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CE142LabManual.pdf
California State University Fresno CE 142L Environmental Quality Laboratory
Laboratory Manual
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PRELUDE: WHAT IS ENVIRONMENTAL CHEMISTRY
INTRODUCTION:
What is environmental chemistry? This question is a little difficult to answer because environmental chemistry encompasses many different topics. Some define it as follows:
“Environmental chemistry is the study of the sources, reactions, transport, effects, and fates of chemical species in water, soil, and air environments." (Stanley E. Manahan. 1991. Environmental Chemistry, 5th ed.). "(The) central position of aquatic chemistry in the natural sciences gives it an increasing popularity in science and engineering curricula; it also makes it a difficult topic to teach for it requires exploring some aspects of almost all sciences." (Francois M. M. Morel. 1983. Preface to Principles of Aquatic Chemistry).
Basically, Environmental Chemistry is the use of chemistry to understand the interactions of environmental systems. Water chemistry is an important aspect of Environmental Chemistry.
A fundamental tool in analyzing water chemistry is total dissolved solids (TDS). The TDS in water consists of dissolved inorganic salts and organic materials. In natural waters, salts are chemical compounds comprised of anions (-) such as carbonates, chlorides, sulfates, and nitrates (primarily in ground water), and cations (+) such as potassium (K), magnesium (Mg), calcium (Ca), and sodium (Na) (EPA, 1986). In ambient conditions, these compounds are present in proportions that create a charge- balanced solution. If there are additional inputs of dissolved solids to the system, the balance is altered and the solution will adjust to achieve charge balance.
This lab manual includes exercises in water chemistry calculations in order to better understand chemical reactions within the aquatic environment. A fundamental understanding of water chemistry is necessary for the remaining laboratory experiments and, later on, for professional practice in civil engineering.
PREPARATION BEFORE ARRIVING AT LAB: 1. The knowledge provided in high school chemistry courses and in CHEM 1A and 3A, while
important, is not adequate for this course or for CE 142 lecture. In view of this, all students in both courses are expected to devote a considerable amount of time outside of class expanding their knowledge of the chemical concepts contained in the lab manual for this course and in the textbook adopted for CE 142.
2. Students are expected to arrive to lab with a complete understanding of the topic assigned for the week. Therefore, students are to (a) study in advance of the lab all relevant material in the lab manual pertaining to the topics and (b) conduct additional reading on the topics from the recommended textbook and other sources (books, Web, etc..). An excellent source of knowledge can be found in YouTube videos.
3. Students are expected to print and bring to class hard copies of each lab exercise, or to have ready access to the material on their laptop.
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Contents LAB 1: WATER CHEMISTRY CALCULATIONS ....................................................... 4 LAB 2: GAS TRANSFER .............................................................................................. 5 LAB 3: BIOCHEMICAL OXYGEN DEMAND .......................................................... 12 LAB 4: ADSORPTION................................................................................................ 21 LAB 5: SOLIDS IN WATER ....................................................................................... 24 LAB 6: CHEMICAL COAGULATION JAR TESTING .............................................. 30 LAB 7: CHEMICAL EQUILIBRIUM, BUFFERING, AND ALKALINITY ............... 36 LAB 8: HARDNESS .................................................................................................... 44 LAB 9: CADILLAC DESERT: WATER & TRANSFORMATION OF NATURE ...... 48
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LAB 1: WATER CHEMISTRY CALCULATIONS Instructions: Do your calculations on engineering paper and then transfer the answers to the spaces indicated near the problem statement. Attach your calculations to the lab sheets when you are finished. Your calculations should be presented in a professional manner -- legible, organized, and complete so that someone else can understand them. Recommended reading from your course textbook (Sawyer, C.N., P.L. McCarty, and G.F. Perkin, Chemistry for Environmental Engineering, 4th ed., McGraw Hill, New York, 1994.): Chapter 2 1. (15 pts.) Molecular weights -- Calculate molecular weights of the compounds:
a. Sodium bicarbonate (baking soda) -- NaHCO3 Answer: b. Ferric Chloride -- FeCl3 Answer: c. Ammonium sulfate -- (NH4)2SO4 Answer:
2. Molar and mass concentrations a. (10 pts.) If the molar concentration of sodium bicarbonate (NaHCO3) is 2.00 mM (millimolar),
what is the mass concentration of NaHCO3 (in mg/L)? For the same molar concentration, what is the mass concentration of Na+? Answer: Answer:
b. (15 pts.) What are the molar concentrations of iron (Fe3+) and chloride (Cl-) in a solution with a mass concentration of 40.7 mg/L FeCl3? Answer:
3. Mass loading and concentration calculations a. (10 pts.) What mass of ammonium sulfate (NH4)2SO4 must you add to water to make a 1.50 L
solution that is 0.045 M? Answer:
b. (10 pts.) Suppose 1.50 L of a solution with a mass concentration of 45 mg/L of some salt is left uncovered in a warm dry room. If, after a week, the volume has dropped to 0.45 L due to evaporation, what is the final concentration? Answer:
c. (10 pts.) A treatment plant discharges a wastewater flow to a lake. If the flow is 3.550 million gallons per day and the concentration is 25.00 mg/L of Nitrate (NO3), what is the annual mass loading of total Nitrogen (N) (in lb/yr). Answer:
4. Ideal Gas Law (15 pts): How many pounds of Argon are there in a 3.880 cubic foot tank if the pressure is 18.00 atm, the temperature is 20.00 �C, and the Argon purity is 99.0%?
