CHAPTER XVII:
LIGHT-SCATTERING AND MOLECULAR SPECTROPHOTOMETRY

 

A. Introduction

 

Light or electromagnetic radiation can be characterized by wave-like behavior. The wavelength, , is the distance between the crest of one wave and the crest of the next one. The frequency, , is the number of complete waves that pass a reference point in a second. These two are related by the velocity of light, c (3 x 1010 cm/sec), in accordance with equation 17.1.

 

(17.1)

 

Alternatively, light may also exhibit behavior resembling that of a particle. These particles, called photons, have a unique energy, E, that is dependent on the frequency (equation 17.2) through Planck's Constant, h.

 

(17.2)

 

As a result of equations 17.1 and 17.2, the energy of a photon of light is inversely proportional to its wavelength.

 

 

 

 

The human eye perceives light of different wavelengths as being comprised of separate colors. When a substance absorbs light of a certain wavelength (color), it depletes the transmitted light in precisely those wavelengths that are absorbed. For this reason, a solution that appears blue to the human eye does not absorb blue light. Rather, it is blue because it absorbed orange light, and allows all other wavelengths to pass unhindered (giving a bluish color to the transmitted light; see Table 17.1).

 


Table 17.1

Relationship Between Color and Absorbance

 

Wavelength of absorbance maximum (nm)

Color Absorbed

Color Remaining

380-420

Violet

Green-yellow

420-440

Violet-blue

Yellow

440-470

Blue

Orange

470-500

Blue-green

Red

500-520

Green

Purple

520-550

Yellow-green

Violet

550-580

Yellow

Violet-blue

580-620

Orange

Blue

620-680

Red

Blue-green

680-780

Purple

Green

 

 

Double bonds are strong absorbers of light, especially when they are arranged in an alternating pattern with carbon-carbon single bonds. This alternating pattern of single bond followed by double bond followed by single bond followed by double bond is called a conjugated double bond system. It has a property known as resonance, which makes it absorb light of high wavelengths (i.e., low energy). The molar absorptivity and wavelength of maximim absorbance generally increases with increasing numbers of conjugated double bonds. This is illustrated by the series of polynuclear aromatic hydrocarbons, below.

 

 

 

 

 

When a solution is placed in a path of light, a certain amount of that light will pass through it unaffected. The fraction of the incident light that reaches the other side of a liquid sample will be dependent on the pathlength through the sample, x, and the ability of the sample components to block or otherwise impede the passage of light, , called the extinction coefficient. This relationship is given by Lambert's Law:

 

(17.3)

or

(17.4)

 

where:

Io = the intensity or radiant power of the incident light

I = the intensity or radiant power of the light after passing through a though a thickness, x, of the liquid sample

 

The blockage or impedance of light may be due to scattering or absorption. Light scattering occurs when small particles deflect the light so that it does not reach the other side of the sample. Absorbance is the process by which a constituent (usually dissolved) absorbs the light energy and releases it as heat or stores it as bond energy. Thus, the extinction coefficient in equation 17.3 and 17.4 is composed of two components, the scattering coefficient () and the absorption coefficient ().

 

(17.5)

 

 

 

B. Light Scattering

 

When light scattering is large compared to absorption, (i.e., >> ), Lambert's Law reduces to:

(17.6)

 

The scattering coefficient may be expressed as the product of the concentration of particles, N, and the scattering cross section of those particles, .

 

(17.7)

 

The scattering cross section is a measure of the amount of light scattered per particle. It is a function of the size and shape of the particle. For particles of extremely small size (i.e., a diameter 10% of the incident light's wavelength), Rayleigh Scattering predominates. The Rayleigh Scattering cross section for an ideal spherical "particle" is given by equation 17.8.

 

(17.8)

 

Thus, the Rayleigh Scattering depends only on the pathlength, the concentration of "particles", the wavelength of the light, , and the polarizability of the light, . Equation 17.8 predicts that the smaller wavelength light will be scattered the most. Therefore, blue light should be scattered more than the other visible wavelengths. As a result, a clear daytime sky appears blue when looking at any direction except directly into the sun.

 

Larger particles are subject to Mie Scattering. With these large particles, light may scatter all over their surfaces at different angles. Mie Scattering is the type of light scattering that is responsible for turbidity in natural waters. It is far more complex than Rayleigh Scattering, and cannot be described by a simple equation. The Mie Scattering cross section is a function of the particle size and shape as well as the wavelength of light. Maximum cross section for a given wavelength occurs when the particles present are of a size that is similar to the wavelength. Less scattering occurs as the difference between the particle size and wavelength increases.

 

 

1. TURBIDITY

 

a. Environmental Significance

 

Historically, turbidity is considered a measure of the attenuation of light through a sample of water. However, modern turbidimeters are based on nephelometry, which is the measurement of true light scattering. For this reason, we will use the terms turbidimetry and nephelometry interchangeably to indicate the measurement of light scattering by small particles in a water. These particles may be composed of clays, bacteria, algae, or colloidal organic molecules. Accordingly, the sizes of these particles may range over several orders of magnitude. Turbidity-producing particles are of concern in drinking water supplies for a variety of reasons. Cloudy or turbid waters are aesthetically unpleasing. Thus, consumers may be tempted to forgo an otherwise acceptable disinfected drinking water for one that is not disinfected, but clear. Turbid waters may be less filtrable as a result of the high load of particles applied to the filters. This would obligate shorter filter runs and add to the expense of the finished water. Finally, turbidity-producing particles may serve to protect micro-organisms (e.g., pathogens) and render disinfection processes less effective.

Turbidity measurements are used to assess the quality of water supplies. Knowing raw water turbidities, an engineer can suggest a treatment sequence that is likely to be economically and technologically feasible for any particular source (e.g., conventional treatment vs direct filtration vs in-line filtration; see Weisner et al., 1988). A knowledge of turbidity may also help the engineer to estimate unit sizes and chemical costs (e.g., for coagulation and sludge handling). However, pilot studies must often be performed before a treatment plant is designed; and here again, turbidity measurements are used to assess the effectiveness of most treatments processes one would test. Finally, turbidities are monitored in full-scale treatment installations both for the purpose of process control and for the purpose of establishing compliance with finished water turbidity standards.

 

b. Measurement: Turbidimetry

 

Modern methods for determining turbidity employ nephelometric turbidimeters which measure the intensity of light scattered at 9030 with respect to the incident light. The detector should be sensitive to light in the range of 400-600 nm wavelength. This measurement is then compared to an identical measurement of an accepted turbidity standard. The standards have pre-designated turbidities expressed in nephelometric turbidity units (NTU). Accepted standard substances include formazin, a co-polymer of hydrazine sulfate and hexamethylenetetramine, and commercially-available polystyrene divinylbenzene beads (AEPA-1).

In the UMass Environmental Engineering Teaching Laboratory we have a Hach Model 2100A laboratory turbidimeter. Light from a tungsten filament lamp is projected up through a flat-bottom sample cell. A 90 slit allows scattered light to reach the photomultiplier tube. The photomultiplier is especially sensitive to the ultraviolet light scattered by small particles (i.e., 0.1 to 0.5 microns). The 2100A has a rated bias and precision of -2% of full scale on all ranges. Thus, it is always advisable to use the highest range while still being able to remain on scale. The ranges are 0-0.2, 0-1.0, 0-10, 0-100 and 0-1000 NTU; and the detection limit is 0.04 NTU on the lowest range.

It is not uncommon that the same sample, measured on two different nephelometric instruments may give different turbidity readings. This can occur even though the two instruments may be calibrated against the same standards. Differences in instrument design, bulb age and output voltage may account for this. For example, if the voltage applied to the bulb drops with time, the tungsten filament will become cooler and the output spectrum will shift toward longer wavelengths. Since larger particles are better at scattering light of longer wavelength, this instrument will become more sensitive to larger particles and less sensitive to smaller ones. A formazin standard which has relatively large particles allows one to calibrate the instrument to the larger particles. However, the instrument will no longer be sufficiently responsive to the smaller particles in a water sample. One way of avoiding this problem is calibrate against two different types of standard materials (e.g., formazin and AEPA-1; these have a diameter range of 1.75-20m and 0.2-0.8m, respectively) and adjust the bulb voltage (or replace the bulb, if necessary) until the instument can be calibrated so they both give correct readings. Another source of error is the presence of stray light reaching the photo sensor. Stray light may be determined using a methanol standard.

