Experiment 15: Potentiometric Titration of

Phosphoric Acid

Objective

Solutions of pure phosphoric acid and of phosphoric acid mixed with its conjugate base, dihydrogen

phosphate, will be titrated with a strong base. The data from the titration of the pure phosphoric acid will be

used to determine the concentration of the acid. The titration of the mixture will be done to illustrate the effect

of the conjugate base on the titration curve.

Background

Acids that contain more than one acidic (ionizable) hydrogen (proton) are called polyprotic or polybasic

acids. The dissociation of polyprotic acids occurs in a stepwise fashion, one proton lost at a time. For

example, the generic triprotic acid will dissociate as shown in Reactions (1) through (3). The equilibrium

constants for these reactions are symbolized by Ka. The trailing subscript "a" indicates that the equilibrium

constant describes an acid dissociation reaction. The trailing subscript "n" is written as a number and indicates

which proton is being dissociated.

H3A + H20 H3O+ + H2A

- Ka1 = [H3O+] [H2A

-] / [H3A] (1)

H2A- + H20 H3O

+ + HA2- Ka2 = [ H3O+] [HA2-]/ [H2A

-] (2)

HA2- + H20 H3O+ + A3- Ka3 = [ H3O

+] [A3-] / [HA2-] (3)

Reactions such as those shown in Reactions (4) and (5) are sometimes used as the basis of chemical

calculations, but these reactions are generally not considered to be representative of what is actually occurring

in solution. These reactions can be useful at times even though they are not rigorously correct chemically. The

equilibrium constant for Reaction (4) is the product of Ka1 and Ka2 while the equilibrium constant for

Reaction (5) is the product of Ka1, Ka2, and Ka3. Prove this to yourself.

H3A + 2H20 2H3O+ + HA2- (4)

H3A + 3H20 3H3O+ + A3- (5)

Remember that the numerical value of an equilibrium constant is an indication of how far to the right a reaction

proceeds. The greater the magnitude of the equilibrium constant, the more complete the reaction. Generally

the numerical value of Ka1 is larger than that of Ka2, which is larger than Ka3 This is reasonable when one

considers that successive protons are removed from species with progressively higher negative charges.

Polyprotic acids are called polybasic because they have more than one conjugate base. A conjugate base is

a species that is produced when an acid loses a proton. For example, the species H2A-, HA2-, and A3- are

all conjugate bases of H3A. The completely deprotonated acid, A3-, is a weak base. The other two conjugate

bases are also weak bases, but in addition they have one or more acidic protons. As a result these species

are amphiprotic or amphoteric, which means that they can act as either an acid or a base. Molecules that are

partially deprotonated polyprotic acids are called acid salts. A common acid salt is sodium bicarbonate,

NaHCO3, which is found in baking soda.

When a strong base is added to a solution of a polyprotic acid, the protons of the acid are neutralized in a

stepwise fashion. That is, Reaction (6) will occur first until all of the H3A is used up. Then Reaction (7) will

begin and will continue until all of the H2A- is gone. Then Reaction (8) will occur until all of the HA2- is gone.

This will only be true when the successive dissociation constants are different by a large enough factor and

when all of the acidic species are strong enough. For example, phosphoric acid has Ka values that are

different by a large enough factor to allow it to react with a strong base in a stepwise fashion, however, the

value of Ka3 for phosphoric acid is so small that the last proton of phosphoric acid is extremely difficult to

remove, and the reaction of phosphoric acid that is analogous to Reaction (8) essentially will not occur in

aqueous solution.

H3A + OH- H2A

- + H20 (6)

H2A- + OH- HA2- + H20 (7)

HA2- + OH- A3- + H20 (8)

The titration curve for the titration of a polyprotic acid such as phosphoric acid will have the form shown in

Figure 1. This curve is a plot of pH (-log[H30+]) versus the volume of base added. The first point on the

curve corresponds to a solution of phosphoric acid only. The pH at this point is due solely to the phosphoric

acid in the solution. As soon as some base is added, some H2PO4- is produced. The solution now contains

both phosphoric acid (a weak acid) and its conjugate base dihydrogen phosphate, and thus it is a buffer. A

buffer is a solution that is resistant to change in pH upon dilution and addition of acid and base. This will be

the situation from the initial point to the first equivalence point, and therefore this region of the curve is called

the first buffer zone. The pH of the point midway between the first and second equivalence points is equal to

the negative logarithm of Ka1, pKa1. At the first equivalence point, all of the H3PO4 has been neutralized and

only H2P04- is present in the solution. When more base is added, some of the H2P04

