Titration of Strong and Weak Acids

SUBDOMAIN 204.3 TITRATION OF STRONG AND WEAK ACIDS

SCA4/5 Task Late Nite Lab

By Nienke Adamse 12/26/10 ABSTRACT

This experiment used the titration technique to determine the unknown concentration of a strong acid (HCl (aq)) and the weak (acetic)acid ((aq)). The base (NaOH (aq)) was used in the quantitative neutralization reaction. The titration technique showed that the unknown concentrations of the strong acid as well as the weak acid could be determined. The concentration of HCl (aq) was 0.6 mol/L and that of (aq) was .16 mol/L. The analytical technique of titration, using a quantitative neutralization reaction to determine the concentration of a weak and a strong acid was fairly accurate; comparing the accepted Ka value with the experimental Ka value, showed a 3.5% error.

EXPERIMENT

For this experiment the following tools, equipment and chemicals were used:

Lab tools: pH Meter

Lab glassware : Burette, Erlenmeyer flask

Lab chemicals :

4 Lab background : Titration is one of the most useful techniques for determining the concentration of an unknown solution, and it is often used with acidic and basic solutions. The equivalence point is reached when a stoichiometric number of moles are added to the unknown solution to neutralize the acid or base.

The determination of the equivalence point will be detected by the following two methods:

* an indicator - a change in the color an indicator added to the solution.

* pH meter - a potentiometric titration, with a sudden change in the pH of the solution as it transitions from being acidic to basic or vice versa. Procedure1 for Late Nite (Laboratory Simulations) Lab:

1. Take a clean Erlenmeyer flask from the Glassware shelf tab and place it on the workbench.

2. Open the Chemical Shelf and add 25 mL HCl (of unknown concentration) to the flask.

3. Add 2 drops of phenolphthalein to the flask.

4. Take a burette from the Glassware shelf and place it on the workbench.

5. Fill the burette with 50 mL of 1M NaOH solution. Record this initial volume.

6. Drag the flask to the lower half of the burette so that the two will be connected. (see figure 1)

7. Open the Data window and click on the flask. Click the Pushpin icon in the blue bar of the Data window to lock its display to the contents of the flask.

8. Take a pH meter from the Tools shelf and drop it on the flask. Record the initial pH of the solution.

9. Open the Properties window and click on the burette. Enter "1" in the amount to add, and click the flow button to drip 1 mL of NaOH into the flask. Record the pH of the solution in the flask.

10. Continue to add NaOH in 1 mL increments. Record the pH for each milliliter added.

11. 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.

12. Continue to add 5 more increments of 1 mL and record the pH at each point.

13. Detach the flask from the burette and drag it to the recycling chute.

14. Take a new flask from the Lab ware Shelf, under the Glassware Tab, and place it on the workbench.

15. Add 25 mL of HCl and 2 drops of phenolphthalein to the flask.

16. Refill the burette to 50 mL NaOH.

17. Based on the results of the previous titration, add enough NaOH solution - all at once - to get to 1 mL BEFORE the endpoint. 18. Using the NaOH from the Chemical Shelf, add .0001 mL to the solution until you reach the endpoint. Record the pH and volume until the endpoint is reached and several drops after it is reached.

Procedure 2 Repeat all steps but use (aq) instead of HCl(aq)

It is important follow the exact procedure because we need to find out what the accurate amount of NaOH added to the acids is, in order to find the equivalence points and thus the concentrations of the unknown acids. The first titration is a rough titration to find out in which area we can make the second titration with the smaller increments of added base. If we start with the smaller increments it might take a long time before we see the change of color. Using two titrations narrows down the area of volume added base and thus makes it possible to add smaller increments of NaOH.