 

 

 

C. Molecular Spectrophotometry

 

1. THEORY

 

When absorption is large compared to light scatter, (i.e., >> ), Lambert's Law reduces to Beer's Law:

 

(17.9)

 

The absorption coefficient may be expressed as the product of the concentration of absorbing substances, c, times their molar absorptivity, a.

 

(17.10)

 

Therefore, the transmittance (T = I/Io) of a liquid sample is given by equation 17.11.

 

(17.11)

 

A more convenient measure of the absorption of a liquid sample is the logarithm of the reciprocal of the transmittance. This value, called the absorbance (A), is linearly related to the concentration of the absorbing substance.

 

A = -log(T) (17.12)

 

A = acx (17.13)

 

Absorbances may also be presented as absorbance units per pathlength, or simply reciprocal pathlength (e.g., as cm 1). We will symbolize this representation of absorbance with Abs.

 

Abs = A/x = ac (17.14)

 

Equation 17.14 predicts that a plot of Abs versus concentration will give a straight line with a slope of "a". The absorptivity is commonly expressed in terms of absorbance units per mole/L. This value is called the molar absorptivity.

 

 

 

 

2. MEASUREMENT

 

a. Spectrophotometers

 

Sample absorbances are measured with a spectrophotometer. This is an instrument composed of a light source, a wavelength selector (monochromator), a sample compartment, and a detector. The type of light source will depend on the desired wavelength. Light in the visible range (350nm-800nm) is best supplied by a tungsten filament lamp. Ultraviolet (UV) light (190nm-350nm) requires a deuterium lamp. The wavelength selector splits the light spectrum into its component colors and selects a narrow band of this spectrum. Either a prism or a grating may be used as a dispersing element to split the light. The spectrum is generally projected on an opaque wall containing a slit. The physical width of the slit determines the band width of the selected light. The wavelength of the selected light may be adjusted by changing the angle of the dispersing element, so that the desired wavelength passes through the slit. The selected light passes into the sample compartment, through the sample (and cell) and to the detector (photomultiplier). By measuring transmittance with (I) and without (Io) the sample present, absorbance can be determined. This is how a single beam spectrophotometer operates.

A more accurate and convenient scheme is embodied in the double beam spectrophotometer. This instrument has a chopper motor which alternately deflects the light beam through a reference cell and the sample cell. This is done many times per second, and the average ratio between the two readings gives the transmittance. This is a more accurate method, because it minimizes variabilities due to rapidly changing lamp output or momentary stray light in the detector compartment. In effect, the analyst is continually monitoring and adjusting for changes in the lamp output as measured by the detector.

In the Environmental Engineering teaching laboratory we have several spectrophotometers. One is a Perkin-Elmer Model 111 Ultraviolet-Visible Spectrophotometer. This is a typical single beam instrument with a spectral range of 200-900 nm. It uses a diffraction grating with fixed bandpass of 2.0 nm. The wavelength accuracy is 0.5 nm and the precision is 0.1 nm. The photometric linearity is -0.005 absorbance units at a reading of 0.400 absorbance units. The photometric reproducibility is better than 0.01% Transmittance, and stray light is less than 0.5% T at 220 nm. This instrument also has a stability characterized by a drift of less than 0.5% T per 5 min after a 15 min warm-up period. As with most UV-Vis spectrophotometers, the PE 111 uses a Tungsten lamp in the visible region (900-340 nm) and a Deuterium lamp in the ultraviolet region (340-200 nm).

For best operation, spectrophotometers should be installed in air conditioned rooms, free from dust, corrosive fumes, vibrations, and large changes in temperature and humidity. Instructions for calibrating and operating the PE 111 are as follows:

 

1. With the operation switch in the off position, verify that the meter mechanical zero is correct.

2. Select the proper lamp for the desired wavelength by adjusting the selector lever to either VISIBLE or ULTRAVIOLET.

3. Adjust the wavelength knob to the correct value.

4. Turn on the appropriate lamp. If the deuterium lamp is chosen, the #1 switch must be turned on first, then 20-30 sec later the #2 switch is turned on.

5. Turn the operation switch to "on".

6. Open the cell compartment and insert both filled sample and reference cells in the cell holder. The reference cell belongs in the position #1.

7. Place the operation switch to the "meter" position.

8. With the cell compartment cover open (this automatically closes the shutter), adjust the meter to infinite absorbance (0 transmittance) with the zero adjusting knob.

9. Close the cell compartment cover, and place the reference cell in the light path by adjusting the cell positioning knob. Adjust the meter to 0 absorbance (100% transmittance) with the zero adjusting knob.

10. Pull out the cell positioning knob and read the absorbance of the sample.

 

 

Also in the environmental engineering laboratory we have a Spectronic 21 capable of measurement in the UV as well as visible range. This is also a single beam instrument that requires adjustment to 100% transmittance (or zero absorbance) prior to use.

 

 

 

 

 

 

 

 

 

 

Figure 17.1 Controls on the Spectronic 21D UV-Vis Spectrophotometer

 

 

 

b. Spectrophotometric cells

 

Cells (or cuvettes) are supplied with varying light paths and different qualities of glass. Most spectrophotometric work is conducted with standard rectangular cells of 1 cm path length. These are square in cross-section and about 4 times as high as they are wide. They generally have a pathlength tolerance of -0.01 mm. In addition, rectangular cells of 0.1, 0.5 and 4 cm are commercially available. For more dilute solutions, cylindrical cells of 5 and 10 cm are available. These have two filler necks into which fit fluoropolymer stoppers. The ability to seal these cells is an attractive feature for volatile or hazardous samples. The cylindrical cells may be supplied with fluoropolymer covers, but these do not make a very good seal. Rectangular cells with extra thick side walls are also available for the analysis of small volumes of liquid. In addition, various types of flow through and jacketed cells are available for kinetic studies, continuous monitoring, and temperature sensitive work.

Cells for absorption spectrophotometry have opposing optically polished windows with frosted glass side walls. Be sure to orient the cells so that the optic windows are perpendicular to the light path. Fluorescence cells are made with all (or adjacent) optically polished windows.

Cells should be cleaned like other laboratory glassware. If detergents prove ineffective, they may be soaked in chromic acid cleaning solution. Note that the windows and side walls are fused forming an acid-proof seal. In addition, quartz cells may also be cleaned by immersing in concentrated nitric acid followed by ultrasonic agitation or 10-15 minutes of boiling in water.

 

 

Table 17.2

Spectrophotometric Cells

 

Window

Wavelength Range (nm)

Letter Code

Lot #s

Color Code

Optical Glass

360 - 1000

OG

 

yellow

Near-UV Glass or Special Optical Glass

300 - 1000

OS or SG

180's

green

Standard Silica

220 - 2500

 

 

 

Supracil Quartz or Quartz UV

165 - 2600, 2850 - 3600

QS or UV

280's

blue

Infracil Quartz or Quartz IR

220 - 3600

QI or IR

300's

red

 

 

Most spectrophotometric cells sold in the US are marked with a series of letters and numbers. Many are marked with and upper case "SCC" (Scientific Cell Company) and number indicating the path length (e.g., 1.000 for 1 cm) followed by a two letter code, a colored dot or a lot number. These markings will identify the type of glass used in the optical windows (see table below). Cells should always be matched with the same markings including lot number. This is especially important when running absorbance scans. It is of highest importance to choose a cell which is useable for the wavelength chosen. Although any of the cells listed below may be used in the visible range, the optical glass cells are preferred because of their lower cost. For near-UV and UV work, near-UV glass, Supracil quartz or Infracil quartz should be employed depending on the wavelength chosen. In general, the lower the minimum wavelength, the more expensive the cell. For work in the infrared region, Infracil quartz is required.

