Chapter 20. Laboratory: Quantitative Analysis
Laboratory 20.1: Quantitative Analysis of Vitamin C by Acid-Base Titration
Vitamin C is a vital part of the human diet. A deficiency of vitamin C causes a wasting disease called scurvy, which until the eighteenth century plagued sailors and others whose regular diets provided few fresh fruits and vegetables. In fact, the chemical name for vitamin C, ascorbic acid, is derived from scorbutus, the Latin name for scurvy.
In 1753, British Royal Navy surgeon James Lind published a book, A Treatise on the Scurvy, which concluded that scurvy could be prevented by including citrus fruits in the diet, and that active cases of scurvy could be treated almost miraculously simply by having the patient consume citrus fruits or juices. The Royal Navy began adding lime juice to the regular tots of rum provided to all sailors, accidentally creating the world’s first rum punch. Scurvy disappeared, and British sailors became known as Limeys.
Scurvy is now almost unheard of in the developed world, because our regular diets include more than sufficient vitamin C to prevent scurvy. But many people take vitamin supplements that provide much higher dosages of vitamin C than the recommended daily minimum.
In this lab session, we’ll use acid-base titration with standardized 0.1000 M sodium hydroxide titrant to determine the actual amount of ascorbic acid (C6H8O6) present in a nominally 500 mg vitamin C tablet. One mole of ascorbic acid reacts with one mole of sodium hydroxide to yield one mole of sodium ascorbate and one mole of water, according to the following balanced equation:
C6H8O6 + NaOH → C6H7O6Na + H2O
Using phenolphthalein indicator provides a sharp endpoint for the reaction. The solution remains colorless while ascorbic acid is in excess. When sodium hydroxide is even slightly in excess, the pH of the solution rapidly rises to a point above the color change range of phenolphthalein, and the solution quickly turns bright pink.
Procedure
Before we begin the analysis, we need to decide what sample sizes are appropriate for our quantitative tests.
We know that a 500 mg vitamin C tablet nominally contains 500 mg of ascorbic acid, and we expect the stated value to be reasonably close to the actual value. The gram molecular weight of ascorbic acid is 176.14 g/mol, so at nominal value one 500 mg tablet should contain about 0.0028+ moles of ascorbic acid. One mole of ascorbic acid reacts stoichiometrically with one mole of sodium hydroxide, so we’ll require about 0.0028+ moles of sodium hydroxide to reach the equivalence point. Our 0.1000 M solution of sodium hydroxide contains 0.1 mol/L or 0.0001 mol/mL, so we should need about 28+ mL of titrant to neutralize the ascorbic acid present in one 500 mg tablet.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh a clean, dry 125 mL Erlenmeyer flask and record its mass to 0.01 g on line A of Table 20-1.
Add one 500 mg vitamin C tablet to the flask, reweigh the flask, and record the combined mass of the flask and tablet to 0.01 g on line B of Table 20-1. Subtract the empty mass of the flask from the mass with the vitamin C table, and enter the mass of the vitamin C table to 0.01 g on line C of Table 20-1.
Use the graduated cylinder to transfer about 50 mL of distilled water to the 125 mL Erlenmeyer flask and swirl the flask gently to break up the tablet.
Place the 125 mL flask on the hotplate and heat it gently (do not boil) for 5 to 10 minutes with occasional swirling to dissolve the vitamin C from the sample. It’s normal for a small amount of binder and other components to remain undissolved. Set the flask aside and allow it to cool to room temperature.
Rinse the 50 mL burette with a few mL of 0.1 M sodium hydroxide, and allow it to drain into a waste container.
Transfer at least 40 mL of 0.1 M sodium hydroxide to the burette. Drain a couple mL into the waste container and make sure that there are no bubbles in the burette, including the tip. Record the initial volume as accurately as possible, interpolating between the graduation marks on the burette, and record that volume on line D of Table 20-1.
Add a few drops of phenolphthalein indicator solution to the flask and swirl the contents until they are thoroughly mixed.
We calculated that we should need about 28 mL of sodium hydroxide titrant, so begin by running 25 mL or so of titrant into the receiving flask with constant swirling. As you near the endpoint, a pink color will appear in the flask where the titrant is being added, but will disappear quickly as you swirl the flask. When you reach this point, slow the addition rate to a fast drip, and continue swirling.