Answer: 5. Adsorption (15 pts): In an adsorption batch test, substance Y is dissolved in the solution with an initial
concentration of 370 mg/L. 0.450 gram of granulate activated carbon (GAC) was added to 100 mL of such solution and had sufficient contact time with the solution. The substance Y solution now has a concentration of 12.0 mg/L. What is the mass of substance Y that is adsorbed to the GAC? What is mass to mass ratio of the adsorbed substance Y to GAC in mg/kg at this equilibrium concentration?
Answer:
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LAB 2: GAS TRANSFER
INTRODUCTION Gas transfer is defined as the process of allowing any gas to dissolve into a fluid, or, the process of promoting the release of a dissolved gas from a fluid. For our purposes the gas of interest is oxygen. Gas transfer of oxygen to water plays an important role in water and wastewater treatment through a process known as aeration. In wastewater treatment, aeration has several functions, which are as follows:
· Prevents the formation of hydrogen sulfide, an odor causing substance. · Stripping of volatile organics and inorganics from the water. · Oxygen supply for aerobic biological treatment of activated sludge. This is also needed for a
process known as nitrification in which bacteria convert ammonia nitrogen into Nitrate NO3-
(Vesilind 1997). · Raising the dissolved oxygen (D.O.) levels in treated effluent to waterways to protect aquatic life. · Increasing oxygen concentration in receiving waterways for emergency situations, to restore
aquatic life in waterways that have been contaminated. In water treatment, Aeration also aids in the following:
· Remove taste and odor causing hydrogen sulfide. · Removal of excess carbon dioxide in groundwater. · Oxidizing ions such as iron and manganese (Schroeder 1987).
Diffusers and mixers, large scale versions of those used in this lab, are often utilized in the Aeration process.
The objective of this lab is to observe and measure gas transfer rates in three different types of systems with the aid of a Dissolved Oxygen meter. The three systems observed are a control system, where the water is not disturbed, a system involving mixing, and a final system utilizing a fine bubble diffuser.
BACKGROUND:
Dissolved oxygen (DO) is the measure of the concentration of free (i.e. not chemically bonded) molecular oxygen, usually measured in mg/L or percent saturation. Considered the most important and commonly employed measure of water quality, DO is primarily used to assess a body of water’s ability to support aquatic life. Adequate concentrations of DO are necessary for aerobic respiration as well as the prevention of offensive odors.
The presence of DO is a result of a continuous circulation of oxygen within the environment, commonly called the oxygen cycle. This process has two primary phases, one where oxygen is produced and one where it is consumed. Oxygen is produced during photosynthesis, a process that converts carbon dioxide (CO2) to oxygen and sugar using energy from the sun. Respiration and decomposition consume oxygen in highly complex processes that covert organic matter into usable energy.
This cycle of production and consumption, for the most part, maintains DO levels throughout the environment at a constant level. On a daily basis, however, noticeable differences can occur. Because photosynthesis is dependent upon sunlight, levels of oxygen production decline during the night. While
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production declines, consumption, from respiration and decomposition, continue at a steady pace. The net result is a decline in DO levels, usually most noticeable during early twilight.
The amount of dissolved oxygen within water is controlled primarily by the concentrations of DO both within the water and within the surrounding air. Oxygen concentrations within air are usually around 21% while those in water are about 10 ppm. This differential in concentrations causes oxygen molecules at the waters surface to dissolve into the water, a result of partial pressures between the two substances. Because this process is directly proportional to the area of air in contact with the water, movement of the water greatly affects the amount of DO present. Movement of the water creates waves on the waters surface. These waves produce a net increase in the available surface area that, in turn, allows for more diffusion of oxygen into the water. The more turbulent the flow, the more surface area and the greater the concentration of DO within the water.