 

 


 

3. ANALYTICAL MOLECULAR SPECTROPHOTOMETRY

 

a. Direct Methods

 

Most specific chemical analytes do not have large enough absorptivity to make their determinations by direct spectrophotometry practical. However, their molar absorptivities may be substantially increased by reaction with color-forming reagents. This is the basis for most colorimetric or spectrophotometric methods. The two exceptions common to environmental engineering are the determination of "color" and "UV absorbing substances". However, color and UV absorbance are somewhat unique in that they are gross parameters which are operationally defined. They are both commonly used to assess the concentration of natural organic matter (e.g., humic substances) in a water. Humic substances absorb strongly over a wide range of wavelengths (see section b). Absorbance measurements will depend on the concentration of the humic substances, the flora and fauna from which the humics were derived, and the solution pH. This last factor is particularly important with respect to the analysis of UV absorbance and color. As pH increases, UV absorbance and particularly, color increase. This is likely due to deprotonation of acidic sites on the humic molecules, causing an unfolding of the molecules and increasing the multiplicity of resonance structures. Regardless of the mechanism, it is important that a water be buffered at a fixed pH (usually 7) prior to analysis for UV absorbance and color.

 

Color. Natural waters derive color from humic materials and sometimes from metals such as iron and manganese. This is a common means of characterizing a natural water's organic content. Color is very easy to measure, requires only the simplest glassware, and correlates well with DOC or TOC in many water treatment systems. Color is also important because it is a direct measurement of a water's disagreeable visual properties. As a result many treatment systems are designed to remove color whereas the specific goal of removing organic carbon is secondary.

Two types of color are often reported, "true color" and "apparent color". True color is that which is attributed to dissolved species, and it is therefore, the color remaining after a sample is filtered. Apparent color is derived from the absorbance of dissolved species and the light scattering of particles. It is, therefore, measured without prior filtration. The actual determination of color is done by visual comparison or with the aid of a spectrophotometer.

The platinum-cobalt method of visual comparison is the classic procedure used for decades. Color standards are prepared by mixing predetermined amounts of potassium chloroplatinate and cobaltous chloride in water, and diluting this stock 10 to 100 fold. The prescribed ratio of cobalt to platinum gives a yellowish hue characteristic of natural waters. Filtered or unfiltered samples are pH-buffered and placed in 50 mL Nessler tubes. Platinum-cobalt standards are added to similar tubes. Concentrations are in Pt-Co color units. These correspond numerically to the concentration of platinum in mg/L in the standard which most closely resembles the water. Color analysis in the field may be more conveniently conducted with pre-calibrated glass disks.

Color may also be conveniently and precisely measured using a spectrophotometer. The range of wavelengths chosen run from 400 nm to 700 nm, with 400 nm being most common. As before, the samples are buffered at pH 7 and pre-filtered if desired. The color value is simply the absorbance of the sample at the designated wavelength.

 

UV Absorbance. Most natural organic matter will absorb sufficient ultraviolet (UV) light to be easily detected by a standard UV-Vis spectrophotometer. By convention, we have chosen 254 nm as the wavelength to measure UV absorbance[1]. This parameter is quite important because: (1) it is inexpensive, rapidly measured, and requires a minimum of training; and (2) it has been found to correlate with certain water quality characteristics, such as DOC and THMFP.

UV absorbance has been successfully used as a means of estimating DOC and THM precursor levels in raw waters. However, its most important contribution is to process monitoring. For a single raw water source, coagulation effectiveness can be effectively monitored by UV absorbance. One can generally develop good linear correlations between UV abs and DOC for raw and treated waters from the same plant (e.g., Edzwald et al., 1985). The interpretation changes, however, when a disinfection or oxidation step is encountered. When monitored across oxidation/disinfection, UV absorbance provides information on the degree of oxidation of the natural organic matter in the water.

The specific absorbance, which corresponds to the absorbance per mg/L of DOC, is a useful tool for rapidly assessing the "humic/non-humic nature" of a water. Some specific absorbances for extracted humic and non-humic fractions are shown in Figure 17.2. Note that humic acid and fulvic acid show the highest SUVA (6.3 and 4.4, respectively). In another study, averages of 10 aquatic humic substances showed SUVA values of 5.8 for the humic acids and 3.6 for the fulvics (Reckhow et al., 1990). Other fractions, especially the hydrophilic acids, show lower SUVA values. For this reason, waters with a high SUVA generally have higher humic contents, and are more amenable to DOC removal by coagulation.

 

Table 17.3 summarizes some attempts to correlate UV absorbance (254 nm) to DOC for raw waters. Note that the reciprocal of the slopes in Table 17.3 correspond to the specific absorbances in Figure 17.2, and that the average slope (~25) gives a specific absorbance of 0.004/cm or 4/m which is similar to those reported for fulvic acids.

 

 

 

Figure 17.2

Specific UV Absorbance of Eight NOM Fractions from Forge Pond

(After Reckhow et al., 1993)

 

 

 

 

 

 

Table 17.3

UV Absorbance (254 nm) to DOC Regression Equations+

 

Water

Slope

Intercept

r

Reference

Tjeukemeer Lake

24.7*

2.7

0.925

De Haan et al., 1982

Grasse River

19.8

0.65

0.93

Edzwald et al., 1985

Glenmore Reservoir

28.1

-0.3

0.71

Edzwald et al., 1985

 

24.1

2.2

0.96

Smart et al., 1976

N.C. Surface Wat.

29.1

1.44

0.83

Singer et al., 1982

Norwegian Lakes

19.1

2.3

0.93

Vik et al., 1985

 

+Model: DOC (mg/l) = (slope)*(UV absorbance, cm-1 ) + intercept

* Based on absorbance at 250 nm, neutral pH.

 

 

 

b. Methods requiring the Formation of a Chromophore

 

As discussed obove, most specific chemical analytes do not have large enough absorptivities to make their determinations by direct spectrophotometry practical. However, through a variety of means, one can increase their molar absorptivities. This is most often done through the formation of a chromophore, a chemical structure that imparts a visible color to a solution. By far the most common group of reactions that is used to form chromophores for spectrophotometric analysis are the complexation reactions (see Table 17.4). Often extractions into organic solvents or redox reactions are needed to improve selectivity (i.e., remove interferences) or sensitivity (i.e., lower detection limit). Other colorimetric methods are based largely on redox reactions (see Table 17.5). This is especially true of analytes that are, themselves, reductants or oxidants (e.g., chlorine). Another class of colorimetric methods relies on the catalysis of an observeable reaction by the analyte (Table 17.6). Here the rate of reaction must be directly proportional to the analyte (catalyst) concentration. These are potentially very sensitive methods. Finally, Table 17.7 lists a number of methods that employ coupling reactions, simple substitutions, or incompletely characterized reactions.


Table 17.4

Colorimetric Methods Based on the Formation of a Complex

 

Analyte

Complexing Metal or Ligand

Additional Treatment

Wavelength

Al

Eriochrome cyanine R dye

 

535 nm

Be

Aluminon (aurintricarboxylic acid ammonium salt)

 

515 nm

Cd

dithizone (diphenylthiocarbazone)

post-extraction (CHCl3)

518 nm

Cr(VI)

diphenylcarbazide

 

540 nm

Cu

neocuproine (2,9,-dimethyl-1,10-phenanthroline)

post-extraction(CHCl3& CH3OH)

457 nm

Cu

bathocuproine disulfonate

 

484 nm

Fe

1,10-phenanthroline

pre-reduction (hydroxylamine)

 

Pb

dithizone

pre-extraction (CCl4)

520 nm

Hg

dithizone

pre-extraction (CHCl3)

490 nm

Ni

dimethylglyoxime

pre-extraction (CHCl3,then H2O)

445 nm

Se

diaminobenzidine

pre-oxidation, pre-reduction post-extraction (C6H5CH3)

420 nm

Zn

dithizone

pre-extraction (CCl4)

535/620 nm

Zn

zincon

 

620 nm

As

silver diethyldithiocarbamate

pre-reduction (Zn)

535 nm

B

curcumin

 

540 nm

B

carmine

 

585 nm

SCN

Fe(+III)

 

460 nm

F

SPADNS

reaction (zirconium)

570 nm

PO4

Mo

post-reduction (SnCl2)

690 nm

PO4

Mo

post-reduction (ascorbic acid)

880 nm

Silica

Mo

 

410 nm

SO4

methylythymol Blue

pre-precipitation (Ba), chelation of excess Ba & detn of excess ligand

460 nm

 


Table 17.5

Colorimetric Methods Based on Redox Processes

 

Analyte

Oxidant or Reductant

Additional Treatment

Wavelength

Mn

persulfate or periodate (to Mn(+VII))

 