When you reach the point where the pink color persists for a few seconds, begin adding the titrant dropwise, and swirl the flask to mix the contents thoroughly after each drop. You’ve reached the endpoint when the pink color persists for at least 30 seconds. Record the final volume on the burette on line E of Table 20-1. Subtract the initial volume of titrant from the final volume, and record the volume of titrant used on line F of Table 20-1.
Using the actual molarity of your nominal 0.1 M sodium hydroxide titrant, calculate the number of moles of sodium hydroxide required to reach the endpoint and enter that value on line G of Table 20-1.
Sodium hydroxide reacts stoichiometrically with ascorbic acid at a 1:1 mole ratio, so the number of moles of ascorbic acid present in the sample is identical to the number of moles of sodium hydroxide needed to reach the endpoint. Calculate the mass of ascorbic acid present in the sample by multiplying the number of moles of sodium hydroxide required by the gram molecular mass of ascorbic acid (176.14 g/mol) and enter that value on line H of Table 20-1.
Dr. Mary Chervenak Comments
An introductory biology class that I took had students consume a known amount of Vitamin C and then collect urine for the 24-hour period following consumption. The urine was then analyzed for Vitamin C. It really was astonishing how little Vitamin C was actually absorbed. Easily one of my favorite experiments.
Item | Value |
A. Mass of empty 125 mL Erlenmeyer flask | __________. ____ g |
B. Mass of empty 125 mL flask + vitamin C table | __________. ____ g |
C. Mass of vitamin C table (B – A) | __________. ____ g |
D. Initial volume of sodium hydroxide titrant | ______._____ mL |
E. Final volume of sodium hydroxide titrant | ______._____ mL |
F. Volume of sodium hydroxide titrant used (C – B) | ______._____ mL |
G. Moles of sodium hydroxide required to reach endpoint | ___.________ mol |
H. Mass of ascorbic acid present in the sample | __________. ____ g |
Review Questions
Laboratory 20.2: Quantitative Analysis of Chlorine Bleach by Redox Titration
Chlorine laundry bleach is an aqueous solution of sodium hypochlorite (NaOCl). Chlorine bleach is produced by running an electric current through a solution of sodium chloride. The electric current oxidizes Cl– ions to Cl02 gas and reduces H+ ions to H02 gas, increasing the concentration of OH– ions. In effect, this electrolysis can be thought of as producing an aqueous solution of sodium hydroxide and chlorine, which immediately react to form sodium hypochlorite.
In solution, sodium hypochlorite dissociates into sodium ions and hypochlorite (OCl–) ions. The bleaching action of chlorine bleach is entirely due to the presence of hypochlorite ion, which is a strong oxidizer. The hypochlorite ion reacts with the colored unsaturated organic compounds that make up stains, oxidizing them to colorless saturated organic compounds.
Inexpensive no-name chlorine bleaches usually contain nothing but 5.25% sodium hypochlorite by mass. (That is, 100.00 g of bleach solution contains 5.25 g of sodium hypochlorite and 94.75 g of water.) Name-brand bleaches are more expensive and often contain scents and brighteners. Name-brand bleaches labeled “extra strength,” “ultra,” or something similar contain slightly higher concentrations of sodium hypochlorite, typically about 6%. Because the hypochlorite ion is the only active ingredient in any chlorine bleach, the sodium hypochlorite content determines bleaching effectiveness.
In this lab session, we’ll use redox titration to quantitatively determine hypochlorite in a chlorine bleach sample. We use a common quantitative analysis procedure called back titration, in which the concentration of an unknown sample is determined by reacting it with an excess of one reagent to produce a product. That product is then back titrated with a second reagent of known concentration until the endpoint is reached. Back titration is used for many reasons, including for samples that are insoluble in water or that contain impurities that interfere with forward titration. But back titration is most commonly used when the endpoint is more apparent by titrating back than by titrating forward.