The amount of DO within the water may be described using Henry’s Law.
))(( CCak dt dC
SL -=
where:
C = the concentration of dissolved gas (mg/L or mole/L) CS = the concentration of gas under saturated conditions kL = gas transfer coefficient (cm/s or m/s) a = ratio of gas/liquid interfacial area to liquid volume (cm-1 or m-1) t = time
CS describes the saturation concentration which is the concentration predicted by Henry’s Law. By integrating this equation and applying the necessary boundary conditions (at t = 0 and C = C0) we are left with the exponential decay function;
)( 0 )(
atk SS
LeCCCC --=-
(CS –C), often called the “deficit” shows how the relationship between the initial concentration and the saturated concentration effects the decay function. Notice that when C is farthest away from CS, the slope of the decay curve is greatest. This implies that as the concentration increases towards saturation, the amount of dissolving decreases. Also note how the decay function is related to “a,” the ratio of surface area to volume. This relationship shows how the creation of waves affects the amount of dissolving.
Another physical attribute of water that greatly affects DO concentrations is that of temperature. Saturation, the point at which a substance has the maximum amount of another substance within in it, is greatly affected by temperature. Colder water, for instance, is capable of holding more oxygen than warmer water (i.e. warmer water becomes saturated more easily than cold water). If a body of water, say
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a lake, is warm it will be capable of holding less oxygen than if it were cold. Even if the water is 100% saturated, it will still have less DO than an equivalent lake at a lower temperature.
This affect of temperature on the levels of DO can lead to stratification of the body of water. Thermal stratification can lead to layers within the water body that are completely devoid of DO while other layers may have too much. This can be of great concern to aquatic life that may be dependent upon specific levels of DO. Too much or too little DO may result in death.
CONCLUSION Gas transfer is defined as the process of allowing any gas to dissolve into a fluid, or, the process of promoting the release of a dissolved gas from a fluid. Gas transfer of oxygen to water plays an important role in water and wastewater treatment through a process known as aeration. Therefore, the objective of this lab is to observe and measure the gas transfer rates in three different types of systems with the aid of a Dissolved Oxygen meter. The three systems observed are a control system, where the water is not disturbed, a system involving mixing, and a final system utilizing a fine bubble diffuser.
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LAB 2: PROCEDURES (Gas Transfer)
Objective: To observe and measure gas transfer rates in several systems.
Equipment and Materials: · Dissolved oxygen meter · Deoxygenated water (provided by instructor) · 1,000 mL beaker (1) · Watch with second hand
Procedures
Your instructor will brief you on the use of the dissolved oxygen (DO) meter. Before lab, your instructor de-oxygenated several beakers of tap water by storing them under a vacuum and/or reacting them with sodium sulfite (Na2SO3) and/or bubbling nitrogen gas in the water. 1. Gently pour approximately 400 mL of deoxygenated water in the beaker with the beaker at a 45
degree slope and the pour spout of the carboy resting on the inner wall of the beaker. Try not to disturb the water excessively while doing so. (Do you know why?)
2. Set up the beaker either on the counter (control), on a mixer, or on a bubbler as directed by your instructor. Assign at least two people to each beaker -- one to measure the DO and one to keep track of the time. Either person (or a third) can record the data.
3. Before starting any testing set the meter measurement mode to “Continuous.” Then measure the initial DO and the temperature using the DO meter. Wait for the reading to stabilize. Once the reading has stabilized and you have an initial DO measurement, you can start the experiment and the timer.
4. The control beaker should be started first because it takes the most time. You can run the mixed beaker test in parallel with the control, except, start it 30 minutes later. When the mixed beaker measurements are complete you can start the bubbler beaker test. When moving the probe from the mixed beaker or bubbled beaker back to the control it will take time for the probe to adjust to the lower DO condition, so you must wait for the reading to stabilize before taking/ recording the first reading.
5. At the beginning, take time and concentration readings every 1 to 5 minutes in the systems with slower transfer (i.e., the control) and every 1 to 10 seconds in the systems that you expect will have a lot of gas transfer (mixed, bubbled). After you see how fast the concentration is changing, you can adjust the time interval. The goals are to make a minimum of 15 to 20 measurements spread out over the test period for each beaker and to have each test period be long enough for the DO to approach saturation (you will not achieve this for the control, so measure that one for only 1.5 to 2 hours).
Data Analysis 1. Using the temperature read in the lab, find and record the saturation concentration from an
environmental textbook or from the Internet. 2. Record the data for all of the water samples. Continued
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3. Use a spreadsheet program and create a table with time in the first column and corresponding values of raw DO concentration in the second column. Create a plot of DO concentration against time.