525 nm

K

dichromate

pre-precipitation (sodium cobaltinitrite) & oxdn of excess

425 nm

Br(-I)

chloramine-T (to Br(+1))

post-reaction (bromination of phenol red)

590 nm

Cl(+I) & ClO2

DPD

 

515 nm

CN

chloramine-T (to CNCl)

post-reaction (pyridine- barbituric acid)

578 nm

I(-I,+I)

Leuco Crystal Violet

pre-oxidation (potassium peroxymonosulfate)

592 nm

NH3

HOCl

post-reaction (phenol) (to indophenol, phenate method)

630 nm

O3

Indigo Trisulfonate

 

600 nm

Tannins & Lignins

tungstophosphoric & molybdophosphoric acids

 

700 nm

 

 

 

 

 

 

 

 

Table 17.6

Colorimetric Methods Based on Catalysis

 

Analyte

Reaction

Additional Treatment

Wavelength

V

oxidation of gallic acid by persulfate

 

415 nm

I(-I)

reduction of Ce(+IV) by arsenious acid

remaining Ce(+IV) reduced by Fe(+II), residual measured at

410 nm

 

 

 

Table 17.7

Colorimetric Methods Based on Miscellaneous Reactions

 

Analyte

Reactions/Reactants

Additional Treatment

Wavelength

NH3

KI + HgI2

 

410 nm

NO3/NO2

diazotization with sulfanilamide then coupling with N-(1-naphthyl) -ethylenediamine)

pre-reduction (Cd) (to nitrite, for nitrate analysis only)

543 nm

NO3

Chromotripic Acid

 

410 nm

S(-II)

S(-II), FeCl3 & dimethyl-p- phenylenedeamine (forms methylene blue)

 

664 nm

Phenols

4-aminoantipyrine, potassium ferricyanide

post-extraction (CHCl3)

460 nm

 

 

b. Measurement of Al and Fe

 

Iron and aluminum are polyvalent metals that are used as coagulants in drinking water treatment. They are also present in natural waters at usually low concentrations. In properly operating drinking water treatment plants, metal coagulants are well removed by sedimentation or filtration. However, extremes of pH, incorrect dosing, or poor hydraulics may result in the breakthrough of high levels of residual iron or aluminum in the distribution system. This is undesirable for several reasons. First these hydrolysing metals may continue to slowly precipitate far downstream of filtration. This will cause hydraulic and water quality problems in any engineered systems that does not have provisions for sludge removal (i.e., clearwells, water mains, etc.). Iron may actually impart a metallic taste at high concentrations, and it is well known to stain plumbing fixtures and laundry. Under certain circumstances iron residuals can support the growth of unwanted bacteria. In the U.S. a secondary maximum contaminant level has been set for iron at 300 g/L (based on aesthetics). The corresponding European recommended standard is 50 g/L, with an MCL of 200 g/L. High concentrations of aluminum have been linked to encepalopathy in kidney dialysis patients. There is also some weak evidence that it may be associated with Alzheimer's Disease as well. In Europe, residual aluminum concentrations in drinking water are limited to 200 g/L, and the recommended limit is 50 g/L. A primary drinking water standard (most commonly 50 g/L) is currently being discussed for the U.S.

 

Aluminum and iron may be determined in aqueous solution together by the simultaneous formation of of two separate colored complexes; Al-ferron, and Fe-phenanthroline (Davenport, 1949). The formation of an Fe-phenanthroline complex must be preceeded by two steps. First iron hydroxides must be dissolved with acid (equation 17.15a), then the iron must be reduced to the ferrous state with hydroxylamine (equation 17.15b).

 

Fe(OH)3 + 3HCl ------> Fe+3 + 3H2O + 3Cl- (17.15a)

 

4 Fe+3 + 2 NH2OH ------> 4 Fe+2 + N2O + H2O + 4H+ (17.15b)

 

The ferrous-iron then forms a strongly-colored complex with three molecules of 1,10-phenanthroline. Binding occurs at the heterocyclic nitrogens. This complex is reported to have a molar absorptivity () of 11,100 M-1cm-1 at max = 508 nm (i.e., orange-red). The intense color is attributed to charge transfer of an electron from Fe(II) to a vacant * orbital on the phenanthroline. Such a transition cannot occur with Fe(III). As a result, this latter complex only possesses a slight bluish color due to the weaker ligand-field transition phenomena.

Aluminum forms a colored complex with several hydroxyquinoline derivatives. One of these ligands, 8-hydroxy-7-iodo-5-quinolinesulfonic acid (or Ferron; abbreviated as "F" below), gives an absorbance maximum at about 370 nm. This reaction is most sensitive at pH 5.5, and it is subject to interference from iron. However iron interference can be conveniently removed by simultaneous complexation with 1,10-phenanthroline (abbreviated as "P" below). In this way, a combined ferron/phenanthroline reagent allows the simultaneous determination of iron and aluminum when absorbances at 370 nm and 520 nm are recorded. For example, the overall absorbance measured at either of these two wavelengths can be broken down as follows:

 

(17.16)

 

where:

= the Total absorbance of a sample or standard at wavelength "x"; i.e. the absorbance that you directly measure with a spectrophotometer.

= the absorbance of the sample or standard prior to addition of complexing reagents (i.e., initial sample absorbance) at wavelength "x"

= the fraction of the total absorbance that is due to free, un-complexed Ferron at wavelength "x"

= the fraction of the total absorbance that is due to free, un-complexed Phenanthroline at wavelength "x"

= the fraction of the total absorbance that is due to Al-Ferron complex

= the fraction of the total absorbance that is due to Fe-Phenanthroline complex

 

 

Invoking Beer's Law and assuming that all aluminum and iron are completely complexed, we presume that:

 

(17.17)

(17.18)

(17.19)

(17.20)

 

where is the absorptivity (usually in: abs units/mg/L) of substance "y" at wavelength "x". The free ligand concentrations are:

 

(17.21a)

 

(17.21b)

 

where "n" and "m" are stoichiometric factors that relate the amount of ligand bound per unit of metal. Now combining equations 17.16 through 17.21b one gets:

 

(17.22)

 

and this reduces to:

 

(17.23)

 

And if we define the absorbance of the added ligands in their free state (i.e. initial ligands) as ,

(17.24)

 

we get:

 

(17.25)

 

Now we know that aluminum can be best measured near the absorbance maximum for the aluminum-ferron complex (370nm), and iron is best quantified at a wavelength near the iron-phenanthroline maximum (520nm). From this we can used equation 17.25 to formulate an expression for Fe and Al concentration:

 

(17.26)

 

(17.27)

 

The first term of both equations 17.26 and 17.27 (AbsT) is simply the final absorbance of the samples or standards after treatment with the ferron-phenanthroline reagent. The second term (Abssi) is just the absorbance of the samples before anything is done to them. The third term (Absli) is obtained from the intercept of the standard curves, and the constant parts of the last two terms are determined from the slopes of these standard curves.

 

 

c. Ammonia-Nitrogen: The Phenate Method

 

Ammonia nitrogen may be determined by reaction with phenol and chlorine which gives a new highly colored condensation product, indophenol (Bolleter et al., 1961). The reaction is thought to occur through the formation of monochloramine (equation 17.28). For quantitative formation of monochloramine, the pH must be around 6.5 to 7.0.

 

NH3 + OCl- NH2Cl + OH- (17.28)

 

Then this powerful electrophile attacks the anionic phenate to give a quinonechloramine (17.29). This intermediate continues to react with phenate forming the highly-conjugated indophenol (17.30). Manganous sulfate is added as a catalyst.

 

(17.29)

 

(17.30)

 

The color of this final product is quite pH-dependent, yellow at low pH, and blue at high. The blue color is most intense at pH 9.9-10.0.

 

 

Procedure

1. To 10 mL of sample in a 100-mL beaker, add 10 mL of high-purity water and 2 drops (~0.1 mL) of manganous sulfate solution.

2. Place on a magnetic stirrer and add 1 mL conc. HOCl solution, followed immediately by the dropwise addition of 1.2 mL phenate reagent. Stir vigorously during the addition of the reagents.

3. Wait 10 min after addition of the reagents, then measure the absorbance at 630 nm.

4. Repeat steps 1-3 for all samples along with a set of 4 standard solutions prepared from the 25 mg/L ammonia solution. The concentrations you choose for the standard solutions should be properly spaced such that the unknowns all fall between the highest and lowest standard.