- is neutralized and some

HP042- is produced. This solution contains both dihydrogen phosphate (a weak acid) and its conjugate base

hydrogen phosphate, and thus is a buffer. At all points between the first and second equivalence points, the

solution will contain these two species, and therefore this portion of the curve is called the second buffer

zone. The pH of the solution at the point halfway between the second and third equivalence points is equal to

pKa2. At the second equivalence point the only species in the solution is HPO42-. The region between the

second and third equivalence points is called the third buffer zone because the solution contains both the

weak acid HPO42- and its conjugate base PO4

3-. Halfway between the second and third equivalence points

the pH of the solution is equal to pKa3. At the third equivalence point, only the weak base PO43- is in the

solution. After the third equivalence point, the solution contains phosphate ion and excess unreacted OH-.

Figure 1. Titration curve for titration of phosphoric acid with sodium hydroxide. The point

marked "a" is the initial point. The pH of the solution equals pKa1, pKa2, and pKa3 at points b,

d, and f, respectively. Points c, e, and g are the first, second, and third equivalence points,

respectively. The points between a and c, between c and e, and between e and g are the first,

second, and third buffer zones, respectively.

Changes in the pH during the titrations you will do in this experiment will be monitored with a glass pH

electrode (probe). This electrode will generate a potential in proportion to the hydronium ion concentration in

the solution. The pH meter will electronically convert this potential to pH. Titrations monitored with

electrodes, such as glass electrodes, that measure a potential are called potentiometric titrations.

In this experiment you will titrate a solution containing only phosphoric acid. The curve from this titration will

be used to calculate the concentration of the phosphoric acid and to calculate the Ka1 and Ka2 values of

phosphoric acid. You will also titrate a mixture of phosphoric acid and potassium dihydrogen phosphate to

discover the effect the added species has on the titration curve of phosphoric acid.

More information on Acid-Base Equilibria can be found at http://www.chem.vt.edu/RVGS/ACT/notes/Acid-

Base_Equilibrium.html

Procedure:

1. Titration of H3PO4

Set up the ULI and pH electrode with one of the Macs according to the directions on

http://www.chem.vt.edu/RVGS/ACT/lab/Vernier-pH.html. Calibrate using pH 7 and 4 buffer

solutions. Set the data rate so that a reading is taken every 10 seconds, and set the x-axis for 350

seconds. The y-axis (the pH) should go from 0 to 14.

Next fill a 50.00 mL buret to the 0.00 mL level with standard 1 M NaOH. Then pipet 20 mL of the

phosphoric acid solution into a 100 mL beaker. Place the pH probe in the acid letting the pH value

equilibrate, then begin adding the NaOH in 1 mL increments for each data time while carefully stirring

with the probe. Watch the computer screen as the titration curve develops and continue adding NaOH

until the curve has peaked and leveled off. You do not need to record the final level off the buret

because you have passed the equivalence point by several milliliters.

Save the data so it can be imported into Graphic Analysis or some other graphing program. Then

repeat the procedure two more times using the same amount of acid for each trial. Adjust the x-axis if

necessary. Be sure to also save the data from these two trials.

2. Effect of KH2PO4 on the titration curve:

Add 20 mL of the phosphoric acid to a 100 L beaker and then mix in 20 mL of distilled water. Titrate

as before. Repeat with another 20 mL of phosphoric acid but this time mix in 20 mL of 0.3 M

KH2PO4 solution. Titrate as before.

Analysis:

1. Titration of H3PO4

Import the data from trial 1 into your graphing program. Convert the values on the x-axis from seconds

to volume. Example: the first data point (taken after one second) should represent 0 mL, then the

second data point, after two seconds should represent 1.00 mL of NaOH, after three seconds it

should represent 2.00 mL, and so on.

Place a small "x" at both equivalence points (label as 1 and 2) and an "x" at the location you think

represents the pKa1 and pKa2 values. Calculate the values of Ka1 and Ka2.

Repeat the previous two paragraphs with the second and third trials. Then take an average of the

values of Ka1 and Ka2, and the average deviation for each. Look up the literature values of Ka1 and

Ka2 for phosphoric acid and calculate your error.

Explain why you selected the locations you marked with an "x" as being the equivalence points and as

representing pKa1 and pKa2. What could have been contributing factors in any error you got,

especially if you were off by more than 10%? DO NOT use "human error" as a factor!