Rather than attempt to determine hypochlorite concentration by forward titration, we’ll first react the hypochlorite sample in an acetic acid solution that contains excess iodide ions, which are colorless in aqueous solution. The hypochlorite ions oxidize the iodide ions, forming elemental iodine, which is brown in aqueous solution. We’ll then back-titrate with a standardized solution of sodium thiosulfate (Na2S2O3), which reduces elemental iodine to iodide ions. Here are the balanced redox equations for these two reactions:
goggles, gloves, and protective clothing | |
balance and weighing papers | |
Erlenmeyer flask, 125 mL | |
volumetric flask, 100 mL | |
graduated cylinder, 10 mL | |
graduated cylinder, 100 mL | |
pipette, Mohr or serological, 2.00 mL | |
burette, 50 mL | |
dropper or Beral pipette | |
funnel | |
ring stand | |
burette clamp | |
chlorine bleach (100 mL) | |
potassium iodide (~2 g per titration) | |
acetic acid, glacial (~5 mL per titration) | |
sodium thiosulfate, 0.1 M (~60 mL per titration) | |
starch indicator (see Substitutions and Modifications) |
OCl–(aq) + 2 CH3COOH0 + 2 I–(aq) → Cl–(aq) + 2 CH3COO– + I02(aq) + H2O0(l)
2 S2O32–(aq) + I02(aq) → 2 I–(aq) + S4O62–(aq)
As we add sodium thiosulfate titrant to the solution of iodine, the brown color of aqueous iodine becomes fainter and fainter. Because there is no abrupt color change, it is very difficult to judge the endpoint of this titration directly. Fortunately, even at very low concentrations, aqueous iodine reacts with a starch solution to produce an intense blue color. By using starch as an indicator, we can judge the endpoint very precisely. We’ll begin the titration without the indicator. When we’ve added enough sodium thiosulfate titrant to reduce the brown coloration to a faint tint, we know the titration is near the endpoint. We’ll then add starch indicator, which turns the solution bright blue, and continue adding sodium thiosulfate titrant dropwise until the blue color disappears.
Dr. Mary Chervenak Comments
Back titration is often used when an appropriate indicator isn’t available. For example, aqueous silver ions react strongly with halides, but indicators don’t respond to the presence of silver ions. Back-titrating the excess silver ions with thiocyanate will enable the determination of the halide ion concentration. Ferric chloride acts as an indicator for the back titration—the ferric ion forms a colorless complex in water, but the ferric thiocyanate complex is bright red.
Caution
Chlorine bleach is corrosive and a strong oxidizer. Glacial acetic acid is corrosive. Wear splash goggles, gloves, and protective clothing.
Procedure
This lab session is in two parts. In Part I, we’ll determine the density of chlorine bleach solution, which we’ll need later to calculate the mass percentage of sodium hypochlorite in the sample. In Part II, we’ll react the chlorine bleach sample with a solution of potassium iodide in acetic acid, and titrate the iodine formed by that reaction with sodium thiosulfate titrant.
Part I: Determine the Density of Chlorine Bleach
In Part I, we’ll determine the mass of a 100.00 mL sample of chlorine bleach, which allows us to calculate its density.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh a clean, dry, empty 100 mL volumetric flask and record its mass to 0.01 g on line A of Table 20-2.
Use the funnel to fill the volumetric flask with the chlorine bleach sample nearly to the index line. Use the dropper or Beral pipette to add the last mL or two of chlorine bleach, bringing the level of the solution up to the index mark on the flask.
Reweigh the filled flask and record its mass to 0.01 g on line B of Table 20-2.
Subtract the empty mass of the flask from the combined mass of the flask and chlorine bleach, and enter the mass of the chlorine bleach on line C of Table 20-2.
Calculate the density of chlorine bleach in g/mL and enter that value on line D of Table 20-2.
Part II: Determine Hypochlorite Concentration
Before we begin the analysis, we need to decide what sample size is appropriate for our quantitative test. We know that 1 g of chlorine bleach nominally contains 0.0525 g of sodium hypochlorite. The gram molecular mass of sodium hypochlorite is 74.44 g/mol, so 1 g of chlorine bleach contains about 0.0007 moles of sodium hypochlorite. In the first redox equation, one mole of sodium hypochlorite reacts with two moles of iodide ion to form one mole of molecular iodine (I2), so the stoichiometric ratio of hypochlorite:iodine is 1:1. In the second redox equation, one mole of molecular iodine reacts with two moles of sodium thiosulfate to produce two moles of iodide ions, so the stoichiometric ratio of iodine: thiosulfate is 1:2. It follows, then, that the stoichiometric ratio for hypochlorite:thiosulfate is 1:2, which means that two moles of thiosulfate ion react stoichiometrically with one mole of molecular iodine.
Our sodium thiosulfate titrant contains 0.1 mol/L or 0.0001 mol/mL. Because two moles of thiosulfate are required per mole of hypochlorite and the chlorine bleach contains about 0.0007 moles of hypochlorite, about 14 mL of titrant should be required to neutralize 1 g of chlorine bleach. As we determined in Part I, the density of chlorine bleach is slightly more than 1 g/mL (my observed value was 1.0841 g/mL), so just over 15 mL of titrant should be required per mL of chlorine bleach. Using a 3 mL sample of bleach would require 45+ mL of titrant, uncomfortably close to the 50 mL capacity of our burette. We decided to use a 2.0 mL sample, which should require just over 30 mL of titrant, well within the capacity of our burette.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh about 2.0 g of potassium iodide and add it to about 25 mL of distilled water in a 125 mL Erlenmeyer flask. (Alternatively, you can use 25 mL of a 0.5 M bench solution of potassium iodide.)
