Your personal safety and the safety of others in the laboratory is your first concern.

The vast majority of injuries in the laboratory are preventable with normal precautions. The following list of laboratory safety rules are to be strictly adhered to. By observing these rules and using common sense you will protect yourself and classmates from many problems.

1. Approved safety goggles are to be worn at all times when you are in the laboratory.

2. No open toed shoes are to be worn in the laboratory.

3. Proper laboratory attire is expected. Full length pants are to be worn in lab. Dresses are strongly discouraged. If you choose to wear a dress you are also required to wear a full length laboratory coat. At all times, a laboratory coat or apron is encouraged.

4. There are to be no food or drinks in the laboratory at any time.

5. Pay close attention to safety notes in your lab manual or on the bulletin board concerning the chemicals you are using.

6. Note the location of all important safety items in your laboratory. You will then be prepared in the event an accident does occur. These items include:

(a) Eye wash stations

(b) Safety showers

(c) Fire extinguishers

(d) First aid kits

7. Follow all the instructions contained in your laboratory manual, contained in handouts, and given by your laboratory instructor.

8. Use hoods whenever using noxious or fuming chemicals.

9. Use pipet bulbs whenever pipetting. Never pipet by mouth!

10. Always maintain a clean and organized work area. This will avoid confusion which can lead to accidents, and it will also save you time in lab. Keep common work areas clean. These include the weighing room, the drying ovens, hoods, reagent benches, and sinks.

11. Always use common sense. If you have any questions or concerns about what you are doing, first consult your laboratory instructor.

Every student is required to maintain a suitable laboratory notebook which is to be a record of all work he or she performs in the laboratory portion of the course. The notebook must be bound. A bound lab notebook is required in which any torn or removed pages will be recorded. Your laboratory notebook is to be a permanent record. Thus, spiral notebooks, looseleaf notebooks, and notebooks with perforated pates are not allowed. You are to organize your lab notebook as follows:

1. All pages in your notebook are to be numbered prior to use of the notebook.

2. The first few pages of the notebook are to be reserved for a table of contents listing the experiment, date(s) the experiment was performed, page number.

3. All writing in your notebook is to be in permanent ink. Any errors are to be marked out with a single line. This will allow you or another person to see what error was made (quite often, the perceived error is actually useful data and needed at a later date). You should never white-out, erase, or completely block out any marks in your laboratory notebook. You are also to never remove pages.

4. All date should be treated with great respect. Enter all raw data directly into your lab notebook. Not on scraps of paper which will later be transposed.

5. Label all data, organize and tabulate data for clarity, and always use the correct number of significant figures.

6. Use the following outline for your laboratory write-ups, which should be contained in your lab manual. The introduction should be written up prior to coming to lab.

1. Introduction - A brief statement of the purpose of the experiment followed by a short statement of the pertinent chemical reactions and measurements you will use. Be brief!

2. Raw Data- Organized Tabulations of all your relevant observations. To facilitate organization you may want to prepare tables which you will need prior to lab. All necessary observations should be included.

3. Report- Indicate any changes or deviations from the procedure as it appears in your lab manual, do not rewrite the procedure in your manual. Discuss any experimental problems you had. Include one sample calculation which shows all the data for arriving at one of your answers. Tabulate all results clearly and label data. Include units, necessary significant figures, and calculated uncertainties in your final result. You may want to very briefly comment on your final result and any sources of error which may have arisen.

Quantitative Analysis

Course Introduction and Pipet Calibration

Quantitative analysis is concerned with the precision and accuracy of chemical measurements. As such, we need to be able to evaluate our instruments. As such, we need to be able to evaluate our measurement tools and the technique which we use. Your first experiment will allow you to practice using the analytical balance and volumetric pipets, and to evaluate the accuracy and precision of your own pipet. You will determine the mass of the water delivered by your pipet and calculate the actual volume of your pipet.