2. Effect of KH2PO4 on the titration curve:

Import the data from the two titrations into one graph so you can more easily compare the two curves.

Describe and explain the differences you see in these two titration curves.

Send comments, suggestions, and questions to

Gwen Sibert at the Roanoke Valley Governor's School

gsibert@rvgs.k12.va.us

Prof. Shakhashiri www.scifun.org General Chemistry

PHOSPHORIC ACID, H3PO4About 10 million tons of phosphoric acid, H3PO4, are produced in this country each year. Most of the

acid (about 80%) is used in the production of agricultural fertilizers, with the remainder being used fordetergent additives (about 10%), cleaners, insecticide production, and cattle feed additives. The commercialmethod of preparation is the addition of sulfuric acid to phosphate rock.

3 H2SO4(R) + Ca3(PO4)2(s) + 6 H2O(R) xv 2 H3PO4(s) + 3 CaSO4@2 H2O(s)Pure anhydrous phosphoric acid is a white solid which melts at 42.35EC to form a viscous liquid. In

aqueous solution, phosphoric acid behaves as a triprotic acid, having three ionizable hydrogen atoms. Thehydrogen ions are lost sequentially.

H3PO4(aq) H+(aq) + H2PO4G(aq) Ka1 = 7.5 10G3H2PO4G(aq) H+(aq) + HPO42G(aq) Ka2 = 6.2 10G8HPO4

2G(aq) H+(aq) + PO43G(aq) Ka3 = 1.7 10G12Phosphoric acid is not a particularly strong acid as indicated by its first dissociation constant. It is a strongeracid than acetic acid, but weaker than sulfuric acid and hydrochloric acid. Each successive dissociation stepoccurs with decreasing ease. Thus, the ion H2PO4G is a very weak acid, and HPO4

2G is an extremely weakacid.

Salts of phosphoric acid can be formed by replacing one, two or three of the hydrogen ions. Forexample, NaH2PO4, sodium dihydrogen phosphate, can be formed by reacting one mole of phosphoric acidwith one mole of sodium hydroxide.

H3PO4(aq) + NaOH(aq) xv NaH2PO4(aq) + H2O(R)[net ionic form: H3PO4(aq) + OHG(aq) xv H2PO4G(aq) + H2O(R)]

Similarly, Na2HPO4 (disodium hydrogen phosphate) and Na3PO4, (trisodium phosphate) could be formedby the reaction of one mole of H3PO4 with two and three moles of NaOH, respectively. (Be sure you areable to write net ionic equations for these processes.)

Salts containing the anion H2PO4G are weakly acidic. The tendency of this ion to dissociate is greaterthan its tendency to hydrolyse, that is, its Ka2, is larger than its Kb.

H2PO4G(aq) H+(aq) + HPO42G(aq) Ka2 = 6.2 10G8H2PO4G(aq) + H2O(R) H3PO4(aq) + OHG(aq) Kb = Kw/Ka1 = 1.3 10G12

Because H2PO4G is weakly acidic and of low toxicity, it is used as the acid in some baking powders. Thesebaking powders contain NaH2PO4 and NaHCO3 (sodium bicarbonate). The leavening action of bakingpowders results from the production of carbon dioxide gas by an acid-base reaction between these twoingredients.

H2PO4G(aq) + HCO3G(aq) xv HPO42G(aq) + H2O(R) + CO2(g)In the reaction between them, H2PO4G acts as the Brnsted-Lowry acid, HCO3G as the base. A comparisonof the ionization constants for these two ions reveals that H2PO4G is a stronger acid than HCO3G.

H2PO4G(aq) H+(aq) + HPO42G(aq) Ka2 = 6.2 10G8HCO3G(aq) H+(aq) + CO32G(aq) Ka2 = 4.8 10G11

Salts containing the anion HPO4G are weakly basic. The tendency of this ion to hydrolyse is greater thanits tendency to dissociate.

HPO42G(aq) H+(aq) + PO43G(aq) Ka3 = 1.7 10G12

HPO42G(aq) + H2O(R) H2PO4G(aq) + OHG(aq) Kb = Kw/Ka2 = 1.6 10G7

Solutions containing the phosphate ion, PO43G, are quite basic. This ion has no acidic hydrogen, and its

base ionization constant (hydrolysis constant) is relatively large.PO4

3G(aq) + H2O(aq) HPO42G(aq) + OHG(aq) Kb = 5.9 10G3As a result, solutions of soluble phosphates tend to have the same slippery, soapy feel as solutions of strongbases, such as NaOH or KOH.