Swirl the contents of the flask until the potassium iodide dissolves and then, with swirling, add about 5.0 mL of glacial (concentrated) acetic acid to the flask.
Use the pipette to transfer 2.00 mL of the chlorine bleach sample to the flask. The solution should immediately turn dark brown as the chlorine bleach oxidizes the iodide ions to free iodine. Record the volume of the sample to 0.01 mL on line E of Table 20-2.
Rinse the 50 mL burette with a few mL of 0.1 M sodium thiosulfate, and allow it to drain into a waste container.
Transfer 50 mL or slightly more of 0.1 M sodium thiosulfate to the burette. Drain a couple mL into the waste container, until the level of the solution is below the zero index mark on the burette, and make sure that there are no bubbles in the burette, including the tip. Record the initial volume as accurately as possible, interpolating between the graduation marks on the burette, and record that volume on line F of Table 20-2.
We calculated that we should need about 30 mL of sodium thiosulfate titrant, so begin by running 25 mL or so of titrant into the receiving flask with constant swirling. As you near the endpoint, the intense brown color of the solution in the flask will begin to fade noticeably. When you reach that point, slow the addition rate to a fast drip, and continue swirling.
When the solution in the flask reaches a pale brown color, stop adding titrant.
Add about 2 mL of the starch indicator to the flask and swirl the flask to mix the contents, which immediately take on an intense blue-black color due to the formation of a starch-iodine complex.
Continue adding sodium thiosulfate titrant dropwise, with swirling. When the color of the solution fades to a slight bluish tint, continue swirling the flask for 30 seconds. (See Figure 20-1.) If the blue tint persists, add one more drop of titrant and swirl the solution. The endpoint occurs when the blue tint disappears.
Record the final volume as accurately as possible, interpolating between the graduation marks on the burette, and record that volume on line G of Table 20-2.
Determine the volume of titrant used by subtracting the initial volume reading from the final volume reading. Record that value on line H of Table 20-2.
Using the actual molarity of your nominal 0.1 M sodium thiosulfate titrant, calculate the number of moles of sodium thiosulfate required to reach the endpoint and enter that value on line I of Table 20-2.
Sodium thiosulfate reacts stoichiometrically with iodine at a 1:2 mole ratio, so the number of moles of iodine (and therefore hypochlorite) present in the sample is half the number of moles of sodium thiosulfate needed to reach the endpoint. Calculate the number of moles of sodium hypochlorite present in the sample by halving the number of moles of thiosulfate required, and enter that value on line J of Table 20-2.
Determine the mass of sodium hypochlorite present in the sample by multiplying the number of moles present by the gram molecular weight of sodium hypochlorite (74.4422 g/mol). Enter your calculated value on line K of Table 20-2.
Using your values for the density of chlorine bleach (line D) and the volume of the chlorine bleach sample (line E), calculate the mass of the chlorine bleach sample and enter the result on line L of Table 20-2.
Calculate the mass percentage of sodium hypochlorite in the chlorine bleach sample, and enter the result on line M of Table 20-2.
Disposal
Return the unused chlorine bleach from Part I to the original bottle. The waste solutions from this lab can be flushed down the drain with plenty of water.