4. Add a column to the spreadsheet containing deficit values (D = CS - C). · Note that the saturation concentration value (CS) is obtained from a reference book and the
value will depend only on the water temperature. · For the control beaker, omit deficit data collected in the early time period when DO values
were constant (before they began increasing). 5. Add another column containing the quantity -ln(D/D0). The initial deficit (D0) is the first value of
deficit in the table. Plot -ln(D/D0) against time (in minutes). Show your plot on a separate worksheet. Determine the transfer coefficient for each test using the slopes of your plots (See the gas transfer notes.) You can use the linear trend line function to get the best-fit regression line. Be sure to force the regression through zero. Format your graph so that the lab data are displayed as points only (no connecting line) and the linear trend line.
6. To check the reasonableness of your calculated transfer coefficient, prepare a plot with your original data points (plotted as points only) and a line (only) generated from the integrated gas transfer equation* using your calculated transfer coefficient.
7. Exchange results with the other groups. Why are the transfer rates different? Are the differences what you would expect? Explain.
* C = CS - (CS - C0)e -(KLa)t
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LAB 2: NOTES ON GAS TRANSFER KINETICS Henry's Law describes an equilibrium situation -- the rate of gas dissolving into the liquid from the atmosphere (rd) equals the rate at which the gas volatilizes from the liquid (rv). Under equilibrium, rd = rv. But what if we are not in equilibrium? Suppose, for instance, a jar of water with little or no oxygen in it is opened to the atmosphere and left on a countertop. If we were to monitor the concentration of oxygen in the water over time, the results might look like:
Figure 1. Variation of DO Level v.s. the Saturated DO Concentration
As can be seen, the concentration approaches some maximum concentration (CS) asymptotically. Experiments of this kind have shown that the rate of change of dissolved gas concentration can be expressed mathematically by:
where C = the concentration of dissolved gas (mg/L or mole/L) CS = the concentration of gas under saturated conditions (called Cequil in your text, p 267) kL = gas transfer coefficient (cm/s or m/s) (kL*a) = overall gas transfer coefficient (1/s) a = ratio of gas/liquid interfacial area to liquid volume (cm-1 or m-1) t = time
CS is called the saturation concentration. It is the concentration when the system is at equilibrium. In other words, it is the concentration predicted by Henry's Law. Let's see if this equation makes sense. If C < CS, then dC/dt > 0, and the concentration rises, indicating that the net movement is from the air into the water (rd>rv). If C > CS, then dC/dt < 0, and the concentration drops, indicating that the net movement is from the water into the air (rd<rv). If we integrate this equation and apply the appropriate boundary condition (at t=0, C=C0), the resulting equation is:(Do you remember how to do this?)
)( 0 )(
atk ss
LeCCCC --=- C0: Dissolved Oxygen Concentration at t = 0 sec In the environment, C is usually less than CS, so the term (CS-C) is often called the "deficit", D. Rewriting the integrated equation in terms of the deficit yields:
D = D0e-(kLa)t D0: Deficit of Dissolved Oxygen at t = 0 sec
0 1 2 3 4 5 6 7 8 9
10
0 5 10 15 Time (min)
C on
ce nt
ra tio
n of
di ss
ol ve
d O
2
(m g/
L)
Cs = 9 mg/L
))(( CCak dt dC
SL -=
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Notice that this is an exponential decay function. Can you see it in the graph of experimental results shown above? It is helpful to think of the deficit (or CS-C) as a "driving force" for net gas transfer. Looking at the differential equation dC/dt = kLa(CS-C), the slope of the C curve is steepest when C is farthest from CS (i.e., when the deficit is largest). As C approaches CS, the slope gets smaller and smaller. In other words, the driving force for net gas transfer diminishes as the system approaches equilibrium. This is another example of Le Chatelier's Principle. When the system's equilibrium is disturbed, it attempts to re-establish that equilibrium. The larger the disturbance, the "harder" the system works to re-establish the equilibrium. Besides the deficit, what other factors do you think might affect the rate of gas transfer? Recall our hypothetical experiment. We left the jar sitting quietly on a countertop and watched the oxygen climb to the saturation value. What if we mixed the jar? What if we bubbled air through the jar? Would the degree of mixing or the size and number of the bubbles affect the results? Yes they would. Physical factors like mixing and bubbles change the gas transfer rate by their effects on the overall gas transfer coefficient, kLa. Identifying and measuring the many factors that affect kLa is usually so difficult that it is easier to measure k directly by experiment. We can do this by linearizing the integrated deficit equation as shown below:
D = D0e-(kLat) D/D0 = e-(kLa)t ln(D/D0) = -(kL a) t (or, y = mx) If we run the hypothetical experiment and graph the results according to the linearized equation, we should get a plot that looks like: The slope of a regression line forced through the origin will be kLa. Measuring the kLa values for several gas transfer systems is the goal of our lab exercise this week.