Phosphoric acid is used primarily in the manufacture of fertilizers, detergents, and pharmaceuticals. Inthe steel industry, it is used to clean and rust-proof the product. It is also used as a flavoring agent incarbonated beverages (read the ingredients list on a can of Coca-Cola), beer, jams, jellies and cheeses. Infoods, phosphoric acid provides a tart, acidic flavor. A recent study reported in the journal Epidemiology(Vol 18, pp 501506, July 2007), found that drinking two or more cola beverages per day doubled the riskof chronic kidney disease. Cola beverages have been associated with kidney changes that promote kidneystones, which may be a result of the phosphoric acid in colas.

In the manufacture of detergents, phosphoric acid is used to produce water softeners. Water softenersremove Ca2+ and Mg2+ ions from hard water. If not removed, these hard-water ions react with soap andform insoluble deposits that cling to laundry and the washing machine. Phosphates produced fromphosphoric acid are used extensively as water softeners (builders) in detergents. The most widely usedphosphorus compound in solid detergent mixtures is sodium tripolyphosphate, Na5P3O10. As a watersoftener, sodium tripolyphosphate binds to Ca2+ and Mg2+, forming soluble chemical species, calledcomplexes or chelates. These complexes prevent the Ca2+ and Mg2+ from reacting with soap and formingdeposits.

Most phosphoric acid is used in the production of fertilizers. Phosphorus is one of the elements essentialfor plant growth. Organic phosphates are the compounds which provide the energy for most of the chemicalreactions that occur in living cells. Therefore, enriching soils with phosphate fertilizers enhances plantgrowth.

Increasing the phosphate concentration in surface waters also enhances the growth of aquatic plant life.Run-off from fertilized farm lands can stimulate plant growth in lakes and streams. Waste water thatcontains phosphates from detergents can have the same effect. Lakes that are rich in plant nutrients sufferfrom accelerated eutrophication. When the lush aquatic plant growth in a nutrient-rich lake dies, thedecomposition of the dead plant material consumes dissolved oxygen. This consumption reduces the levelof dissolved oxygen to a point where it is insufficient to support animal life. To reduce the threat of lakeeutrophication, many localities have banned the use of phosphates in detergents. In some cases, thephosphates have been replaced by carbonates. In others, new detergents have been developed that do notreact with the Ca2+ and Mg2+ ions of hard water.

Revised: 6 Feb 2008

Experiment 11 POTENTIOMETRIC ANALYSIS OF ACID IN SOFT DRINKS: COLA VS. UNCOLA

2 lab periods Phosphoric acid is a common ingredient in cola drinks; it provides a taste that is both sweet and sour, but does not compete with other flavors. There is some variability in both the amount and composition of the acid in cola drinks. The composition is affected by the equilibrium

H3PO4 + OH- H2PO4- + H2O In this experiment, we will determine the H3PO4 and H2PO4- in a sample of cola drink using a potentiometric titration. We choose a potentiometric detection method over an acid-base indicator for two reasons: (1) the color of the cola obscures indicator changes, and (2) the use of a pH meter permits a more accurate location of the equivalence points in a titration than an indicator. We will also determine the citric acid concentration in an uncola drink. Citric acid, which is also tribasic, is another common ingredient in many soft drinks. The acid dissociation constants for phosphoric and citric acids are as follows: Phosphoric acid Citric acid K1 7.11 x 10-3 7.44 x 10-4

K2 6.32 x 10-8 1.73 x 10-5

K3 7.10 x 10-13 4.02 x 10-7 From these data, it is clear that the constants for phosphoric acid are more than a factor 1,000 apart and three distinct endpoints can therefore be observed in a titration. In contrast, the citric acid constants are closer together and the titration has no definite breaks between the endpoints. Moreover, several of the citric and phosphoric acid endpoints are mutually close. It is therefore advisable to carry out this experiment with soft drinks that do not contain both phosphoric and citric acid. A possible choice is Coca Cola for the phosphoric acid determination, and Squirt for the citric acid determination. If you choose other brands, you should make sure from the ingredient panel that only one or the other of the acids is present. The drinks also should not contain lactic acid or aspartame (Nutrasweet) so do not choose diet drinks!) The potentiometric response of the glass electrode is described by the equation:

Eglass = k - 0.059 pH where k is a constant. Clearly, there is a simple linear relationship between the measured potential and the pH of the solution. For convenience, the pH meter is calibrated in pH units, so that the appropriate values can be read off directly. You will calibrate the meter with pH 4 and 7 (or 10) buffers, following the instructions that are provided with the pH meter. Once the meter is calibrated, the pH of the H3PO4 solution is easily

followed as a function of added NaOH. At pH 10.5 - 11, the glass electrode begins to respond to other ions (mainly Na+ in this case) since so few H3O+ ions remain. This effect, which makes it appear that the pH is lower than it really is, is called the alkaline error. Its occurrence makes it advisable not to carry the titration beyond pH 10.5, meaning that you will not observe the third equivalence point of phosphoric acid: - HPO42- + OH- PO43- + H2O Prelaboratory Assignment The phosphoric acid in a 100.00-mL sample of cola drink was titrated with 0.1025 N NaOH. If the first equivalence point occurred after 13.11 mL of base was added, and the second equivalence point occurred after 28.55 mL of base, calculate the concentrations of H3PO4 and H2PO4- in the cola sample. (Hint: where would the second equivalent point have occurred if only H3PO4 were present?) Apparatus stirrer and (large) stir bar pH meter and glass electrode two 250 mL beakers 50-mL buret 25-mL pipet 1000-mL bottle 25-mL graduated cylinder 1000-mL boiling flask stirring rod 400-mL beaker watch glass Chemicals sodium hydroxide KHP cola unknown (e.g.,Coca Cola) uncola unknown (e.g.,Squirt) pH 4 and 7 buffers phenolphthalein indicator

Procedure 1. Standardize the pH meter with the buffers. See page 40 for procedure. 2. Prepare a standard 0.10 N NaOH solution as described in Experiment 1 (using a KHP

primary standard). For the titration, use your pH electrode in conjunction with the phenolphthalein indicator and compare the electrode response to the indicator color change. Add small increments of titrant, reading both the stabilized pH value and the total volume added after each addition. Initially, the additions should be large enough to cause pH changes of about 0.2 units. When the pH starts to change rapidly, reduce the size of the NaOH aliquots. As you near the equivalence point, the pH will change considerably upon the slightest addition of base. To develop the entire titration curve (pH vs. volume of titrant), you need to proceed somewhat beyond the equivalence point. The electrode response is going to be the principal indicator of the endpoint in this experiment, but you should observe that the indicator changes color at the point where the greatest pH change occurs (note this volume). Any difference is called the indicator error. It should be small. Stop the titration at pH 10.5.

3. Add 100.00 mL of cola to a 250-mL beaker and cover it with a clean watch glass. Bring the

solution just to boiling and keep it warm for five minutes. This expels the CO2 which otherwise would interfere with the titration of H3PO4. Cool the solution by placing ~200 mL of cold water in a 400-mL beaker and carefully resting the beaker with the cola in the cold water.

4. Rinse the electrodes. Refill the buret with the NaOH. 5. Place the glass electrode in the beaker. Add the stir bar and enough water to cover the

electrode. Start the stirrer. 6. Proceed with the titration of the cola solution as you did for the NaOH standardization

(except that there is no indicator here). Expect two equivalence points, one near pH 4 and the other near pH 8. Continue to pH 10.5.

7. Repeat steps 3-6 with the uncola. Now only one equivalence point should be found, near pH

6. Calculations 1. Plot pH (ordinate, i.e, y-axis) vs. volume of NaOH (abscissa, i.e., x-axis) for the

standardization and the two unknowns. 2. Construct first-derivative plots for these titrations. This is accomplished by plotting (pH2

pH1)/(V2 V1) vs. (V1 + V2)/2, where V1 and V2 are two successive titration volumes (totals) and pH1 and pH2 are the corresponding pH values. These plots have peaks where the original graphs have inflection points (i.e. the end points of the titrations). Use them to

estimate the equivalence points. 3. Calculate the molarity of the titrant. 4. Use the equivalence point volumes obtained for the cola titration, along with the NaOH

molarity, to calculate the moles of H3PO4 present. Remember that at the first equivalence point one proton has been titrated, while at the second equivalence point, two protons have reacted. If your results show that Veq2 > 2 Veq1, then not only H3PO4 but also H2PO4- was in present in the drink (see Prelab Assignment). Calculate the concentrations of both.

