Water Characterization

Reference Chapter: 6

Objective

After performing this experiment, the student shall be able to

• Measure several parameters that indicate the characteristics and differences of various types of natural water samples: surface water, groundwater (mineral water) and seawater.

Introduction

The parameters to be explored in this experiment will help us determine the main differences among water samples. They include pH, conductivity, chloride and sulfate concentration, and hardness level (as measured by the total amount of calcium and magnesium ions). Besides showing the pH, these parameters reveal the salt content of each sample, which normally varies depending on the source. For example, the main differences between surface and groundwater may lie in their salt content and turbidity, which is an indirect measure of suspended solids. The characteristics of surface and groundwater depend mainly on the nature of the catchment area, the type of soil present, and the materials in the confinement rock that retains the aquifer. In contrast, the characteristics of seawater are more constant and well known.

Experimental Procedure

The method for this experiment is to perform, on different samples, sequential measurements of the parameters discussed above.

The first step is to obtain samples. Preferably, you should obtain samples from original sources, such as river or lake water. For groundwater, take a sample from a well if possible, but if this is unfeasible, use a sample of bottled mineral water (preferably from a natural source). Seawater is an ideal sample, if it is available nearby; if it is not, prepare it synthetically or obtain it from a commercial source.

Samples must be collected in clean polyethylene bottles and analyzed immediately. This is mandatory for pH. For the other parameters, if immediate analysis is not possible, refrigerate the samples at 4°C and analyze within 48 hours.

Next, measure the conductivity/salinity and pH of the samples and compare the values with those for the same parameters of tap and distilled or deionized (D.I.) water. The conductivity measurements can be done with a conductivity meter, which in some cases is equipped for reporting the percentage of salinity in the sample as well.

The pH of the samples can be measured with a portable potentiometer or pH meter, preferably on the sampling site, because a maximum of 2 h is advisable for this measurement. Salinity (or an indirect measurement, such as the chloride ion concentration obtained by titration), provides interesting data for sample comparison.

Another parameter that will be measured and that also affects salinity values is the concentration of sulfate ions.

Total hardness, which helps to differentiate samples, as well as the amount of hardness due to calcium and magnesium ions in each sample, can be determined either by titration or by atomic absorption.

A. pH and Conductivity Measurements

Estimated time required: 1 min per sample Safety Measures

No special precautions are needed with these measurements because a measuring probe is introduced directly into the sample and the sample is not modified.

Materials

Reagents or samples

Conductivity

Take river or lake water samples. Do the

meter

same with seawater, bottled

commercial drinking mineral water,

and —when available, use also samples

of groundwater extracted from wells.

(Note: for synthetic seawater, prepare a

solution with commercial sea salt mix

for seawater aquariums)

pH meter with a

small-diameter

electrode

(for test tubes)

50-mL beakers

Experimental Sequence

Calibrate the conductivity and pH meters. Place each sample in a clean beaker.

1. Measure the conductivity/salinity percentage of the samples by introducing the conductivity probe. Rinse the probe with D.I. water between samples, and collect the rinses in a separate beaker.

2. Measure the pH of the samples by introducing the pH probe. The pH meter must be set up and ready in advance.

Note: When not in use, the pH meter must remain in STANDBY mode and the probe bulb must stay submerged in an appropriate solution (typically a pH

7 buffer, or a KC1 solution). The pH probe must be rinsed perfectly between measurements and before storage. Blot dry it before each measurement.

B. Chloride Concentration by Titration Applying the Mohr Method

In this part of the experiment the student will measure the chloride ion concentration in the water samples by promoting the formation of a white silver chloride precipitate, with the chromate ion as indicator of the endpoint (signaled by the formation of a red-brownish silver chromate).

Estimated time required: 15 min per sample

Safety Measures

Keep the titrant and the indicator from coming into contact with the skin or eyes because silver nitrate produces brown spots on the skin and the chromate indicator is toxic. All the residues generated in this experiment must be collected in a heavy metal residue bottle.

Materials Reagents

1 microburet (i.e., a 2-mL graduated pipet in 1/100, with a syringe connected to it by means of latex or Tygon® tubing)

2 2-mL volumetric pipets

1 propipet or a syringe (e.g., a 3-mL syringe with latex or Teflon tubing)

2 25-mL or 10-mL Erlenmeyer flasks

6 50-mL beakers (for the sample, for the titrant and for the buret rinsing) 1 support and clamps for the microburet 1 Beral pipet

1 wash bottle with distilled or

D.I. water 1 25-tnm plastic filter holder 1 10-mL syringe 1 wash bottle with D.I. water 0.7 |xm Nitrocellulose filter membranes (25 mm diameter) pH meter or pH indicator paper 1 thin spatula 1 bottle for residues

— Potassium chromate indicator (dissolve 0.05 g of K2CrC>4 in 10 mL of D.I. water)

- 0.01 N AgN03 solution

- Standard 0.01 NNaCl solution

— Activated carbon

- CaC03

Experimental Sequence

1. Secure the microburet in the stand with a clamp. Pour a small amount of standardized 0.01 N silver nitrate solution (see the *Note below) into a beaker and from this, take a small amount with the microburet (using the syringe); rinse the buret. Repeat this step with another small amount of the solution. Collect the rinses in a separate beaker. Fill the microburet with the same silver nitrate titrant up to the 2 mL mark.