Figure 1

OBSERVATION

First titration HCl (1 ml increments of NaOH) ml added NaOH pH Color solution 2 0.32 no 3 0.37 no 4 0.42 no 5 0.48 no 6 0.54 no 7 0.60 no 8 0.67 no 9 0.75 no 10 0.84 no 11 0.95 no 12 1.09 no 13 1.28 no 14 1.59 no 15 7.00 no 16 12.39 pink 17 12.68 pink 18 12.84 pink 19 12.96 pink 20 13.05 pink 21 13.12 pink

When adding 16 ml of NaOH, the solution turned pink and the pH jumped to 7. In order to find a more accurate point at which the solution becomes pink, a second titration is necessary with smaller increments of NaOH added. End point: 16 ml added NaOH , pH of 12.39

Second titration HCl (0.0001ml increments of NaOH)

ml NaOH pH color 15.0001 8.40 no 15.0002 8.70 no 15.0003 8.88 pink

End point: 15.0003 ml added NaOH, pH of 8.88

Half equivalence point: 7.5 ml, pH of about 0.635 First titration (1ml increments of NaOH) ml NaOH pH color 0 2.78 no 1 4.28 no 2 4.76 no 3 5.23 no 4 8.95 pink 5 12.52 no 6 12.81 No 7 12.97 No 8 13.08 No 9 13.17 No

End point: 4 ml added NaOH, pH of 8.95

Second titration (.05 ml increments of NaOH) starting 1ml before the end point. ml NaOH pH color 3.05 5.26 No 3.10 5.29 No 3.15 5.33 No 3.20 5.36 No 3.25 5.39 No 3.30 5.43 No 3.35 5.47 No 3.40 5.51 No 3.45 5.55 No 3.50 5.60 No 3.55 5.65 No 3.60 5.71 No 3.65 5.78 No 3.70 5.85 No 3.75 5.93 No 3.80 6.04 No 3.85 6.17 No 3.90 6.35 No 3.95 6.65 No 4.00 6.95 pink 4.05 11.23 Pink 4.10 11.53 Pink 4.15 11.71 pink

End point: 4.00 ml added NaOH, pH of 6.95 Half equivalence point: 2.05 ml NaOH added, pH of about 4.76 (did not measure exact pH)

ANALYSIS The equivalence point from this graph = 15 ml added NaOH with a pH of 7 The equivalence point of this graph= 4.5 ml NaOH , pH 9.5 The difference between end-point and equivalence point is, if one is trying to get precise, that the equivalence point is the exact point where the acid and the base in the reaction have canceled each other out and when acid and base have the same number of moles (it is almost impossible to know when this point has been reached with the use of a color indicator). The end-point is the perceptual event at which the indicator actually does change, and you decide you are "done", and you actually read the quantity of measuring reagent you have added. Such as when the solution finally changes color and does not change back right away. The end-point may be slightly different from the true equivalence point, due to the indicator you use and the slowness of pH changing with reagent addition In the Acid-Base titration of NaOH and HCl, using a phenolphthalein indicato,r there is an equivalent amount (same number of moles) of HCl and NaOH together in the flask when the pH is near pH=7. Nothing visible happens right then, because the phenolphthalein indicator changes color closer to pH=10. If adding base from the burette to acid in the flask, you might need to add an extra drop of NaOH solution to make the indicator change color. Then the volume you have added would be "wrong" by about 1 drop in 100ml. This means an error of about 1%. When titrating weak acids and bases, it matters a lot more. Because the pH changes more slowly, the error can be bigger than 20%.

When titrating a monoprotic acid with NaOH, the molar ratio for the reaction is 1:1 and all of the following are true at the equivalence point:

(a) Moles of acid in flask = moles of base added from the burette

(b) (molarity of acid) x (volume of acid) = (molarity of the base) x (volume of added base)

(c) Molarity of the acid = (molarity of the base) x (added volume of the base) / (volume of the acid)

A weak acid, such as CH3COOH (acetic acid), exists in equilibrium with water, and only some of the acid donates the protons to form hydronium ions:

- + CH3COOH + H2O <-> CH3COO + H3O

This also explains why the amount of NaOH added to the weak acid to reach equilibrium was about half of the volume of NaOH that needed to be added to the strong acid to reach equilibrium.