1. Cleaning the pipet.

You will be using your 10 ml and 25 ml pipets for this experiment. Clean pipets are essential to the validity of your measurements. The markings on the volumetric pipet are only valid if the pipet is properly used and the pipet is clean. When water is drawn into and delivered from a clean pipet, the inside surface of the pipet will be coated with a smooth, unbroken film of water. The formation of droplets and/or a broken film is an indication of a dirty pipet. These pipets should be cleaned by using cleaning solution (either detergent or chromerge (chromic acid)). To clean the pipet, draw the cleaning solution into the pipet using a pipet bulb until the pipet is about 1//3 full of the solution. Tilt the pipet to a nearly horizontal position and rotate the pipet until all inside surfaces are coated with the solution. Allow the solution to remain in the pipet for about one to two minutes. Drain the cleaning solution back into the cleaning solution bottle or a waste disposal bottle. Finally rinse the pipet repeatedly (three or more times) with distilled water.

Caution: Cleaning solution is a mixture of sulfuric acid and (typically) chromium trioxide. It is highly corrosive and can cause chemical burns. Do not expose your skin to this solution. In the event you get this acid on your skin, treat it as a strong acid burn and flush with water for 10 to 15 minutes.

2. Calibrating the pipet.

Weigh a clean, dry, empty stoppered erlenmeyer flask to an accuracy of 0.1 mg. Use a flask sufficiently large to hold 100 ml.

Draw a small amount of distilled water into the pipet with a pipet bulb and rinse the inside of the pipet. Repeat at least three times. After the pipet is rinsed, fill the pipet to the graduated mark. Gently wipe the outside of the pipet with a paper wipe. Properly dispense the liquid into your weighing flask. Touch the tip of the pipet to the side of the weighing flask and rotate the tip to remove any drops adhering to the tip. Do not blow the last drop out of the inside of the pipet with the pipet bulb.

Now weigh the flask and the water together and determine the weight of the water by difference. Repeat this process at least three times. You may add the second and subsequent aliquots of water directly to the water remaining in the flask from the previous addition. Record each weight and calculate by difference.

3. Calculation of true volume.

In the determination of the mass of water there is an experimental artifact due to buoyancy effects. When the water is added to the flask, it displaces a volume of air equal to the volume of water. In the initial weighing of the flask the measured weight is due to the flask plus the air contained in the flask. The water measurement gives a value of the flask weight, the water, and the air minus the volume of air displaced by the water. A further complication occurs with two pan balances due to the buoyancy effects of the weights used in the measurement (for a complete discussion see, Journal of Chemical Education, 1984, 61, 51). We will assume the balances were calibrated using stainless steel weights. The calculation proceeds as follows:

The true volume of the pipet can be calculated by:

m_meas. = the measured mass of water

m_air = buoyancy of the H₂O - buoyancy of equivalent mass of standard weights

V_wgt = Approximate volume of stainless steel weights of a mass equivalent to the mass of the H₂O delivered to the pipet.

Suppose you obtain the following data:

mass of flask = 35.5864 g

mass of flask plus water from 10 ml pipet = 45.5527 g

measured mass (m_meas. ) of water delivered from pipet = 9.9663 g

Now, calculate the true mass (m_t) of the water delivered:

Thus, the true mass of the water delivered is given by:

m_t = 9.9663 + (10.0013 x 0.00110 - 1.265 x 0.00110)

= 9.9663 + 0.0110 - 0.0014 = 9.9759 g

this gives:

V_t = 9.9759/0.9965 = 10.011 ml

the experimentally determined volume of the pipet.

Repeat the calculation for all trials, tabulate your results, and report a final value along with the confidence limits.

Quantitative Analysis

Gravimetric Analysis

Determination of Sulfur In a Soluble Sulfate

The total soluble sulfur in a sample (for example, seawater) can be determined by dissolution of the sample and oxidation to the sulfate. By addition of a solution of a soluble salt of barium (in the current experiment barium chloride is employed), the precipitation reaction:

Ba²⁺ + SO₄^2- ---> BaSO₄(s)

is observed. The product barium sulfate is only sparingly soluble (k_sp = 1.3 x 10^-10) and the solid which is produced is amenable to weighing. In this experiment a sample containing sulfur will be analyzed by this reaction, and some of the considerations in obtaining the weighable solid product will be evident.