 

 

 

Reagents

1. Conc. HOCl solution: Dilute 20 ml of the 5% NaOCl stock to 100 mL. Adjust pH to 6.5 - 7.0 with HCl.

2. Manganous sulfate solution (0.006 N): Dissolve 50 mg monohydrate manganous sulfate in 100 mL distilled water.

3. Phenate reagent: Dissolve 2.5 g NaOH and 10 g phenol in 100 mL distilled water.

 

d. Nitrite-Nitrogen

 

Nitrite can be measured easily and with good sensitivity by a coupling reaction known as diazotization. Under acidic conditions, nitrite ions and aromatic amines can form reactive diazonium salts. In the test commonly employed for nitrite, the aromatic amine used is sulfanilamide. This forms p-diazobenzenesulfonamide (equation 17.31).

 

(17.31)

 

To improve the intensity of the absorbance band, this is reacted with N-(1-naphthyl)ethylenediamine to form a reddish purple azo dye, p-benzenesulfonamide azonaphthylethylenediamine (equation 17.32).

 

(17.32)

 

Absorbance of this azo dye is measured at 543 nm. It obeys Beer's law up to 180 g-N/L. This method requires the use of nitrite-free dilution water and accurate nitrite standards. If its not certain that distilled water is nitrite-free, it can be treated with a small amound of potassium permanganate. This common oxidizing agent converts nitrite to nitrate. Residual permanganate can be removed by fractional distillation of the water, discarding any pink distillate. Because nitrite is so easily oxidized, even fresh solutions must be standardized. This is done by oxidation with a standard permanganate solution and back titration with a standard reducing agent (oxalate or ferrous ammonium sulfate).

 

Procedure

1. To 50 mL of sample in a 100 mL beaker, add 2 mL of the color reagent and mix..

2. Wait 10 min after addition of the reagents, then measure the absorbance at 543 nm. If the absorbance is greater than that measured for the 0.025 mg/L standard, dilute the original sample and repeat steps 1-2.

3. Repeat steps 1-2 for all samples along with three standard solutions prepared from the 250 mg/L nitrite-N solution.

 

Reagents

1. Color Reagent: Place about 800 mL of distilled water to a 1-liter volumetric flask. To this slowly add 100 mL of 85% phosphoric acid and then 10 g sulfanilamide. Once the sulfanilamide is completely dissolved, add 1 g N-(1-naphthyl)-ethylenediamine dihydrochloride and dissolve. Dilute to 1 liter. This reagent must be stored in a dark bottle in a refrigerator. It can be used for 1 month.

2. Nitrite Standard (250 mg/L as N) Dissolve 1.232 g NaNO2 in 1 liter of distilled water. Add 1 mL chloroform as a preservative. Standardize with permanganate and oxalate or ferrous ammonium sulfate.

 

 

 

 

Figure 17.3

Nitrite Standard Curve

 

 

 

Figure 17.4

Determination of Nitrite in an Unknown from Several Dilutions


e. Ozone: Gas phase

 

Ozone is powerful oxidant that is widely used in Europe for the treatment of drinking water. Its use in the U.S. is less common, but it is a rapidly growing technology. Principal applications of ozone include disinfection, oxidation of Fe and Mn, oxidation of industrial pollutants, bleaching of color, and improvement of coagulation/filtration.

 

1. Direct UV Absorption

 

Commercial gas phase ozone monitors are based on the direct measurement of ultraviolet absorbance. With bench-scale studies, it is often convenient to use a laboratory UV-Vis spectrophotometer equiped with a flow-through quartz cell (0.1-0.2 cm pathlength) as a substitute for a dedicated ozone gas monitor. The ozone concentration may be calculated based on Beer's Law and the Ideal Gas Law.

 

(17.33)

 

where o is the absorptivity in atm-1cm-1 of ozone at the wavelength of measurement, Ta is the absolute temperature of the gas being measured, P is the pressure in mm Hg of the gas, and L is the cell pathlength in cm. At 253.7 nm, the absorptivity of ozone in the gas phase (at 760 torr, 293oK or 20oC) is 134 atm-1cm-1 (Inn & Tanaka, 1959; Hearn, 1961; DeMore & Patapoff, 1976). There is quite a bit of fine structure to ozone's broad UV absorbance band in the gas phase. For this reason, narrow band widths are preferable. When expressed as mass per volume, the terms for pressure and temperature drop out and one gets equation 17.34a.

 

(17.34a)

 

which for a wavelength of 253.7 nm and a pathlength of 0.2 cm reduces to equation 17.34b.

 

(17.34b)

 

Use of equations 17.32 and 17.33 is more convenient in the laboratory. If necessary, one can convert back to percent-based concentrations by equation 17.34c.

 

(17.34c)

 

 

 

 

 

2. Iodometric Methods

 

Gas phase concentrations of ozone are also easily measured iodometrically. Wet chemical analysis of ozone in the gas phase is less convenient, however, it allows one to calibrate or check spectrophotometric analyzers. Also, the wet chemical tests measure mass flow rather than concentration directly. If there are uncertainties in the gas flow rate, the wet chemical methods may provide greater accuracy in estimating mass ozone application rates.

Many variations on the classic iodometric method have been proposed, Basically, a portion of the gas stream is directed to a gas bubbler filled with 2% KI solution for an exact period of time. Ozone reacts stoichiometrically to form an equivalent amount of iodine. In the presence of excess iodide, the triiodide ion (I3) is formed.

 

O3 + 2KI + H2O ---------> I2 + O2 + 2OH- + 2K+ (17.35)

 

I2 + I- <---------> I3- (17.36)

 

 

The iodine formed is then titrated with sodium thiosulfate using starch as an indicator to accentuate the endpoint (APHA et al., 1985). Phosphate buffered KI solutions are to be avoided, as the phosphate anions appear to catalyze the formation of hydrogen peroxide (Flamm, 1977). This can lead to a shifting endpoint and unpredictable stoichiometry. Flamm (1977) reports that a borate buffered modification of the standard KI procedure has shown excellent agreement with direct gas-phase UV absorption. He measured the triiodide by spectrophotometry (352 nm). Calibration of this method requires the spectrophotometric analysis of standard triiodide solutions. These solutions are prepared by oxidation of iodide using standard iodate solutions (equation 17.37). The stoichiometry of this reaction is preserved when conducted in 0.1 to 1.0 N acid. The direct preparation of triiodide from iodine crystals and iodide is less accurate.

 

(17.37)

 

From equations 17.35 and 17.37, one concludes that one mole of iodate liberates the same amount of iodine/triiodide as three moles of ozone. Since iodine or triiodide has two equivalents of oxidizing potential, ozone also has two, and iodate has six.


 

a. Procedure

 

1. Bubble ozone gas through a gas washing bottle containing a convenient volume of BKI solution (usually 250-500 mL). Record the bubbling time and gas flow rate or settings. A sample of the original (time zero) BKI solution should be saved for determination of UV blank. For low-level measurements, it is recommended that the BKI solution be very slightly preozonated. This removes small amounts or reducing materials that are invariably present in commercial potassium iodide.

2. At the end of the ozone trapping period, disconnect the gas washing bottle, and pour a small sample (e.g., 10 mL) into a 50 mL beaker.

3. Measure absorbance of the ozonated BKI at 352 nm after 1 minute (Abss). If the absorbance is greater than 1.2, the sample must be diluted. The final absorbance value must then be corrected for this dilution.

4. Obtain a blank measurement by determining the absorbance at 352 nm of an aliquot of the same BKI solution that was present in the gas washing bottle at time zero (Absb).

5. Calculate concentration with equation #8. Obtain the calibration factor, 1, by the standardization procedure in "b".

 

Ozone Concentration (mg/L as O3) = (Abss-Absb)/1 (17.38)

 

b. Standardization

 

1. Prepare a set of standard triiodide solutions. This is done by first cleaning a series of 100 mL volumetric flasks, and half filling them with the Acidic KI solution. Add aliquots ("V" mL) of the Standard Potassium Iodate Solution to each and fill to the mark with additional acidic KI solution. Each mL of iodate will be roughly equivalent to 1.5 mg/L ozone.

 

2. Measure the absorbance of each of the standard triiodide solutions at 352 nm. Plot absorbance vs equivalent ozone concentration.