Item | Value |
A. Mass of empty 100 mL volumetric flask | __________. ____ g |
B. Mass of volumetric flask + 100.0 mL of chlorine bleach | __________. ____ g |
C. Mass of chlorine bleach (B – A) | __________. ____ g |
D. Density of chlorine bleach (C/100) | ______._____ g/mL |
E. Volume of chlorine bleach sample | ______._____ mL |
F. Initial volume of sodium thiosulfate titrant | ______._____ mL |
G. Final volume of sodium thiosulfate titrant | ______._____ mL |
H. Volume of sodium thiosulfate titrant used (F – G) | ______._____ mL |
I. Moles of sodium thiosulfate required to reach endpoint | ___.________ mol |
J. Moles of sodium hypochlorite present in the sample (G/2) | ___.________ mol |
K. Mass of sodium hypochlorite present in the sample | ______. ________ g |
L. Mass of sample | ______. ________ g |
M. Mass percentage of sodium hypochlorite | ______.______ % |
Review Questions
Laboratory 20.3: Quantitative Anion Analysis of Seawater
Seawater is a complex solution of dissolved solid salts. Although salinity varies with location, depth, time of year, and other factors, one kilogram of standard seawater as defined for analytical purposes contains about 964.83 g of water and about 35.17 g of dissolved salts. That 35.17 g is made up of about 19.35 g (55.0%) chloride ions, 10.78 g (30.7%) sodium ions, 2.71 g (7.71%) sulfate ions, 1.28 g (3.64%) magnesium ions, 0.41 g (1.17%) calcium ions, 0.40 g (1.13%) potassium ions, and 0.24 g (0.68%) other ions, including strontium, bromide, fluoride, bicarbonate, carbonate, and others.
In this lab, we do a quantitative analysis for chloride and sulfate ions, the only two anions that are present in seawater in sufficient quantity to be realistic targets for quantitative analysis in a home chemistry lab. For illustrative purposes, we’ll use different methods to analyze chloride ions and sulfate ions. We’ll determine chloride ion content by titrimetric (volumetric) analysis, and sulfate ion content by gravimetric (mass) analysis.
As we learned in the preceding chapter, chloride ions can be precipitated by silver ions in the form of insoluble silver chloride. For a qualitative analysis, the simple fact that adding silver nitrate to a solution of chloride ions causes a precipitation is sufficient to establish the presence of those two ions. But a quantitative analysis requires that we add just sufficient silver ions to precipitate all of the chloride ions, and no more.
Unfortunately, it’s almost impossible to determine that endpoint directly. Silver chloride precipitates as a fluffy whitish solid that forms clumps and takes a long time to settle. Visually, it’s impossible to tell if you’ve added insufficient silver nitrate, just the right amount, or too much. Fortunately, there’s an elegant solution to this problem, based on the differential molar solubilities of silver chloride and silver chromate. (See Laboratory 13.3: Determine a Solubility Product Constant for more details.)
We’ll add a small amount of lemon-yellow potassium chromate solution to our seawater sample as an indicator. Although silver chromate is considered insoluble, it is slightly more soluble than silver chloride. As we titrate the seawater sample with standard silver nitrate solution, the less-soluble silver chloride precipitates first. (That precipitate appears yellow rather than white because of the chromate ion present in the solution.) Only after all of the silver chloride has precipitated does the silver nitrate titrant begin to react with the chromate ion, forming insoluble, brick-red silver chromate. The end point is readily visible as an abrupt color change from lemon yellow to a distinct orange that resembles orange juice.
goggles, gloves, and protective clothing | |
balance and weighing papers | |
hotplate | |
beaker, 150 mL | |
Erlenmeyer flask, 125 mL | |
volumetric flask, 100 mL | |
graduated cylinder, 100 mL | |
evaporating dish | |
burette, 50 mL | |
serological or Mohr pipette, 5 mL | |
stirring rod | |
funnel | |
ring stand | |
support ring | |
wire gauze | |
burette clamp | |
filter paper | |
seawater (200 mL) | |
silver nitrate, 0.1000 M (~50 mL) | |
potassium chromate, 0.1 M (a few mL) | |
barium nitrate, 0.1000 M (~40 mL) |
The sulfate ion is present in seawater at a much lower concentration than the chloride ion, but at a level high enough to be easily analyzed quantitatively in a home lab. We’ll analyze the sulfate ion by precipitating it with a slight excess of barium nitrate solution. Barium ions react with sulfate ions to produce a precipitate of insoluble barium sulfate, which we’ll separate by filtration. After determining the mass of the barium sulfate, we can easily calculate the concentration of sulfate ions in the seawater sample.
Caution
Silver nitrate is toxic, corrosive, and stains skin, clothing, and other organic materials. Potassium chromate is extremely toxic and a known carcinogen. Soluble barium salts are toxic. Wear splash goggles, gloves, and protective clothing. Wear a disposable respirator mask if you work with solid potassium chromate or barium nitrate. (An N95 or N100 disposable mask from the hardware store or home center is fine.)
Procedure
Before we begin the analysis, we need to decide what sample sizes are appropriate for our quantitative tests.