The equilibrium expression for acetic acid is:

+ - Ka= [H3O ][CH3COO ] / [CH3COOH] While titrating the acid with NaOH, the reaction is:

- NaOH(aq) + CH3COOH(aq) -> H2O(l) + Na+(aq) + CH3COO (aq)

During the titration there is a point where EXACTLY HALF of the acid has reacted with NaOH. The concentration of the acetate ion is exactly the concentration of the unreacted acid at this point and it is at the HALF EQUIVALENCE point. From the equilibrium expression above, we have:

+ - Ka = [H3O ] (at the half equivalence point, where [CH3COO ] = [CH3COOH])

Taking the log of both sides at this point, we have: pKa = pH, with pKa defined as pKa = -log(Ka).

The acid's dissociation constant, Ka, then, can be found from the pH of the solution at the half equivalence point. On the titration curve, this is the point where the number of moles of added NaOH is equal to half of the number of moles added to reach the equivalence point.

The experimental Ka value of HCl = 1/-logpH (at the half equivalence point) therefore is 1/-log0.64= 0.23

The experimental Ka value of CH3COOH = 1/-logpH (at the half equivalence point)= 1/- log4.76 = 0.0000174

Errors

When calculating the percent error of the experimental Ka values of HCl and [CH3COOH I looked up the accepted values in the chemistry handbook. The percent error between the experimental and accepted values are, according to %error= experimental Ka – acceptedKa/acceptedKa x100%

For the HCl Ka value the percent error therefore becomes: 0.23- 1300000/1300000= 100%

The Ka value for a strong acid is a really big number. It does not matter in an aqueous solution, since everything will be converted into hydronium and chloride ions: HCl (aq) ==> H+ + Cl-

The acid is 100% dissociated, so all of HCl will be converted into H+ and Cl-. If you are dealing with aqueous solutions, talking about Ka values for strong acids is irrelevant.

Any strong acid when dissolved in water is converted into H+, since H+ is the strongest acid that can exist in water. That is called the leveling effect. In essence, using this approach, you can say that the Ka of any strong acid approaches infinity, any number divided by zero is in essence infinity.

For the [CH3COOH Ka value the percent error becomes: 0.0000174-0.000018/0.000018= 0.035x100% = 3.5% error

The use of the interactive “Nite Lab” simulation equipment eliminates any major human errors, especially when the experiment can be done again and again following the exact same procedures.

DISCUSSION AND CONCLUSION

The titration method to find the unknown concentration of a substance is a fairly accurate method. A conjugated acid-base pair consists out of two substances related to each other by the donating and accepting of a single proton. The equilibrium expression for Ka (the acid dissociation constant) in simple form of a conjugated acid-base pair is: (When H represents the acid and A the base:)

HA (aq) ------> H+ (aq) + A- (aq)

When titrating a monoprotic acid with the base NaOH, the precise equivalence point (where the equivalence point of a chemical reaction occurs when the amount of tyrant (NaOH) added is stoichiometrically equal to the amount of analyte (acid) present in the sample: the smallest amount of titrant that is sufficient to fully neutralize or react with the analyte) can be determined by a titration curve, a graph where the change of pH of the solution is plotted versus the number of moles (or milliliters of) base added. When working with a color indicator, the equivalence point is the point where the neutralization is indicated by a change of color.

At the equivalence point the moles of acid = the moles of base added. With equation M1V1= M2V2, the unknown concentration of the acid can be calculated.

The molarity of the HCl in the first experiment therefore was calculated as 1M (given molarity of NaOH) x volume added NaOH (15 ml)/ volume HCl (25 ml)=

0.6 mol/L or 0.6 M

The molarity of the CH3COOH in the second experiment was calculated as 1Mx 4.5 ml/25 ml= 0.18 mol/L or 0.18 M The difference in acid strength does not influence the found concentrations with the titration method, however the differences in found Ka values show that a weak acid has a strong conjugate base which means that not all acid is dissociated in the aqueous solution, which explains that only a small amount of titrant (base) is needed to reach the equivalence point.

REFERENCES

Late Nite Simulation Labs http://66.246.168.132/csp/lab/ViewLabReport pH Titration curves

http://www.chemguide.co.uk/physical/acidbaseeqia/ph curves .html

Zumdahl, S.S and Zumdahl, S.A (2009) Chemistry. Seventh Edition Boston, MA: Houghton Mifflin Company