 

Abs = 1 (Equiv. Conc.) (17.39)

 

where the equivalent concentration is given by:

 

Equiv. Conc. (mg/L as O3) = MIO3 (100) (3Mole IO3) (48,Mole O3O3) (17.40)

 

= 1,440 MIO3 V

 

Note that MIO3 is the exact molar concentration of the Standard Potassium Iodate Solution (see equation 11). Try to work only in the linear range, or the range from 0-1.2 absorbance units.

 

 

c. Reagents

 

1. BKI Reagent (0.1 M Boric Acid, 1% Potassium Iodide): Add 6.2 g H3BO3 and 10.0 g KI to 1 liter of distilled water. Stir to dissolve.

 

2. Standard Potassium Iodate Solution: Dry the primary standard at 120oC for 2 hours. Then, weigh out about 0.021 g of the dried material. Record the exact weight to 4 significant figures. Dissolve in distilled water in a 100 mL volumetric flask and fill to the mark. Calculate exact molar concentration:

 

MIO3 = wt in grams / 21.402 (17.41)

 

3. Acidic KI Solution (1% potassium iodide in 0.1N acid): Add 5.6 mL of concentrated H2SO4 to a 1 liter volumetric flask. Slowly fill the flask about half-way with distilled water and stir. Then add 10 g KI, stir, and fill to the mark with distilled water.

 

 

Strictly speaking, all of these iodometric methods are non-selective. That is, they measure a wide range of oxidants, not just ozone. For this reason, they should not be used to measure aqueous ozone concentrations. Because ozone is by far the major oxidant species produced by corona ozone generators, and because ozone is far more easily stripped from water, the measurement of ozone in the gas phase is not subject to significant problems with interferences. Thus, iodometric methods may be used in this case without reservation.

 

 

 

e. Ozone: Aqueous Phase

 

1. Direct UV Method

 

Aqueous ozone concentrations in pure (e.g., distilled) water may be conveniently determined by direct spectrophotometric measurement at 260 nm.

 

CO3 (mg/L as O3) = 14.59*(Abs @260 nm) (17.42)

 

Equation 12 is based on a molar absorptivity of 3290 M-1cm-1 (Hart et al., 1983). Unfortunately, most solutes will interfere at this wavelength, so with actual environmental samples another method must be used. Since the iodometric method is too nonspecific for aqueous determinations, the indigo method of Bader and Hoigne (1981) is recommended.

 

2. Indigo Method

 

This colorimetric procedure uses solutions of indigo trisulfonate (Bader & Hoigne, 1981). Ozone will stoichiometrically bleach this intense blue dye, and the loss in absorbance at 600 nm may be translated directly into an ozone concentration. The reaction product is relatively unreactive to further ozonation. The reaction is best carried out at low pH to minimize ozone decomposition, and preserve the 1:1 stoichiometry. Bader and Hoigne (1981) report a sensitivity factor or apparent absorptivity for indigo trisulfonate of 20,000 M-1cm-1. This is based on an aqueous ozone molar absorptivity of 2900 M-1cm-1. If one adopts the higher value reported by Hart et al. (i.e., 3290 M-1cm-1), the sensitivity factor for indigo trisulfonate becomes 22,700 M-1cm-1.

This method is quite selective, however, it is subject to a few notable interferences. The presence of residual chlorine will cause a positive bias. Addition of 500 mg/L malonic acid to the Indigo Reagent solves this problem by out-competing the indigo for the chlorine. Oxidized manganese species will also result in bleaching of the indigo. Here it is recommended that duplicate samples be analyzed, one according to the standard procedure, and one following addition of glycine. The glycine selectively reduces residual ozone without affecting oxidized manganese species. The true ozone concentration may then be estimated from the difference of these two measurements.

 

a. Procedure

1. Prepare an indigo blank by adding 1 mL of the Standard Indigo Stock to a 25 mL volumetric flask and filling to the mark with Phosphate Buffer. Stopper and mix. Measure the absorbance of this solution at 600 nm (Absi). When the Indigo Stock is new, it should be about 0.650. With time this value will drop. When it falls below 80% of the original value, prepare a new Indigo Stock and repeat procedure. If low ozone concentrations are anticipated (i.e., < 0.3 mg/L) prepare a Secondary Indigo Stock by adding 20 mL of the Standard Indigo Stock to a 100 mL volumetric flask and diluting to the mark with Super-Q water.

2. Soak a series of volumetric pipets in a dilute ozone solution for several minutes. These pipets are to be used to transfer the solution to be measured to the indigo-containing flasks. They must therefore be rendered ozone-demand-free. The capacities of the pipets will depend on the range of anticipated ozone concentrations (see table below).

3. Assemble a series of clean, glass-stoppered 25 mL or 50 mL volumetric flasks (see table below) and fill each with 1.00 mL of Standard Indigo Stock (25 mL flasks) or Secondary Indigo Stock (50 mL flasks) using a volumetric pipet. Wash this down from the inner surfaces of the neck with about 10 mL of Phosphate Buffer.

4. Quickly pipet the recommended sample volume to an indigo-containing volumetric flask (see table below). Be sure that the tip of the pipet is below the meniscus. Fill the flask to the mark with Phosphate Buffer, cap and invert several times to mix.

 

(Vs)

(Vt)

(L)

Anticipated Ozone

Recommended

Recommended

Recommended

Concentration

Sample Volume

Total Volume

Pathlength

0 - 0.2 mg/L

25 mL

50 mL *

10 cm

0.1 - 0.3 mg/L

15 mL

50 mL *

10 cm

0.2 - 0.5 mg/L

10 mL

50 mL *

10 cm

0.3 - 2.0 mg/L

15 mL

25 mL

1 cm

1.5 - 3.0 mg/L

10 mL

25 mL

1 cm

1.8 - 3.5 mg/L

8 or 9 mL

25 mL

1 cm

2.3 - 4.5 mg/L

6 or 7 mL

25 mL

1 cm

3 - 6 mg/L

5 mL

25 mL

1 cm

4 - 7 mg/L

4 mL

25 mL

1 cm

5 - 10 mg/L

3 mL

25 mL

1 cm

7 - 15 mg/L

2 mL

25 mL

1 cm

15 - 30 mg/L

1 mL

25 mL

1 cm

*Secondary Indigo Stock must be used with 10 cm pathlength cells

 

 

5. Measure absorbance (Absf) of each sample at 600 nm using cells of the indicated pathlength (L). Concentration is calculated from a slope or calibration factor determined by calibration against the direct UV method (see b. Calibration).

 

Ozone Conc. (mg/L as O3) = Vt(Absi - Absf)/b1VsL (17.43)

 

b. Calibration

1. Prepare a series of aqueous ozone standards in slightly acidified (0.1 mM HNO3) Super-Q water. This is generally done by first bubbling ozone gas through about 500-1000 mL of the acidified water until near saturation (~1 hour). Under optimal conditions for high ozone output (i.e., high voltage, low gas flow, low temperature) this should result in aqueous concentrations of about 10 mg/L. Then, aliquots of this water are removed and diluted with varying amounts of un-ozonated, acidified Super-Q water. The degree of dilution will depend on the range of ozone concentrations anticipated in the samples of interest.

2. One-by-one measure each of these diluted solutions by the indigo method above and by the direct UV method (equation 17.42). Use the same sample volume, Vs, for all standards.

3. Plot absorbance (from indigo method) versus aqueous ozone concentration (from direct UV method). The slope of this line multiplied by Vt/Vs gives b1L (see equation 17.44). Based on the presumed sensitivity factor for indigo trisulfonate of 22,700 M-1cm-1, the calibration factor, b1, should be about 0.47 abs/cm per mg-O3/L.

 

Absf = Absi - (b1LVs/Vt) (Ozone Conc.) (17.44)

 

c. Reagents

1. Standard Indigo Stock (1 mM in 20 mM phosphoric acid): Dissolve 1.36 mL conc. H3PO4 in 1 liter of super-Q water and mix. To this add 0.6 g indigo trisulfonate, mix and store in a brown glass bottle.

2. Phosphate Buffer (pH 2): Dissolve 28 g NaH2PO4.H2O and 20.6 mL (35 g) conc. H3PO4 in Super-Q water and dilute to 1 liter.