5. Calculate the concentration of citric acid in the uncola.

Questions 1. In the phosphoric acid titration, could Veq2 < 2Veq1? Explain. 2. Assume that you could titrate to the third equivalence point of H3PO4. What would be the

relationship of Veq3 to Veq2 and Veq1? 3. What is the structure of citric acid? 4. How could CO2 interfere with the titration of H3PO4? 5. The glass electrode that you used appears to be a single device, while it is actually two

electrodes. Explain.

Experiment 11 Name ______________________________ Determination of acid in soft drink Purpose Procedure Explain the procedure used to prepare the cola and uncola samples. Calculations A. Titration of NaOH and KHP

Concentration of KHP: ________________ M Volume of NaOH at equivalence point: ________________ mL

Concentration of NaOH: ________________ M Plot pH vs. volume of NaOH for this titration. Use graph paper and label the equivalence

point. B. Data for titration of cola and uncola with NaOH Make a table with the following columns:

Vol. NaOH Added

pH V2 V1 pH2 pH1 (V1 + V2)/2 pH2 pH1 V2 V1

Plot the first-derivative plots using the above data. C. Results Cola: volume of first equivalence point: _________ml; pH: _________ volume of second equivalence point: ________ml; pH: _________

molarity of H3PO4: _________M; molarity of H2PO4-: _________M Uncola: volume of equivalence point: _________ml; pH: _________

molarity of citric acid: _________M

DETERMINATION OF PHOSPHORIC ACID CONTENT IN SOFT DRINKS LAB PH 8 From Chemistry with Calculators, Vernier Software & Technology, 2000 INTRODUCTION Phosphoric acid is one of several weak acids that present in carbonated beverages. It is a component of all cola soft drinks. Phosphoric acid has a much higher concentration than other acids present in a soft drink, so its concentration can be determined by a simple acid-base titration.

In this experiment, you will titrate a sample of a cola soft drink with a sodium hydroxide solution and determine the concentration of phosphoric acid, H3PO4. Hydrogen ions from the first dissociation of phosphoric acid react with hydroxide ions from the NaOH in a one-to-one ratio in the overall reaction:

H3PO4(aq) + OH(aq) H2O(l) + H2PO4(aq) In this experiment, you will use a pH Sensor to monitor pH as you titrate. The region of most rapid pH change will then be used to determine the equivalence point. The volume of NaOH titrant required to reach the equivalence point will be used to determine the molarity of the H3PO4.

PURPOSE The purpose of this experiment is to determine the amount of phosphoric acid,H3PO4, in a variety of soft drinks by titrating each sample with sodium hydroxide, NaOH.

MATERIALS

LabPro interface 50-mL buret TI Graphing Calculator 100-mL graduated cylinder DataMate program 250-mL beaker pH Sensor ring stand various cola soft drinks, decarbonated utility clamp 0.050 M NaOH magnetic stirrer (if available) deionized water stirring bar ( or stirring rod)

SAFETY

Always wear goggles and an apron in the lab. Do not eat or drink in the lab. Sodium hydroxide is caustic. Avoid spilling the solution on your skin or clothing.

Westminster College SIM PH8-1

Determination of Phosphoric Acid Content in Soft Drinks PROCEDURE 1. Use a graduated cylinder to measure out 40 mL of a decarbonated cola soft drink

and 60 mL of distilled water into a 250-mL beaker

2. Place the beaker on a magnetic stirrer and add a stirring bar. If no magnetic stirrer is available, you need to stir with a stirring rod during the titration.

3. Plug the pH Sensor into Channel 1 of the LabPro interface. Use the link cable to connect the TI Graphing Calculator to the interface. Firmly press in the cable ends.

4. Use a utility clamp to suspend the pH Sensor on a ring stand as shown in Figure 1. Position the pH Sensor in the beverage solution and adjust its position so that it is not struck by the stirring bar.

5. Obtain a 50-mL buret and rinse the buret with a few mL of the 0.050 M NaOH solution. Dispose of the rinse solution as directed by your teacher. Use a utility clamp to attach the buret to the ring stand as shown in Figure 1. Fill the buret a little above the 0.00-mL level of the buret with 0.050 M NaOH solution. Drain a small amount of NaOH solution so it fills the buret tip and leaves the NaOH at the 0.00-mL level of the buret. Record the precise concentration of the NaOH solution in your data table.

Figure 1

6. Turn on the calculator and start the DATAMATE program. Press CLEAR to reset the program.

7. Set up the calculator and interface for the pH Sensor. a. Select SETUP from the main screen. b. If CH 1 displays PH, proceed directly to Step 8. If it does not, continue with this

step to set up your sensor manually. c. Press ENTER to select CH 1. d. Select PH from the SELECT SENSOR menu.