1. Preparation of barium and sulfate solutions.

The unknown sample must be dried in the oven for at least one hour. Weigh out accurately three portions of about 0.5 g (weighed to the nearest 0.1 mg) and dissolve each sample in a 250-ml beaker with about 200 ml of distilled water containing 1 ml of concentrated HC1. Prepare your precipitating solution by dissolving 5 g of barium chloride dihydrate (BaCl₂^.2H₂O) in 100 ml of distilled water. 20 ml of this solution will be used for each of the three samples.

2. Formation of barium sulfate.

Now heat both the sample solution and the barium chloride solution nearly to boiling. Pour the hot barium chloride solution quickly but carefully into the hot sample solution and stir vigorously. Cover the beaker with a watch glass and allow the precipitate to digest for 1 to 2 hours (overnight is acceptable), keeping the solution hot (80 to 90 C) on a hot plate.

3. Filtration of barium sulfate.

A fine porous porcelain filtering crucible will be used to collect the precipitate. Clean, rinse, and heat to constant weight three crucibles and lids or porous porcelain. These crucibles must be heated within a regular crucible to prevent damage to the porous bottoms. The full heat of a Meker burner is required, and the outside crucible should become orange colored. After heating to constant weight, the crucibles may be cooled and stored in a desiccator.

The solution must be hot at the time of filtration. Decant the clear supernatant solution through the filter. Discard the clear filtrate. Then rinse the precipitate into the filtering crucible with hot distilled water. Remove any precipitate from the walls of the beaker with a rubber policeman and rinse such particles into the filter with the hot water. If the filtrate is cloudy, it must be refiltered. Continue to rinse the precipitate in the filter with hot distilled water until a drop of 1% silver nitrate solution added to a test portion of the washings collected in a test tube shows that chloride is absent.

After washing is complete, dry the filtering crucible and its contents in the oven. Then heat the filtering crucible inside a regular crucible with full heat again. Cool, and determine the weight of the precipitate.

Calculate the percentage of sulfur expressed as sulfur trioxide and as percent sulfate in the unknown sample and report the results.

Quantitative Analysis

Precipitation Titrations

Mohr Method of Chloride Determination

The Mohr titration uses the chromate ion as an indicator for the titration of chloride with silver ion. After the chloride is consumed the excess silver reacts with the chromate to

Ag⁺ + Cl^- ---> AgCl (s) k_sp = 1.8 x 10^-10

2Ag⁺ + CrO₄^2- ---> Ag₂CrO₄ (s) k_sp = 1.1x10^-12

form an orange yellow precipitate. As the indicator and analyte reactions are competitive equilibria, the concentration of the indicator must be carefully chosen. Ideally, a concentration of 6.1x10^-3 M CrO₄^2- will initiate precipitation of Ag₂CrO₄ at the equivalence point of the titration. Experimentally, it has been found that `2x10^-3 M is the optimum concentration.

1. Standardization of silver nitrate solution.

Weigh a clean, dry beaker on a triple beam balance and add 3 to 4 grams of AgNO3 to the beaker. Add distilled water to make 400 ml of solution. Calculate the approximate molarity of your silver nitrate solution.

Caution: Silver nitrate is very corrosive. It is never to be weighed on an analytical balance. It may damage the balance. It can also burn the skin, use caution. (It is also expensive - do not waste it).

Use the analytical balance to weigh out 0.75 - 1.0 g (weighed to the nearest 0.1 mg) of pure, dry NaCl and dissolve it in 250 ml of distilled water in a volumetric flask. Pipet 25.0 ml of this solution into a titration flask and add 4 drops of 0.5 M K₂CrO4. Titrate with the silver nitrate solution until you observe the appearance of an orange-yellow precipitate while vigorously stirring the solution. Vigorous stirring of the solution is necessary to maintain equilibrium throughout the solution. Repeat the titration at least three times. Calculate the concentration of your silver nitrate solution.

2. Determination of chloride in an unknown.

Use the analytical balance to weigh out 1-1.5 g of your unknown (weighed to the nearest 0.1 mg) and dilute in 250 ml in a volumetric flask. Titrate 25 ml aliquots of your unknown chloride solution with the standardized silver nitrate solution. Repeat and calculate the % Cl in your sample.

