracy so important here?)
3. Rinse* and fill a second buret with the unknown sulfamic acid solution. Open the stopcock briefly to allow any air bubbles to pass through. (Again, why is this important?)
To rinse the buret, add 2-3 mL of the HCl solution and allow it to run through the buret. Repeat 3 times.
4. Add 10.00 mL of unknown sulfamic acid solution (delivered from the buret) into a 250 mL Erlenmeyer flask. Add 2 drops of phenolphthalein indicator. Why is indicator required? What is its function? Then add 25 mL of distilled water to this solution. What is the purpose of adding water? Would the titration be different if 50 mL of water was added instead?
5. Record the initial colors of both the NaOH and H2NSO2OH solutions. Also, draw diagrams of the flask and buret, showing the molecules present in each solution prior to the titration.
6. Predict what the pH of the unknown H2NSO2OH solution will be. Explain your prediction. Then, using the digital pH meter, record the initial pH of the H2NSO2OH solution in the Erlenmeyer flask.
Note: The pH will be recorded every 2 mL, continuing past the equivalence point until 45-50 mL of NaOH have been added. (However, when a large change on pH is observed, record the pH after every 0.2 mL until it levels off again). Why is it necessary to record the pH past the equivalence point?
7. Gradually dispense some of the NaOH solution drop-by-drop from the buret into the solution in the Erlenmeyer flask. Swirl the flask constantly as the drops are added. Note any color changes observed, and do so constantly as NaOH is added to the H2NSO2OH solution. Continue to record the pH after every 2 mL of NaOH have been added.
8. As the NaOH is being added to the H2NSO2OH solution, what is happening at the molecular level? Draw diagrams of the flask and buret at this point, showing the molecules present in each solution. Also, symbolically write the equation for the chemical reaction that is taking place.
9. As the NaOH is being added, you will notice that the pH is increasing gradually. At the molecular level, describe why there is a change in pH. Also describe the change in pH using a chemical equation.
10. As the equivalence point is approached, a pinkish color will appear and dissipate more slowly as the titration proceeds. Why does this occur? Now add the NaOH drop-by drop until the endpoint of the titration is reached (this is the point at which a very light pink color is obtained after 20 seconds of swirling the flask). Also, begin measuring the pH more often (after every 0.2 mL of NaOH added).
Why is it necessary to swirl for 20 seconds?
11. Record the volume of NaOH required to reach the endpoint of the titration. Predict what the pH of the solution in the Erlenmeyer flask will be and then record the pH at this point. This value is extremely important. Why? Also, is the endpoint the same as the equivalence point? Explain.
12. At the molecular level, illustrate the species present in the flask at the equivalence point. As the titration progressed, which chemical species increased in number and which decreased?
13. Continue to add NaOH to the HCl solution until 45-50 mL of NaOH has been added. Continue to record the pH every 2 mL and also record any observations. Predict whether the final solution will be acidic, basic or neutral. Record the final pH.
14. Draw a molecular representation of the species in the Erlenmeyer flask at this point.
15. Repeat the titration until 3 accurate trials have been completed. The volumes of NaOH required to reach the endpoint should agree within +/- 0.1 ml.
16. Prepare a data table of your results, including the initial, final and total volumes of NaOH required for each titration.
17. Plot the pH as a function of the volume of NaOH added for one good trial.
18. The equivalence point can be found by taking the midpoint of the steep part of the titration curve. What is the pH at the equivalence point for this titration? What does this value tell you about the strength of the acid and the base involved in the titration? Clearly mark and label this point on your graph.
19. Is the unknown H2NSO2OH solution acidic basic or neutral at the equivalence point?
20. Calculate the [H+] at the equivalence point, using [H+] = invlog(-pH).
21. Write a balanced chemical equation for the reaction of H2NSO2OH with NaOH. Also write the net ionic equation and indicate which two species are undergoing a chemical reaction. Which species are spectator ions?
22. Making use of your Info and the Healthy Substance equation from 21 above, Determine the mass of sulfamic acid (H2NSO2OH) in the 10 mL sample Getting titrated. Use Both the Quantity of NaOH from your Preferred Demo, or Typical the 3 Preferred Demos and use this Typical Quantity of NaOH. As a final point, multiply your End result by 10 to get the mass of sulfamic acid in the Authentic 100 mL Alternative (Identical to the mass of sulfamic acid in the Authentic vial).
23. Compare your result with the actual mass of H2NSO2OH revealed by your teacher. What are some possible reasons for the discrepancies? Suggest some ways to reduce the errors.
24. Describe the Principle and Procedure of titration in as A great deal detail as Feasible. Be Positive to use molecular and symbolic representations. Consist of the Conditions titration, neutralization, titrant, Finish Place, equivalence Place, indicators, analyte and titration curve.
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