Westminster College SIM PH8-2

Determination of Phosphoric Acid Content in Soft Drinks 8. Set up the data-collection mode.

a. To select MODE, press once and press ENTER . b. Select EVENTS WITH ENTRY from the SELECT MODE menu. c. Select OK to return to the main screen.

9. You are now ready to perform the titration. This process goes faster if one person

manipulates and reads the buret while another person operates the calculator and enters volumes. a. Select START to begin data collection. b. Before you have added any NaOH solution, press ENTER and type in 0 as the

buret volume, in mL. Press ENTER to save the first data pair for this experiment. c. Add 0.5 mL of NaOH solution. When the pH stabilizes, press ENTER and enter

the current buret reading. You have now saved the second data pair for the experiment.

d. Continue to add 0.5-mL increments, entering the buret level after each increment. When the pH has leveled off (near pH 10), press STO to end data collection.

10. Examine the data on the displayed graph to find the equivalence pointthat is, the

0.5-mL volume increment that resulted in the largest increase in pH. As you move the cursor right or left on the displayed graph, the volume (X) and pH (Y) values of each data point are displayed below the graph. Go to the region of the graph with the large increase in pH. Find the NaOH volume (in mL) just before this jump. Record this value in the data table. Then record the NaOH volume after the 0.5-mL addition producing the largest pH increase.

11. Print a copy of the graph of pH vs. volume. Then print a copy of the NaOH volume data and the pH data for the titration.

12. Dispose of the beaker contents as directed by your teacher. Rinse the pH Sensor and return it to the storage solution.

Westminster College SIM PH8-3

Determination of Phosphoric Acid Content in Soft Drinks

DATA SHEET Name ________________________ Name ________________________ Period _______ Class ___________ Date ___________ DETERMINATION OF PHOSPHORIC ACID CONTENT IN SOFT DRINKS DATA TABLE

Concentration of NaOH M

NaOH volume added before the largest pH increase mL

NaOH volume added after the largest pH increase mL

Volume of NaOH added at equivalence point

Moles NaOH

molMoles H3PO4

molConcentration of H3PO4

mol/L

Westminster College SIM PH8-4

Determination of Phosphoric Acid Content in Soft Drinks PROCESSING THE DATA

1.Use your printed graph and data table to confirm the volume of NaOH titrant you recorded before and after the largest increase in pH values upon the addition of 0.5 mL of NaOH solution.

2.Determine the volume of NaOH added at the first equivalence point. To do this, add the two NaOH values determined above and divide by two.

3.Calculate the number of moles of NaOH used using the volume of NaOH added at the equivalence point and the concentration of the NaOH.

4.See the equation for the neutralization reaction given in the introduction. Determine the number of moles of H3PO4 reacted.

5. Recall that you measured out 40.0 mL of the beverage for the titration. Calculate the H3PO4 concentration.

Westminster College SIM PH8-5

Potentiometric Analysis of Acid in Coke Phosphoric acid is a common ingredient in cola drinks; it provides a taste that is both sweet and sour, but does not compete with other flavors. There is some variability in both the amount and composition of the acid in cola drinks. The composition is affected by the equilibrium H3PO4 + OH- H2PO4- + H2O In this experiment, we will determine the H3PO4 and H2PO4- in a sample of cola drink using a potentiometric titration. We are choosing a potentiometric detection method over an acid-base indicator for two reasons: (1) the color of the cola obscures indicator changes, and (2) the use of a pH meter permits a more accurate location of the equivalence points in a titration than an indicator. The potentiometric response of the glass electrode is described by the equation: Eglass = k - 0.059 pH where k is a constant. Clearly, there is a simple linear relationship between the measured potential and the pH of the solution. For convenience, the pH meter is calibrated in pH units, so that the appropriate values can be read off directly. You will calibrate the meter with pH 4 and 7 (or 10) buffers, following the instructions that are provided with the pH meter. Once the meter is calibrated, the pH of the H3PO4 solution is easily followed as a function of added NaOH. At pH 10.5 - 11, the glass electrode begins to respond to other ions (mainly Na+ in this case) since so few H3O+ ions remain. This effect, which makes it appear that the pH is lower than it really is, is called the alkaline error. Its occurrence makes it advisable not to carry the titration beyond pH 10.5, meaning that you will not observe the third equivalence point of phosphoric acid: HPO42- + OH- PO43- + H2O Prelaboratory Assignment Predict the shape of the titration curve for phosphoric acid. Show first equivalence point near pH 4 and second near pH 8. Label your curve with the correct reactions at each equivalence point. Now lets do a sample calculation. The phosphoric acid in a 100.00-mL sample of cola drink was titrated with 0.1025 M NaOH. If the first equivalence point occurred after 13.11 mL of base was added, and the second equivalence point occurred after 28.55 mL of base, calculate the concentrations of H3PO4 and H2PO4- in the cola sample. One small problem, the second equivalence point is a combination of H2PO4- originally in cola sample plus H2PO4- generated during the titration of H3PO4. To get the H2PO4- originally in the cola you must be an ICE table.