We know that a seawater sample should contain roughly 20 g of chloride ions per liter. The atomic mass of chlorine is about 35.45 g/mol, so the molarity of seawater with respect to chloride ion is about 20/35.45 = 0.56+ M. If we use a 5 mL aliquot, that sample should contain about 0.0028 moles of chloride ion. One mole of chloride ion reacts stoichiometrically with one mole of silver nitrate, so we’d need about 0.0028 moles of silver nitrate to react with that 5 mL aliquot of seawater. Our 0.1 M silver nitrate titrant contains 0.1 mol/L or 0.0001 mol/mL, so we should need about 28 mL of titrant, well within the range of our 50 mL burette.
We know that a seawater sample should contain roughly 3 g of sulfate ions per liter. The formula weight of the sulfate ion is about 96.06 g/mol, so the molarity of seawater with respect to sulfate ion is about 3/96.06 = 0.03+ M. If we use another 5 mL aliquot of seawater, that sample should contain about 0.00016 moles of sulfate ion. One mole of sulfate ion reacts stoichiometrically with one mole of barium nitrate, so about 0.00016 moles of barium nitrate would react with that 5 mL aliquot of seawater to produce about 0.00016 moles of barium sulfate. The formula weight of barium sulfate is 233.43 g/mol, so this reaction would produce only about 0.04 g of barium sulfate.
Such a small amount of product is likely to introduce significant errors, so we’ll use a much larger sample. It’s convenient to use the 100 mL volumetric flask to obtain that sample. If we start with 100.00 mL of seawater in the volumetric flask and remove a 5.00 mL aliquot for chloride ion testing, that leaves 95.00 mL of seawater in the flask. If a 5 mL aliquot would yield about 0.04 g of barium sulfate, a 95 mL aliquot will yield about 0.76 g of barium sulfate, an amount that is much easier to weigh accurately. (Of course, if you have plenty of seawater available, you can simply use a 100.0 mL aliquot. Use the 95 mL aliquot only if you’ve made up an artificial seawater sample.)
We can also calculate about how much 0.1 M barium nitrate titrant should be needed. If about 0.00016 moles of barium nitrate are needed to react completely with a 5 mL aliquot of seawater, the amount needed to react with a 95 mL aliquot of seawater is (95 mL/5 mL) · 0.00016 mol = 0.0030 mol. Our 0.1 M barium nitrate titrant contains 0.1 mol/L or 0.0001 mol/mL, so we’ll need about 30 mL to react completely with the 95 mL aliquot of seawater. To ensure that all of the sulfate ions are precipitated, we’ll add a slight excess of barium nitrate.
Part I: Determine Seawater Density and Total Dissolved Solids
In Part I, we’ll determine the mass of a 100.00 mL sample of seawater, which allows us to calculate its density. We’ll then evaporate the water to determine the mass of the dissolved solids. (If you make your own artificial seawater from purchased sea salt, you can skip the evaporation step and simply enter the mass of sea salt you dissolved to make your artificial seawater on line G of Table 20-3.)
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh the clean, dry, empty 100 mL volumetric flask and record its mass to 0.01 g on line A of Table 20-3.
Fill the 100 mL volumetric flask to the index line with seawater. Weigh the filled flask, and record its mass to 0.01 g on line B of Table 20-3.
Subtract the mass of the empty flask from the mass of the filled flask, and record the mass of 100.0 mL of seawater to 0.01 g on line C of Table 20-3.
Calculate the density of your seawater sample in g/mL, and enter that value on line D of Table 20-3.
Weigh a clean, dry 150 mL beaker and record its mass to 0.01 g on line E of Table 20-3.
Transfer the contents of the 100 mL volumetric flask to the weighed beaker. To make sure you’ve done a quantitative transfer, rinse the volumetric flask several times with a few mL of distilled water, and add the rinse water to the beaker.
Place the beaker on the hotplate and boil the solution gently to evaporate the water. Keep a close eye on the beaker. When the liquid is nearly boiled off, use the stirring rod to break up any solid masses, reduce the heat and continue heating until the contents of the beaker appear to be completely dry.
Dr. Mary Chervenak Comments
If you have access to an oven, use it to dry the sample. Although the temperature inside the oven is lower than the temperature of the hotplate, extended drying even at a lower temperature is the most effective way to eliminate all moisture from the sample because the heat is uniform across the sample, instead of being concentrated on the bottom.
Allow the beaker to cool completely and reweigh it. Jot down the mass to 0.01 g and return the beaker to the hotplate. Heat the beaker strongly for several minutes, allow it to cool completely, and reweigh it. If the mass is the same as the mass you jotted down earlier, you can assume the sample is completely dry. If the second reweighing yields a smaller mass, the sample is not yet completely dry. Heat the beaker again and reweigh it until two sequential weighings yield the same mass. Record that mass to 0.01 g on line F of Table 20-3.