 

 

f. Aqueous Total Oxidant Concentration

 

Ozone reacts in aqueous solution to give a variety of oxidant species. These may include simple inorganic oxygen-containing free radicals, hydrogen peroxide, organic free radicals, and organic peroxides. The iodometric method of Flamm (1977) coupled with the use of a catalyst (Taube & Bray, 1940) is sufficiently non-specific to be useful for the determination of total oxidant concentration.

 

 

a. Procedure

 

1. Place between 1 and 9 mL of BKI reagent in a 10 mL volumetric flask.

2. Add 1 mL of the Molybdate Catalyst solution.

3. Quickly add sample to the mark, stopper and invert several times to mix.

4. Measure absorbance at 352 nm after 1 minute (Abss)

5. Repeat steps 1-4 substituting distilled water for the sample (Absb)

6. Calculate total oxidant concentration:

 

Total Oxidant Concentration (mg/L as O3) = b1*(Abss-Absb)/Vs (17.45)

 

Follow the procedure in part "b" using iodate standards to get the calibration factor, b1. Experience indicates that it should be about 21.6 mg/L as O3 per mL of sample per cm pathlength.

 

b. Standardization

1. Prepare a set of standard triiodide solutions. Clean a series of 100 mL volumetric flasks, and half fill them with the Acidic KI solution. Add aliquots ("V" mL) of the Standard Potassium Iodate Solution to each and fill to the mark with additional acidic KI solution. Each mL of iodate will be roughly equivalent to 1.5 mg/L ozone.

2. Measure the absorbance of each of the standard triiodide solutions at 352 nm. Plot absorbance vs equivalent ozone concentration.

 

Equiv. Conc. (mg/L as O3) = MIO3 (100) (3Mole IO3) (48,Mole O3O3) (17.46)

 

= 1,440 MIO3 V

 

Try to work only in the linear range, or the range from 0-1.2 absorbance units.

 

c. Reagents

1. BKI Reagent (0.1 M Boric Acid, 1% Potassium Iodide): Add 6.2 g H3BO3 and 10.0 g KI to 1 liter of distilled water. Stir to dissolve.

2. Molybdate Catalyst Solution: Add 0.9 g Ammonium Molybdate to 10 mL distilled water and stir to dissolve.

3. Standard Potassium Iodate Solution: Dry the primary standard at 120oC for 2 hours. Then, weigh out about 0.021 g of the dried material. Record the exact weight to 4 significant figures. Dissolve in distilled water in a 100 mL volumetric flask and fill to the mark. Calculate exact molar concentration:

 

MIO3 = wt in grams / 21.402 (17.47)

 

4. Acidic KI Solution (1% potassium iodide in 0.1N acid): Add 5.6 mL of concentrated H2SO4 to a 1 liter volumetric flask. Slowly fill the flask about half-way with distilled water and stir. Then add 10 g KI, stir, and fill to the mark with distilled water.

 

g. Hydrogen Peroxide

 

Hydrogen peroxide is one of the secondary oxidants produced by the action of ozone. It may be produced as a byproduct of aqueous ozone decomposition. It is also an important byproduct of the reaction of ozone with unsaturated organic compounds. Hydrogen peroxide is also important, because it may be added to waters undergoing ozonation for the purposes of accelerating the formation of hydroxyl radicals. These radical species are capable of oxidizing many types of organic structures that are not affected by molecular ozone. The two preferred methods for hydrogen peroxide analysis involve the formation of heavy metal complexes (Parker, 1928; Masschelein et al., 1977). Both are susceptible to interference from aqueous ozone. Therefore, ozone must first be purged. The titanium method (Parker, 1928) is recommended here.

 

a. Procedure

1. Sparge sample to remove residual ozone as necessary.

2. Pipet 1 mL of the Titanium Reagent into a 10 mL volumetric flask.

3. Add sample to the mark, stopper and invert several times.

4. Measure absorbance at 415 nm.

5. Calculate concentration from a calibration curve prepared as below.

 

b. Standardization

1. Prepare two serial dilutions of the commercial 30% solution. Add 1 mL of the commercial solution to a 500 mL volumetric flask and dilute to the mark with distilled water. This is the concentrated peroxide stock solution. In a second 500 mL volumetric flask dilute 5 mL of the concentrated stock to 500 mL with distilled water. This is the dilute peroxide stock.

2. Standardize the dilute peroxide stock using the BKI spectrophotometric method for total oxidants. In this case the standard curve is most conveniently based on equivalent hydrogen peroxide concentrations in mg/L.

 

Equiv. Conc. (mg/L as H2O2) = MIO3 (100)(3 Mole IO32)(34,Mole H2O2O2) (17.48)

 

= 1,020 MIO3 V

or:

 

Equiv. Conc. (mg/L as H2O2) = 0.708 Equiv. Conc. (mg/L as O3) (17.49)

 

3. Into a series of 10 mL volumetric flasks, place 1 mL of the Titanium Reagent. Add varying amounts of the dilute peroxide standard, and fill to the mark with distilled water. Measure the absorbance at 415 nm. Plot concentration versus absorbance. The data should be linear with a slope of about 57 mg/L per absorbance unit per cm.

 

 

c. Reagents

1. Titanium (IV) Reagent: Place 10 mL of 6N HCl in a beaker on a crushed ice bath under a fume hood. Place TiCl4 in an ice bath and allow both to cool. Carefully add 10 mL of the TiCl4 dropwise to the chilled HCl. Allow the mixture to stand in the ice bath until the solids that form dissolve. Transfer this mixture to a 1 liter volumetric flask using 6N HCl. Dilute to the mark with additional HCl.

2. 30% Hydrogen Peroxide Solution (Commercially available).

3. BKI Reagent: (This is the same solution prepared for ozone gas phase analysis and the total oxidant analysis; 0.1 M Boric Acid, 1% Potassium Iodide): Add 6.2 g H3BO3 and 10.0 g KI to 1 liter of distilled water. Stir to dissolve.

4. Molybdate Catalyst Solution: (This is the same solution prepared for the total oxidant analysis) Add 0.9 g Ammonium Molybdate to 10 mL distilled water and stir to dissolve.

5. Standard Potassium Iodate Solution: (This is the same solution prepared for ozone gas phase analysis and the total oxidant analysis) Dry the primary standard at 120oC for 2 hours. Then, weigh out about 0.021 g of the dried material. Record the exact weight to 4 significant figures. Dissolve in distilled water in a 100 mL volumetric flask and fill to the mark. Calculate exact molar concentration:

 

MIO3 = wt in grams / 21.402 (17.50)

 

6. Acidic KI Solution (1% potassium iodide in 0.1N acid; This is the same solution prepared for the ozone gas phase analysis and the total oxidant analysis): Add 5.6 mL of concentrated H2SO4 to a 1 liter volumetric flask. Slowly fill the flask about half-way with distilled water and stir. Then add 10 g KI, stir, and fill to the mark with distilled water.

 

 

 

 

 

REFERENCES

 

Bader, H.; Hoigne, J. Wat. Res. 1981, 15:449-456.

 

Bolleter, W.T., C.J. Bushman and P.W. Tidwell (1961) "Spectrophotometric Determination of Ammonia as Indophenol," Anal. Chem. 33(4)592-594.

 

Davenport, W.H. (1949) Anal. Chem. 21.

 

DeMore, W.B.; Patapoff. M. Environ. Sci. Technol. 1976, 10:897.

 

Flamm, D.L. Environ. Sci. Technol. 1977, 11:10:978-983.

 

Hart, E.J.; Sehested, K.; Holcman, K. Anal. Chem. 1983, 55:46.

 

Hearn, A.G. Proc. Phys. Soc. 1961, 78:932.

 

Inn, E.C.Y.; Tanaka, Y. In Ozone Chemistry and Technology 1959, ACS Advances in Chemistry Series #21, ACS, Washington, pp.263-268.

 

Masschelein, W.J.; Denis, M.; Ledent, R. Wat. Sewage. Wrks. 1977, August, 69-72.

 

Parker, G.A. In Colorimetric Determination of Nonmetals, 1928, Boltz, D.R.; Howell, J.A., Eds. Wiley, New York, pp.301-303.

 

Taube, H., Bray, W.C. J. Am. Chem. Soc. 1940, 62:3357.