Apparatus stir bar pH meter and glass electrode pH 4 and 7 buffers two 250 mL beakers 50-mL buret 100-mL pipet 25-mL graduated cylinder 400-mL beaker watch glass Chemicals sodium hydroxide KHP cola unknown (e.g.,Coca Cola) pH 4 and 7 buffers phenolphthalein indicator Procedure 1. Standardize the pH meter with the buffers. 2. Prepare a standard 0.10 N NaOH solution (using a KHP primary standard). For the titration, use your pH electrode in conjunction with the phenolphthalein indicator and compare the electrode response to the indicator color change. Add small increments of titrant, reading both the stabilized pH value and the total volume added after each addition. Initially, the additions should be large enough to cause pH changes of about 0.2 units. When the pH starts to change rapidly, reduce the size of the NaOH aliquots. As you near the equivalence point, the pH will change considerably upon the slightest addition of base. To develop the entire titration curve (pH vs. volume of titrant), you need to proceed somewhat beyond the equivalence point. The electrode response is going to be the principal indicator of the endpoint in this experiment, but you should observe that the indicator changes color at the point where the greatest pH change occurs (note this volume). Any difference is called the indicator error. It should be small. Stop the titration at pH 12. 3. Add 100.00 mL of cola (using the 100 mL volumetric pipet) to a 250-mL beaker and cover it with a clean watch glass. Bring the solution just to boiling and keep it warm for five minutes. This expels the CO2 which otherwise would interfere with the titration of H3PO4. Cool the solution by placing ~200 mL of cold water in a 400-mL beaker and carefully resting the beaker with the cola in the cold water.

4. Rinse the electrodes. Refill the buret with the NaOH. 5. Place the glass electrode in the beaker. Add the stir bar and start stirring. 6. Proceed with the titration of the cola solution as you did for the NaOH standardization (except that there is no indicator here). Expect two equivalence points, one near pH 4 and the other near pH 8. Continue to pH 10.5. Calculations 1. Plot pH (ordinate, i.e, y-axis) vs. volume of NaOH (abscissa, i.e., x-axis) for the standardization and the cola sample. 2. Construct first-derivative plots for these titrations. This is accomplished by plotting (pH2 pH1)/(V2 V1) vs. (V1 + V2)/2, where V1 and V2 are two successive titration volumes (totals) and pH1 and pH2 are the corresponding pH values. Create a data table similar to the one below to perform first-derivative calculations. These plots have peaks where the original graphs have inflection points (i.e. the end points of the titrations). Vol. NaOH Added pH V2 V1 pH2 pH1 (V1 + V2)/2 pH2 pH1 V2 V1 3. Calculate the standardized molarity of the titrant. 4. Use the equivalence point volumes obtained for the cola titration, along with the NaOH molarity, to calculate the moles of H3PO4 present. Remember that at the first equivalence point one proton has been titrated, while at the second equivalence point, two protons have reacted. If your results show that Veq2 > 2 Veq1, then not only H3PO4 but also H2PO4- was present in the drink. Calculate the concentrations of both. Questions 1. Using the titration curve for KHP, estimate its pKa. The pKa is equal to the pH at half the equivalence point volume. 2. Using the titration curve of Coke, estimate the two pKa (s) for phosphoric acid. The second pKa is equal to the pH at 3/2 times the first equivalence point volume. 3. pKa and Ka values can be used to identify an acid. Using Wikipedia, lookup the pKa for KHP and phosphoric acid. Are these pKa values similar to your estimated values?

DETERMINATION OF PHOSPHORIC ACID CONTENT IN SOFT DRINKSLAB PH 8SAFETYDETERMINATION OF PHOSPHORIC ACID CONTENT IN SOFT DRINKS