3. Tap water determination.

Pipet 100 ml of tap water into a titrating flask. Titrate with the standardized silver nitrate solution. You will most likely need to adjust the amount of tap water used in your titration for optimum conditions. Determine the % chloride in your tap water.

Quantitative Analysis

Acid-Base Titrations

Determination of Percent Carbonate

A solution containing CO₃^2- can be analyzed by titration with a strong acid such as HC1. The titration can be stopped either when one mole of HCl per mole of CO₃^2- has been added (the first equivalence point):

H⁺ + CO₃^2- ---> HCO₃^-

or it can be stopped after two moles of HC1 per mole of CO₃^2- have been added (the second equivalence point):

2H⁺ + CO₃^2- ---> H₂CO₃ --(heat)---> CO₂ + H₂O

The first equivalence point can be detected by the one-color indicator, phenolphthalein, going from pink to colorless and the second equivalence point can be detected by the two-color indicator, bromcresol green, going from blue to yellow. One can take advantage of the dissociation of H₂CO₃ into H₂O and CO₂ to obtain a very sharp endpoint with bromcresol green. If dissociation into HC1 is added until the solution is blue-green in color (this occurs when the pH =4.5) and the solution is warmed gently, the CO₂ will be expelled and the buffer capacity of the solution will be lowered. The addition of a fraction of a drop of HCl will cause a sharp drop in the pH due to this loss in buffering capacity. The second endpoint can then be located very accurately.

1.Standardization of the HC1 solution.

An approximately 0.1 M solution of HC1 may be prepared by adding 9 ml of concentrated HCL to about 1 liter if water. Mix the solution thoroughly to ensure uniform concentration.

The HC1 is standardized with the primary standard Na₂CO₃. It is important to dry the Na₂CO₃ overnight for 110 C prior to weighing. After cooling in a desiccator, weigh by difference at least three portions of dried Na₂CO₃ (0.1 to 0.15 g) into separate titration flasks. Add about 4 drops of bromcresol green indicator (0.1%) and titrate with the HC1 solution until the blue color of the indicator becomes blue-green. Warm the solution gently and swirl the flask to expel the CO₂. Cool and continue the titration by the addition of HC1 in fractions of a drop until the solution becomes yellow. Calculate the molarity of the HC1 from the three (or more) replicates. The precision of these measurements should be on the order of two parts per thousand.

2. Determination of the % carbonate in an unknown.

Dry the unknown sample at 110 C for several hours (or overnight). Weigh by difference three portions (or more) (0.3 to 0.5 g) into separate titration flasks. Dissolve three portions in distilled water and titrate with the standard HC1 solution prepared in section 1. Using bromcresol green as the indicator and titrate using the method used for the standardization of the HC1.

3. Calculate and report the % sodium carbonate and in your unknown.

Quantitative Analysis

Non-Aqueous Titrations

The perchloric acid solution will be standardized using pure potassium hydrogen phthalate, which acts as a base in this nonaqueous solvent. The standard acid is then used to titrate an amine, which is too weak a base to be titrated in water. Methyl violet is used as the indictor.

1. Preparation of acid and indictor.

(1) 0.1 M HC1O₄. Add 5 ml of concentrated perchloric acid to 145 ml of glacial acetic acid, mix well, add 15 ml of acetic anhydride, and allow the solution to stand for 30 min. Dilute to 500 ml with glacial acetic acid and allow the solution to cool to room temperature.

(2) Methyl violet indicator. Dissolve 0.2 g of methyl violet in 100 ml of chlorobenzene.

2. Standardization of perchloric acid.

Weigh samples of about 0.5 g (weighed to the nearest 0.1 mg) of dried, pure KHP into three Erlenmeyer flasks and add 60 ml of glacial acetic acid to each flask. Heat the sample cautiously until the sample is in solution. Then cool and add 2 drops of the methyl violet indicator. Titrate with the perchloric acid solution to the disappearance of the violet tinge. Repeat the titration with the other two samples and calculate the molarity of the perchloric acid solution.