Subtract the mass of the empty beaker from the combined mass of the beaker and solid residue to determine the mass of the total dissolved solids. Record that mass to 0.01 g on line G of Table 20-3. Calculate the mass percentage of total dissolved solids and enter that value on line H of Table 20-3.
Part II: Determine Chloride Ion Quantitatively
In Part II, we’ll quantitatively analyze the chloride ion content of a seawater sample by using silver nitrate as a titrant to precipitate chloride as silver chloride. Because the endpoint of this titration is difficult or impossible to determine directly, we’ll use potassium chromate as a titration indicator. A change in color from bright yellow to orange indicates that all of the chloride ion has been precipitated and that silver chromate is beginning to form. At its endpoint, the titrated solution resembles orange juice.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Use the pipette to transfer 5.00 mL of seawater to the 125 mL Erlenmeyer flask. Note the volume of the aliquot as accurately as possible, interpolating between the graduation marks on the pipette. Record the volume on line I of Table 20-3.
Use the graduated cylinder to transfer about 45 mL of distilled water to the 125 mL Erlenmeyer flask and swirl the flask to mix the contents. Place a stirring rod in the flask, and leave it there until the titration is complete.
Add sufficient 0.1 M potassium chromate solution to the flask to give the solution a bright lemon-yellow color. (The exact amount is not critical. The yellow chromate ion serves as a titration indicator, so if in doubt, use more rather than less.)
Rinse the 50 mL burette with a few mL of 0.1 M silver nitrate solution, allow it to drain into a waste container.
Transfer at least 40 mL of 0.1 M silver nitrate solution to the burette. Drain a couple mL into the waste container and make sure there are no bubbles in the burette, including the tip. Record the initial volume as accurately as possible, interpolating between the graduation marks on the burette, and record that volume on line J of Table 20-3.
We calculated that we should need about 28 mL of silver nitrate titrant, so begin by running 20 mL or so of titrant into the receiving flask with constant swirling. As the silver nitrate is added, it reacts with the chloride ions in the seawater, forming a milky precipitate of silver chloride, which is tinted bright yellow by the chromate ions present in the indicator.
Continue adding titrant more slowly, and with constant swirling and stirring. (The stirring helps to break up the curds of silver chloride that otherwise prevent complete mixing of the solutions.) As you approach the endpoint of the titration, you’ll see a orange tint begin to appear in the yellow liquid at the point where the titrant comes into contact with it (Figure 20-2). That signals that the endpoint of the titration is near. Continue adding titrant dropwise with constant swirling and stirring until a noticeable color change occurs, from yellow to a distinct orange.
Continue stirring vigorously for at least 30 seconds to make sure that no unreacted chloride ions remain trapped within the silver chloride curd. As you stir, the contents of the flask will probably return to the original bright yellow color as additional chloride ions are released from the silver chloride curd.
Continue adding titrant dropwise with vigorous stirring until the orange coloration remains for at least 30 seconds before fading. At that point, you’re about one drop short of the end point. Add that drop of titrant with stirring and verify that the orange tint persists for at least a full minute. Record the final volume of silver nitrate on line K of Table 20-3. Subtract the initial volume reading from the final volume reading, and enter the volume of titrant required on line L of Table 20-3.
Using the actual molarity of the nominally 0.1 M silver nitrate titrant and the actual volume used, calculate the number of moles of silver ions required to reach the endpoint. Enter that value on line M of Table 20-3.
One mole of silver ions react with one mole of chloride ions, so the number of moles of chloride ions in the 5.00 mL aliquot is identical to the number of moles of silver ions in the volume of titrant required to reach the endpoint. On that basis, calculate the mass of the chloride ions present in the aliquot, and enter that value on line N of Table 20-3.
Calculate the mass of chloride ions per liter of seawater, and enter that value on line O of Table 20-3.
Dispose of the contents of the flask as noted in the Disposal section. Wash and dry all of the glassware you used. The silver chloride precipitate can be quite tenacious, covering the inside of the flask with a milky white film even after you’ve washed it. To remove this film, rinse the inside of the flask with a 6 M or higher solution of aqueous ammonia solution.