2. If the sample is colored and shows high turbidity, add a small amount of activated carbon and pass the sample through a 0.7 |xm nitrocellulose filter. Use a clean, 10-mL plastic syringe to draw

5. Titrate drop wise, swirling the flask gently until the sample turns a pink-yellowish color. Note that a grayish precipitate forms, and at the end it takes on the color of the indicator. Record the volume of titrant used. Place the residues in the corresponding bottle, together with the rinse water used to wash the flasks.

Repeat the process with each water sample.

6. All the residues generated in the experiment must be collected in the bottle labeled for that purpose and disposed of according to local regulations.

The concentration of chloride ions in the sample is calculated with the following equation:

(V titrant, mL) (Weight-equiv. of CI") (Normality of AgN03) (1000 mg/g)

in the sample. Connect the plastic syringe to the filter holder containing the filter membrane, and let the liquid flow through by pushing the syringe plunger softly. Collect the filtered sample in another beaker. Filter approximately 20 mL of the sample (colorless and free of suspended solids).

3. Measure the pH of the sample (using either pH indicator paper or the pH meter). If the pH is below 5, add a small amount of sodium carbonate and swirl gently before the next step. If the pH values are strongly basic (>10), neutralize first with dilute sulfuric acid and then add some sodium carbonate.

4. Measure 2 mL of the water sample (previously filtered if necessary) with a volumetric pipet, and put it in a 25-mL or 10-mL Erlenmeyer flask. Add two to three drops of the chromate indicator and swirl. Observe the color. It must be greenish-yellow.

C. Sulfate Concentration Applying the Turbidity Method

In this part of the experiment the student determines the presence of sulfate ions in water by observing the generation of a barium sulfate precipitate. Then the turbidity produced in the sample is measured and related to the sulfate concentration through a calibration plot. The minimum detection limit with this method is 1 mg/L of sulfate ion.

Estimated time required: 35 min Safety Measures

Barium chloride is toxic and must not come into contact with the skin nor be inhaled. The residues from the experiment must be collected and deposited in a bottle labeled for disposal.

*Note: To standardize the silver nitrate solution, put a 2-mL sample of the 0.1 N standard NaCl solution in an Erlenmeyer flask and repeat the procedure outlined above. The volume of titrant used for titration of the standard sample will allow us to calculate the real concentration of the silver nitrate solution. Remember that at the equivalence point, the number of equivalents of the analyte equals the number of equivalents of the titrant. Use the equation below for this calculation:

Normality of AgN03 titrant:

(V, mL of NaCl std. sample) (Normality of the NaCl std.) (mL of the AgN03 titrant used)

Materials

2 2-mL, graduated pipets (1/100)

5 10-mL flat-bottom tubes or vials equipped with a cap and a micro magnetic stirrer

1 propipet or a syringe adapted with latex or Tygon tubing 1 1-mL graduated pipet 1 Beral pipet 1 10-mL beaker

1 spectrophotometer set at 420 nm

2 spectrophotometer cuvettes (5-mL or smaller)

1 5-mL syringe

0.7 (Jim Nitrocellulose filters

(25 mm diameter) 1 25-mm plastic filter holder 1 wash bottle with D.I. water

6 50-mL beakers pH meter or pH indicator paper 1 thin spatula magnetic stirring plate 1 bottle for residues

Reagents

- Na2SC>4 standard solution (150 mg/L). Note: 1 mL of this solution is equivalent to 100 micrograms of sulfate ion

- Conditioning reagent*.

'Dissolve 5 mL of glycerin in 3 mL of concentrated HC1 and 30 mL of distilled water; add to this mixture, 10 mL of isopropyl alcohol and 7.5 grams of sodium chloride. Mix everything until perfectly dissolved. This reagent must be prepared in advance.

Experimental Sequence

A. Prepare at least five dilutions of the sulfate standard solution. Apply the technique described below to analyze each dilution, and generate the data to build the calibration curve. Because the total volume of sample used with this method is 4 mL, prepare the following dilutions of the standard:

All of these dilutions must undergo the entire procedure that applies to the water samples.

All the water samples must be pre-filtered to free them of any solids.

B. Use the following technique with each dilution of the standard as well as with each water sample.

1. With a pipet, pour exactly 4 mL of sample into a vial with a magnetic microstirring rod and cap.

2. With the graduated pipet, add 0.6 mL of the conditioning reagent and stir magnetically. (Note: The conditioning reagent acidifies the medium in order to favor the precipitation reaction and to eliminate the possibility of precipitation of the barium carbonate that may form in highly alkaline waters. The glycerin favors the dispersion of the colloidal precipitate formed in the liquid medium, allowing a better turbidimetry measurement).