3. Analysis of unknown amines.

Weigh three samples of the amine to be determined, taking about 3 mmol for each sample. Dissolve each in about 50 ml of acetic acid. Add 2 drops of methyl violet indicator and titrate with standard perchloric acid to the first appearance of the violet color.

Calculate the number of mmol of amine in the sample.

Quantitative Analysis

Complexometric Titrations

Determination of % Calcium and Water Hardness

EDTA is a polydentate ligand which forms a soluble 1:1 complex with Ca²⁺ and most other metal ions in aqueous solution. The ligand is commonly obtainable as the hydrated disodium salt, Na₂H₂Y^.H₂O which is the form in which it is prepared commercially and which dissolves reasonably readily in water. The reaction of this ligand with the metal ion can be represented as:

An increase in pH will favor the forward reaction and an optimum pH for the titration of any metal ion can be calculated from the formation of the metal - EDTA complex and the four acid dissociation constants of the tetraprotic acid of EDTA. For the reaction of Ca²⁺ or Mg²⁺ to be considered quantitative, the pH of the solution must be greater than or equal to 10. An ammonia-ammonium chloride buffer is used to maintain this pH throughout the titration of Ca²⁺ or Mg²⁺ with EDTA.

The equivalence point of the titration is located by means of a visual indicator. Calmagite is an indicator that may be employed for the titration of Ca²⁺ or Mg²⁺. The indicator itself is a complexing agent and forms pink or red complexes with most metal ions. When the indicator is added to a solution of the metal ion buffered at pH 10, the solution turns pink because the metal ion-indicator complex is formed. When the EDTA solution is added from a buret, all the free metal ion is first complexed by the EDTA. Next, the indicator is displaced by the EDTA from the metal-indicator complex:

Mⁿ⁺ - indicator + EDTA ---> Mⁿ⁺- EDTA + indicator

This occurs because the EDTA binds to the metal much more strongly than does the indicator. When all the indicator has been displaced by the EDTA, the solution assumes the color of the free indicator at pH 10, blue.

The calcium-indicator complex is relatively weak at pH 10. It is necessary, therefore, to add a trace of magnesium to form the pink magnesium to form the pink magnesium-indicator complex. The Mg²⁺ can either be added to the solution that is being titrated, in which case its concentration must be known exactly, or it can be added to the EDTA solutions. Since the EDTA solution is the titrant the amount of Mg²⁺ which has been added is automatically taken into account in the standardization process.

1. Standardization of EDTA solution.

The disodium salt of EDTA is not a primary standard and a solution of EDTA is usually standardized with pure dry CaCO₃.

Dissolve about 6 g of Na₂H₂Y^..2H₂O in about 800 ml of water in a 1 liter bottle. Add 20 ml of a 1% MgCl₂ solution and 3 ml of 6 M NH₃. Add another 150 ml water and make sure that the disodium salt is completely dissolved. Calculate the approximate molarity of your solution.

Dissolve an appropriate amount of pure CaCO₃, (weighed to the nearest 0.1 mg) in a minimal amount of 6 M HC1 and dilute to volume with distilled, deionized water in a 250 ml volumetric flask. (An appropriate amount is that amount of Ca(CO)₃ required to react with 40.0 ml (nearly a full buret) of your previously prepared EDTA solution, remembering that only 25 ml of the 250 ml CaCO₃ solution will be used for each titration).

Pipet 25 ml of the CaCO₃ solution into a 250 ml Erlenmeyer flask. Add 15 ml of NH₃-NH₄Cl buffer at a pH of 10 and a sufficient amount of calmagite indicator to give a pink solution. A very small amount of the indicator will color the solution strongly. Before addition of indicator, you may need to add a masking agent to bind any transition metal ions present. See the instructor.

Titrate the calcium carbonate solution with EDTA until the pink to blue-violet endpoint is reached. Repeat at least three times. Calculate the molarity of your EDTA solution based on these titrations.

2. Determination of % calcium in an unknown

Weigh our approximately 0.7 g (weighed to the nearest 0.1 mg) of unknown an prepare in the same manner as the calcium carbonate standard. Titrate with EDTA using the same method as the standard. Repeat at least three times.

Calculate the % CaO in the sample including the standard deviation and the relative standard deviation.