Part III: Determine Sulfate Ion Quantitatively
In Part III, we’ll quantitatively analyze the sulfate ion content of a seawater sample by using barium nitrate as a titrant to precipitate sulfate as barium sulfate. After we add sufficient barium nitrate solution to the aliquot of seawater to ensure that all of sulfate ions present in the sample have been precipitated, we’ll filter that precipitate, rinse it thoroughly with water to remove any traces of soluble salts, dry the precipitate, and determine its mass. With the mass of barium sulfate known, it requires only a simple calculation to determine the mass of sulfate ions present in the seawater sample.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Transfer a 100.0 mL aliquot (or a 95.0 mL aliquot; see the introduction to this section) of seawater from the 100 mL volumetric flask to the 250 mL beaker. Record the volume of the aliquot on line P of Table 20-3.
Rinse the flask two or three times with a few mL of distilled water to make sure you’ve transferred all of the sample to the beaker.
Measure about 40 mL of 0.1 M barium nitrate in the 100 mL graduated cylinder, and pour it slowly, with stirring, into the 250 mL beaker that contains the seawater sample. The solution immediately assumes a white, milky appearance as insoluble barium sulfate is formed by the reaction of the barium nitrate with the sulfate ions present in the seawater sample.
Continue stirring the solution for 30 seconds or so to make sure that the solutions are completely mixed and that the reaction is complete. When you remove the stirring rod from the beaker, use your wash bottle to rinse any solids on the stirring rod into the beaker.
Fold a piece of filter paper to prepare it for use, weigh it, and record its mass to 0.01 g on line Q of Table 20-3.
Insert the folded filter paper in the funnel, place a receiving vessel under the funnel, and pour the contents of the beaker into the funnel. Rinse the beaker thoroughly several times with a few mL of distilled water and transfer the rinse to the filter funnel to make sure that all of the solids have been transferred to the filter paper.
After all of the liquid has passed through the filter paper, rinse the filtrand two or three times with several mL of distilled water to ensure that any soluble salts are eliminated from the filtrand.
Remove the filter paper from the funnel and place it in a drying oven or under a heat lamp to evaporate any remaining liquid. After the filter paper and filtrand are completely dry, reweigh the filter paper and product and record the mass to 0.01 g on line R of Table 20-3. Subtract the initial mass of the filter paper from the combined mass of filter paper and filtrand to determine the mass of barium sulfate, and record that value to 0.01 g on line S of Table 20-3.
The gram molecular mass of barium sulfate, BaSO4, is 233.43 g/mol, of which barium constitutes 137.33 g. Use these values to calculate the mass of sulfate ions present in the aliquot, and enter that value on line T of Table 20-3. Calculate the mass of sulfate ions present per liter of seawater, and enter this value on line U of Table 20-3.
Disposal
The waste solutions from this lab contain small amounts of silver and barium, both of which are toxic heavy metals. Any excess silver or barium ions present in solutions can be precipitated with excess sulfate or carbonate ions (for example, from a solution of sodium sulfate or sodium carbonate) and filtered. Dispose of the solids as hazardous waste. Any remaining solutions can be flushed down the drain with plenty of water.
Item | Value |
A. Mass of empty 100 mL volumetric flask | __________. ____ g |
B. Mass of flask + 100.00 mL seawater | __________. ____ g |
C. Mass of 100.00 mL seawater (B – A) | __________. ____ g |
D. Density of seawater (C/100) | __________. ____ g/mL |
E. Mass of 150 mL beaker | __________. ____ g |
F. Mass of 150 mL beaker plus residue after evaporation | __________. ____ g |
G. Mass of total dissolved solids | __________. ____ g |
H. Mass percentage of dissolved solids ([G/C] · 100) | __________. ____ % |
I. Volume of seawater aliquot for chloride analysis | ______._____ mL |
J. Initial volume of silver nitrate titrant | ______._____ mL |
K. Final volume of silver nitrate titrant | ______._____ mL |
L. Volume of silver nitrate titrant used | ______._____ mL |
M. Moles of silver ion required to reach endpoint | ___.________ mol |
N. Mass of chloride ions in 5.00 mL sample | __________. ____ g |
O. Mass of chloride ions per liter of seawater | _____.______ g/L |
P. Volume of seawater aliquot for sulfate analysis | ______._____ mL |
Q. Mass of filter paper | __________. ____ g |
R. Mass of dry filter paper and filtrand | __________. ____ g |
S. Mass of barium sulfate (R – Q) | __________. ____ g |
T. Mass of sulfate ions present in the aliquot | __________. ____ g |
U. Mass of sulfate ions per liter of seawater | _____.______ g/L |