Add a small amount of barium chloride with the spatula. It is preferable to add an excess of this reagent, in order to favor the common ion effect and to accomplish the complete precipitation of the sulfate ions. Stir constantly.

3. Keep stirring a minute longer; then allow the mixture to stand for 2 minutes. Watch for any turbidity that may form. A turbidity lasting several minutes (while stirring) signals the presence of sulfate ions.

4. Turn on the spectrophotometer and set it at 420 nm. Allow it to warm up. Calibrate to 100% transmittance with a spectrophotometer cuvette containing D.I. water.

Any barium-treated sample that shows turbidity must be well mixed and poured into a spectrophotometer cuvette and inserted into the spectrophotometer. Immediately afterwards, read the transmittance or the absorbance continuously at 30-second intervals for 4 minutes.

The maximum absorbance (turbidity) is generally obtained within 3^4- minutes, and sometimes this value lasts for one minute or longer. This maximum value is used for calculating the sulfate concentration.

The experimental conditions must be the same for all samples in order to obtain reliable values.

5. To determine the sulfate concentration of the water sample, you must first plot the calibration curve with the data of absorbance vs the corresponding concentrations (i.e., from the dilutions of the standard ion). Once you know the absorbance value of the sample, you can calculate the concentration of sulfate ions in the water sample.

6. All the residues generated in the experiment must be collected in the bottle labeled for that purpose.

D. Total Hardness, Calcium, and Magnesium Ion Concentrations

The student will determine total hardness as well as the calcium and magnesium ion content in a water sample applying the EDTA titration method.

Estimated time required: 10 min per sample

Safety Measures

In this case, the student must proceed with caution in handling the sodium hydroxide for the pH adjustment. In the event of a spill or skin contact, wash with abundant water. The residues generated can be neutralized and flushed down the drain.

Materials

Reagents

1 microburet (i.e., a 2-mL graduated pipet in 1/100, with a syringe connected to it by means of latex or Tygon® tubing)

2 small spatulas

1 2-mL volumetric pipet 1 2-mL graduated pipet (in 1/100) 1 propipet or a 3- or 5-mL syringe (adapted with a small latex or Tygon® tube) 6 25- or 10-mL Erlenmeyer flasks 4 10-mL beakers 1 pH meter equipped with a small-diameter combination pH electrode (for test tube insertion), or pH indicator paper

1 stand and a clamp for the microburet

2 Beral pipets

1 wash bottle with distilled or D.I. water

2 M NaOH solution

(for pH adjustment) Solid murexide indicator Solid eriochrome black indicator 0.01 M or 0.001 M Na2EDTA (sodium ethylenediamine tetra acetate) solution pH 10 buffer (NH3 / NH+)

Experimental Sequence

1. Measure the pH of the sample with a pH meter or pH indicator paper. Fill the microburet with the concentrated EDTA solution, and adjust to a known volume.

2. In order to measure the total hardness value (i.e., the Ca2+ + Mg2+ concentration), place a 2-mL, solids-free water sample (measured with a volumetric pipet) in an Erlenmeyer flask. Add 23 mL of the pH 10 buffer, swirl and add one or two crystals (or a small amount of powder) of the Eriochrome black solid indicator. Swirl until total dissolution. The mixture should now appear with a red wine color. Titrate this with the EDTA solution to a dark-blue endpoint. If the amount of titrant needed to reach the endpoint is too small to be measured, repeat the titration with another sample using a more dilute titrant (e.g., 0.001 M EDTA). Note: If the blue color appears from the start, this means there is no measurable hardness in the sample.

3. With a volumetric pipet, put 2 mL of the sample (free of solids) in a 25-mL Erlenmeyer flask. Add 1 mL of 2 M NaOH to ensure that the pH is frankly basic (pH — 11 >; add one or two crystals of solid murexide indicator (or a small amount of its powder), and swirl softly until they dissolve. Titrate with 0.01 M EDTA to a violet endpoint. If the amount of titrant needed to reach the endpoint is too small to measure, repeat the titration with another sample using a more dilute titrant (e.g., 0.001 M EDTA). This value will let us know the calcium ion concentration in the sample. Repeat this method with each water sample.

4. One mole of EDTA is consumed for each mole of Ca2+ or Mg2+. Because the MW of calcium carbonate is virtually equal to 100, then the concentration of Ca or Mg (expressed as mg/L of calcium carbonate) can be calculated with the following equation:

Total or Ca2+ hardness, mg/L as CaCC>3

_ (Vtitra„, x EDTA Molarity x 100 g/mol x 1000 mg/g)

^sample

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Section

1. Water Characterization Date__

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