3. Determination of water hardness

Pipet 100 ml of tap water into a titrating flask, add ml of the NH₃-NH₄Cl buffer and a masking agent followed by the Calmagite indicator. Titrate the solution with EDTA as above. Report the total water hardness as parts per thousand of CaCO₃.

Quantitative Analysis

Redox Titrations

Determination of Iron in an Iron Ore

The basis of a redox titration is the quantitative preparation of the analyte in an appropriate oxidation state followed by titration with a standardized oxidizing or reducing agent. The equivalence point of the titration can be determined by one of two methods. The first method involves the addition of a visual indicator which is a compound which changes structure (and color) when the oxidizing potential of the solution reaches an appropriate level. The second method uses the potential of a known electrochemical half-cell, such as the standard calomel electrode, as a reference with the titration reaction serving as the other half-cell. The cell potential is monitored as a function of added titrant and a titration curve is generated.

The titration reaction used in the current experiment is the oxidation of the analyte Fe²⁺ with potassium dichromate, reaction 1. As an exercise, balance reaction 1.

Fe²⁺ + K₂Cr₂O₇ ---> Fe³⁺ + 2K⁺ + Cr³⁺ (1)

The iron in the ore sample must first be quantitatively be converted to Fe²⁺ prior to carrying out reaction 1. The sample is initially dissolved in acid and is present as Fe³⁺ in solution. The reduction to Fe²⁺ is affected by addition of an excess of SnCl₂. The Sn²⁺ reduces the Fe³⁺ via reaction 2 (balance the reaction).

Fe³⁺ + Sn²⁺ ---> Fe²⁺ + Sn⁴⁺ (2)

It is necessary to remove the excess Sn²⁺ after pre-reduction of the Fe³⁺ since the remaining Sn²⁺ will react with K₂Cr₂O₇. The removal of the excess Sn²⁺ is accomplished by addition of excess HgCl₂, reaction 3. (balance the reaction)

Sn²⁺ + HgCl₂ ---> Hg₂Cl₂ (s) + Sn⁴⁺ (3)

After reaction 3, the remaining unreacted HgCl₂, as well as the Sn⁴⁺ and the Hg₂Cl₂ (s) do not react with K₂Cr₂O₇ and so do not interfere with the main titration reaction. Most importantly, HgCl₂ does not re-oxidize the Fe²⁺.

One final problem with the pre-reduction of iron occurs with the addition of too large an excess of Sn²⁺. If too much Sn²⁺ is added, reaction 4 will occur when the HgCl₂ is added, and an

Sn²⁺ + HgCl₂ ---> Hg⁰ + 2Cl^- + Sn⁴⁺ (4)

interference from Hg⁰ will be present in the main titration reaction. The presence of Hg⁰ is identified by its grey or black color.

An alternative method for pre-reduction of the analyte is a Jones (or a Walden) reductor. This is a packed column of Zn(Hg) amalgam (Jones) or finely divided Ag (Walden). The analyte solution is passed over the column of these strong reducing agents to quantitatively reduce the analyte. This process will not be used in the current experiment.

Now that the iron has been completely reduced to Fe²⁺, the solution is prepared for the titration reaction. Phosphoric acid is added to stabilize the Fe³⁺ complex product and decrease the reduction potential for the corresponding half-reaction.

During the titration it is important that there are no interferences in the solution which will react with the K₂Cr₂O₇. Redox titrations are often performed using KMnO₄. However, the Cl^- in the solution which you are titrating (from the HCl) will oxidize to Cl₂ with KMnO₄. K₂Cr₂O₇ is a milder oxidizing agent and will not react with Cl^-. As an exercise, show why Cl^- reacts with KMnO₄ but not K₂Cr₂O₇. It is possible to titrate Fe²⁺ in the presence of Cl^- with KMnO₄ by the addition of Zimmermann-Reinhardt reagent (a solution of Mn²⁺ in a mixture of concentrated sulfuric and phosphoric acids). The Mn²⁺ binds the Cl^- as a ligand and inhibits its oxidation.

The final consideration is the determination of the equivalence point of the reaction. This can be accomplished using potentiometry or a visual indicator. Though you will only use a visual indicator n your experiment, both methods will be discussed.