 

 

 


 

 

 

 

 

 

new stuff

 

 

III. MEASUREMENT OF OZONE

 

A. Gas Phase Concentrations

 

Commercial gas phase ozone monitors are based on the direct measurement of ultraviolet absorbance. With bench-scale studies, it is often convenient to use a laboratory UV-Vis spectrophotometer equiped with a flow-through quartz cell (0.1-0.2 cm pathlength) as a substitute for a dedicated ozone gas monitor. Gas phase concentrations of ozone are also easily measured iodometrically. A portion of the gas stream is directed to a gas bubbler filled with 2% KI solution for an exact period of time. Ozone reacts stoichiometrically to form an equivalent amount of iodine. In the presence of excess iodide, the triiodide ion (I3) is formed.

 

O3 + 2KI + H2O ---------> I2 + O2 + 2OH- + 2K+ (20)

 

The iodine formed is then titrated with sodium thiosulfate using starch as an indicator to accentuate the endpoint (APHA et al., 1985). Flamm (1977) reports that a borate buffered modification of the standard KI procedure has shown excellent agreement with direct gas-phase UV absorption. He measured the triiodode (iodine + iodide) formed by spectrophotometry (352 nm).

 

B. Aqueous Concentrations

 

Aqueous ozone concentrations in pure (e.g., distilled) water may be conveniently determined by direct spectrophotometric measurement at 260 nm.

 

CO3 (mg/L) = 14.59*(Abs @260 nm) (21)

 

Equation 47 is based on a molar absorptivity of 3290 M-1cm-1 (Hart et al., 1983). Unfortunately, most solutes will interfere at this wavelength, so with actual drinking waters another method must be used. Since the iodometric method is too nonspecific for aqueous determinations, the indigo method of Bader and Hoigne (1981) is recommended. This colorimetric procedure uses solutions of indigo trisulfonate. Ozone will stoichiometrically bleach this intense blue dye, and the loss in absorbance at 600 nm may be translated directly into an ozone concentration.

 

 

old stuff:

 

 

******************************** Research Topic *****************************

 

a. UV Absorbance

 

Very little structural information may be obtained from the ultraviolet-visible absorption spectrum of humic substances. Figure 17.1 shows a typical scan of an aquatic fulvic acid. Note the total lack of features. One sees only a monotonic decrease in absorbance with increasing wavelength. Nevertheless, the absolute value of the absorbance has been found to correlate with certain water quality characteristics, such as DOC and THMFP, and this is often the reason for measuring UV absorbance.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Figure 17.1

UV-Visible Absorbance Scan of an Aquatic Fulvic Acid

 

 

It is now reasonably well established that the absorbance of light in the UV-Visible range by aquatic humic substances conforms to Beer's law (Black & Christman, 1963a,b, Packham, 1964). That is, the concentration of humic materials in a natural water which has been diluted with distilled water is linearly proportional to its absorbance. Therefore, one can calculate a specific absorbance which corresponds to the absorbance per mg/l of DOC. Some specific absorbances for extracted humic materials are listed in Table 17.6. As a result of the conformance to Beer's law the DOC of a raw water can be estimated by measuring its absorbance at a fixed wavelength provided that the relative concentration of all the constituent organics remains fixed. Although this is never strictly true, changes in the nature of the raw water may be small enough to be unimportant. Table 17.7 summarizes some attempts to correlate UV absorbance (254 nm) to DOC for raw waters. Note that the reciprocal of the slopes in Table 17.7 correspond to the specific absorbances in Table 17.6, and that the average slope (~25) gives a specific absorbance of 0.004 which is similar to those reported for riverine fulvic acids.

 

Note that groundwater humics and humics from eutrophic lakes are less colored than riverine humics. Humics from bogs and swamps are most colored.

 

Since color increases with increasing pH, it is important to make all color measurements at the same pH. By convention, a pH 7 phosphate buffer is generally used.

 

 

Table 17.6

Specific Absorbance of Aquatic Humic Substances

 

 

Sample Specific Color (400nm) Specific UV abs (254nm) Ref.

Average Range Average Range

Groundwater

Fulvic Acid 0.001 1

Humic Acid 0.002 1

Rivers & Streams

Fulvic Acid 0.005 1

0.0033 0.0023-0.0047 0.035 0.029-0.043 2

Humic Acid 0.007 1

0.0081 0.0071-0.0089 0.054 0.049-0.059 2

Wetlands

Fulvic Acid 0.010 1

Humic Acid 0.014 1

 

 

References

1. Thurman, 1985

2. Reckhow, 1988

 

 

 

Table 17.7

UV Absorbance (254 nm) to DOC Regression Equations+

 

Water Slope Intercept r Reference

Tjeukemeer Lake 24.7* 2.7 0.925 De Haan et al., 1982

Grasse River 19.8 0.65 0.93 Edzwald et al., 1985

Glenmore Reservoir 28.1 -0.3 0.71 Edzwald et al., 1985

24.1 2.2 0.96 Smart et al., 1976

N.C. Surface Wat. 29.1 1.44 0.83 Singer et al., 1982

 

 

+Model: DOC (mg/l) = (slope)*(UV absorbance, cm ) + intercept

* Based on absorbance at 250 nm, neutral pH.

 

 

 

Estimation of THMFP Trihalomethane formation potential (THMFP) is the trihalomethane formation that occurs under a specified set of extreme conditions. Reaction conditions are chosen so that THMFP approaches the maximum possible yield. A number of researchers have proposed the use of simple linear models for THMFP, based on a single easily-measured property of the water (equation 7).

 

THMFP = b0 + b1(independent variable) (17.12)

 

Ultraviolet absorbance of a water has also been used as a surrogate for THMFP. Its utility is based on the presumption that conjugated and electron-rich structures that are responsible for much of the absorption of 254 nm light, are also readily attacked by chemical electrophiles such as chlorine. A certain fraction of these reactions (e.g., ~3%) will ultimately lead to the formation of THMs. UV absorbance is generally determined using light of 254nm wavelength, and it should be measured only on samples that have been previously filtered. Table 17.8 shows results of some studies where UV absorbance measurements were made on filtered samples.

 

 

 

Table 17.8

Correlations Between UV Absorbance* and THMFP*

 

 

Sample Type b0 b1 r2 Origin Ref

Humic Substances 1570-150 Aq. Fulvic Acids 4

1320-220 Aq. Humic Acids 4

Raw Waters 102 2106 0.86 N. Carolina Waters 1

38 1573 0.94 Grasse River 2

49 1596 0.82 Glenmore Res. 2

101-74 2120-180 0.80 Colored Waters 3

-62 1975 0.741 Colored Waters 8

Treated Waters 38 1629 0.97 Canton WTP 2

42 1578 0.94 Oneida WTP 2

 

 

*THMFP is in units of ug/l, UV Absorbance is in abs. units/cm.

 

 

References

1. Singer et al. (1981); THMFP conditions - 7 days, pH 6.7, sufficient chlorine to maintain a residual of 15-20 mg/l; Chlorine demand - 7 days

2. Edzwald et al. (1985); THMFP conditions - 7 days, pH 7.5, 3-5 mg chlorine per mg TOC; Chlorine demand - 2 hours, 10-35 mg/l dose.

3. Singer (1987); THMFP conditions - 7 days, pH 7, 2-3 mg chlorine per mg TOC.

4. Reckhow (1984); THMFP conditions - 3 days, pH 7, 20 mg/l chlorine, 20oC; values adjusted to 7-day reaction time.

5. Urano et al. (1983); General THM Model adapted to the following conditions: 7 days, pH 7, 20 mg/l chlorine dose, 20oC

6. Morrow & Minear (1987); General THM Model adapted to the following conditions: 7 days, pH 7, 20 mg/l chlorine dose, 20oC

7. Oliver & Thurman (1983); THMFP conditions - 7 days, pH 7, 10 mg/l chlorine, 20oC.

8. Amy et al., 1987); THMFP conditions - 3 days, pH 7, chlorine to carbon mass ratio of 3.0, 20oC; General THM Model adapted to the following conditions: 7 days, pH 7, 20 mg/l chlorine dose, 20oC

 

 

 

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[1]The old mercury vapor lamps, once common in spectrophotometers, had an especially high intensity at 254 nm. Therefore, measurements at this wavelength were favored, because they were especially free of interferences posed by stray light, etc.