Potentiometry involves the measurement of the relative redox potential of an unknown solution. A full discussion of this method may be found in your text. Briefly, upon addition of K₂Cr₂O₇ the Fe²⁺ is quantitatively oxidized to Fe³⁺. Prior to the equivalence point the solution is a mixture of Fe²⁺ and Fe³⁺. This solution then constitutes an electrochemical half-cell when it is placed in an electrical circuit with another half-cell. If the potential of the second half-cell is known, the potential of the Fe²⁺/Fe³⁺ cell may be determined. What equation relates the cell potential back to the concentrations of Fe²⁺ and Fe³⁺? From the cell potential as a function of added titrant, a titration curve may be generated.

The method you will use in your experiment relies on a change in structure of a dye (colored organic molecule) as a function of the redox potential of the solution. The indicator you will use is diphenylamine sulphonate which is oxidized to diphenylbenzidine. The endpoint is signaled by the change to a deep violet color.

K₂Cr₂O₇ is a primary standard and can be used as an oxidizing reagent in redox titrations in the presence of concentrated HCl (< 2M) and in the presence of organic matter. For these reasons, it is a better reagent than KMnO₄.

1. Preparation of standard K₂Cr₂O₇ solution.

Weigh out exactly (to nearest 0.1 mg) between 1.2 and 1.3 g of pure, dry K₂Cr₂O₇. Dissolve this in water to make 250.0 ml of solution.

2. Dissolution of iron ore sample.

Weigh out approximately 0.2 g of ore (weighed exactly, to the nearest 0.1 mg) and transfer to a 500 ml Erlenmeyer flask. Add 10 ml of distilled water and 10 ml of 12 M HCl to each flask and warm gently. Continue heating the solution until the ore is completely dissolved. A small amount of undissolved silica may remain after dissolution. Additional acid may be added to aid in dissolution but do not allow the total volume to exceed 25 ml.

3. Reduction of Fe⁺³

SnCl₂ is used to reduce the Fe⁺³ to Fe⁺². The SnCl₂ solution is prepared by dissolving 110 g of SnCl₂^.2H₂O in 250 ml of 12 M HCl. This solution is then diluted to 1 liter with distilled water. 15 grams of Sn metal is then added to this solution to minimize the air oxidation of Sn⁺². You lab instructor will have prepared this solution and placed it in one of the fume hoods.

The iron ore sample solution is heated to boiling and the SnCl₂ solution is added in a dropwise manner until the yellow color of the FeCl₃ just disappears. After this, add 1-2 drops excess SnCl₂ solution.

Cool the solution to room temperature and slowly add 10 ml of HgCl₂ solution. The HgCl₂ solution is prepared by dissolving 50 g of HgCl₂ in water liter of distilled water. This solution should be available in one of the fume hoods. After the addition of the HgCl₂ solution a milky white precipitate of mercurous chloride should form. If the precipitate is grey (indicating the presence of Hg metal) or no precipitate forms, the amount of SnCl₂ added was incorrect and the sample must be discarded. After the addition of the SnCl₂ the solutions cannot be stored for the next lab period.

4. Preparation of Fe²⁺ solution for titration

After the milky white precipitate forms, let the solution stand for ~2 minutes and then add 200 ml of deionized water, then 5 ml concentrated H₂SO₄, then 10 ml concentrated H₃PO₄, and then 6-8 drops of indicator solution. The indicator is prepared by dissolving 0.3 g barium diphenylamine sulphonate in 100 ml deionized water. Add ~0.5 g Na₂SO₄ and allow the solution to stand. Decant the clear solution for use as the indicator.

5. Titration of Fe²⁺ with K₂Cr₂O₇.

Add the K₂Cr₂O₇ solution from a buret into the Fe²⁺ solution. Stir the solution continuously with a magnetic stirrer. The green color of the solution will turn grey-green just prior to the endpoint which is indicated by the change to a deep violet.

6. Calculate and report the %Fe in your iron ore sample.

Quantitative Analysis

Ion-Exchange Equilibria

Determination of Cation Content

Ion-exchange resins are chemically stable, insoluble solids which may be polymers or hydrous metal oxides which have acidic or basic sites. The acid/base sites may include sulfate, carboxylate, phosphate, amine, or other functionality. Some of the common commercial ion-exchange resins include Amberlite and Dowex. For the case of a sulfate cation-exchange resin, the ion-exchange equilibrium is represented in reaction 1.

Many naturally occurring materials, including clays, act as ion-exchangers. In fact, their ion-exchange properties play crucial roles in the chemistry of natural water systems.

The position of the equilibrium in reaction 1 depends on the concentrations of the various ions in solution in contact with the resin. A concentrated solution of the metal salt MX may be used to displace the protons from the resin. Alternatively, a strong acid may be used to displace the metal from the resin and re-acidify the resin. Thus, the equilibrium in reaction 1 may be used to obtain the concentration of an unknown metal halide salt, MX, by displacement of protons from the ion-exchanger and titration of the displaced protons with NaOH. Conversely, a known solution of MX may be used to determine the ion-exchange capacity or the total number of acidic sites of the resin. In the present experiment, the equilibrium in reaction 1 will be exploited to standardize a base solution, and to determine the number of equivalents of a monovalent cation in an unknown salt.

1. Preparation of carbonate-free NaOH solution.

Prepare 1 liter of CO₂-free distilled water by boiling 1 liter of water, cooling, and storing in an air tight bottle. Next, deliver ~5.3 ml of 50% w/w NaOH solution to the CO₂-free water. Mix the solution thoroughly and store in a polyethylene bottle. This solution should have a concentration of ~0.1 M.

2. Preparation of ion-exchange resin.

Either add 25 ml of resin to a 50 ml buret, or add sufficient resin to a 20 cm by 1.5 cm (diameter) column to fill 75% full. Place a small amount of glass wool at the top of the column to prevent disturbance of the resin upon addition of solutions.

Note: Never let the liquid in the column fall below the surface of the resin in the column.

Acidify the resin by addition of 100 ml of 3.0 M HCl through the column. After acidifying, all of the excess HCl solution must be rinsed from the column. Rinse the column and test the effluent every 20-30 ml by addition of one drop of methyl red indicator and one drop of 0.1 M NaOH. If the effluent is neutral, addition of base should change the solution from red to yellow.

3. Standardization of NaOH solution.

Accurately weigh (to nearest 0.1 mg) 0.3 g of pure, dry KCl and dissolve in 1-2 ml of distilled water. Prepare three solutions in this manner. Quantitatively add one of the three KCl solutions to the column using the smallest amount of distilled water possible. Allow the solution to drain to a level just slightly above the resin surface. This process will displace protons with K⁺ from the resin.

Next, add 3-5 ml of distilled water to the column, allow the column to again drain to just slightly above the resin surface. Collect and save the effluent. Repeat this procedure and continue collecting the effluent so the column drains at the rate of about one drop per second. Periodically, test portions of the effluent for neutrality using the same procedure as above (addition of methyl red indicator and 1 drop of 0.1 M NaOH) until the effluent is neutral. At this point, you will have collected all of the acid displaced form the ion-exchange resin. 75 to 100 ml of distilled water should be sufficient to wash all, the free HCl from the column.

Note: Record all of the drops of NaOH added to the small test solutions and consider them in the titration of the displaced acid.

Combine all of the effluent, including the test solutions, and titrate with the 0.1 M NaOH solution. Use this titration (accounting for the test drops of NaOH added) to calculate the exact molarity of your NaOH solution.

Repeat this procedure for all three KCl solutions and calculate and average molarity for your NaOH solution. You should have good agreement in your results.

Note: At the beginning of each day, and after every 2-3 elutions, the ion exchange column should be recharged with HCl.

4. Analysis of unknown salt.

Accurately weigh three (to nearest 0.1 mg) 0.2 to 0.3 g portions of your unknown sample. Dissolve each portion in 1-2 ml of distilled water. Pass each portion through the ion-exchange column using the same procedure as above. Titrate the eluted solutions with your standardized NaOH solution to the methyl red endpoint.

5. Calculate and report the number of millimoles of monovalent cation M⁺ per gram of your unknown sample.