Figure 2. Linear Regression of the –Ln (D/D0) Time Sequence Data
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LAB 3: BIOCHEMICAL OXYGEN DEMAND OBJECTIVES The objectives of this lab are to measure the seven-day Biochemical Oxygen Demand (BOD7) of a dairy wastewater sample, to learn how BOD is related to the amount of organic substances present the wastewater, and to learn the importance of dilution.
INTRODUCTION1
Biochemical Oxygen Demand (BOD) is the amount of oxygen used in the metabolism of biodegradable organics. Typically, BOD tests are conducted over a five day period. However, for convenience, this experiment will be conducted over a seven day period. To understand how BOD is related to organic substances in water, it is necessary to know about the degradation of these substances. Bacteria and other forms of microbes feed on organic waste. These microbes, like most life forms, consume oxygen while metabolizing food, causing a demand for oxygen, or BOD. In water, there is only a limited amount of oxygen, dissolved oxygen (D.O.), present for consumption. If all the available oxygen is consumed by microbes, it can be catastrophic to the aquatic life in the water that also depend on the D.O. for survival. In wastewater treatment, these same microbes are used to aid in the removal of organic substances from the water. This is done through aeration, by transfer of oxygen to the water through diffusers or mixers, to keep the D.O. constant until the microbes consume all the organic matter present. For these reasons it is important for engineers to be able to determine the BOD of waters as an indirect measure of the organic substances in the water.
BACKGROUND
Biochemical Oxygen Demand BOD measures the amount of oxygen required by aerobic bacteria and other microorganisms which decompose organic matter. Microorganisms use oxygen directly proportional to the amount of organic matter oxidized in biochemical oxidations. BOD determination is used in studies to measure the purification capacity of streams. It also serves regulatory authorities as a means of checking the quality of effluents discharged into such waters. The BOD test is the only test that can measure the amount of biologically oxidizable organic matter present. It determines the rates at which oxidation will occur or BOD exertion in the receiving bodies of water. This is the reason why BOD is a major criterion used in stream pollution control where organic loading must be restricted to maintain desired dissolved-oxygen levels. BODt is the amount of BOD exerted during the duration of a test, while the BODu is proportional to the total biodegradable organic content of the water. The BODu is the ultimate or limiting BOD, meaning the DO reaches a point where it stops dropping and remains constant. The total DO drop at this point represents the ultimate BOD. This means that the bacteria consumed all of the organic material or that there was not a sufficient supply of oxygen for decomposition.
Dissolved Oxygen (DO) DO is a major determinant of water quality in streams, lakes, and other water courses such as groundwater. The saturation of oxygen in water is a function of temperature and pressure. Also, the concentration of dissolved solids in the water affects the saturation levels of oxygen in water. For
1 Adapted from the Technical Information Series—Booklet No. 7 by Clifford C. Hach, Robert L. Klein, Jr., Charles R. Gibbs
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example, the solubility is less in saline waters. Also, solubility of oxygen varies greatly with the temperature. DO will be measured using an oxygen probe and meter for this experiment. For a successful experiment, two factors must be taken into consideration. Since the rate of oxygen consumption is temperature dependent, temperature must remain constant, ± 2 �C for an accurate analysis of BOD. Also the reaction needs to take place in the dark. This is very important because the presence of algae will produce an error in the BOD measurements. Algae photosynthesizes in the light, producing oxygen in the testing bottles, creating an error in the end results.
Dilution The size of the sample volume plays an important factor in the success of a BOD measurement. If the sample volume is too large, oxygen supply will be diminished before the five day mark yielding an inaccurate BOD analysis. Bacteria will be overwhelmed by the massive amount of organic material available and will quickly deplete the oxygen supply since decomposition of organic material is proportional to oxygen consumption. In the other case, if the sample volume is too small, the change in DO will be too small to give reliable results. The change in DO will be minimal. Therefore, the dilution is crucial to ascertain reliable results. The following table shows the standard DO thresholds, which will be used to discredit sample volumes falling into the two DO concentrations conditions:
Threshold Test Problem Final DO < 1 mg/L Sample volume too large DOo – DOt < 2 mg/L Sample volume too small
Any test that falls in either of the conditions listed above is considered invalid. The effect of sample volume on the D.O. in the bottle during a BOD test is shown in the Figure 1.
Figure 1. Effect of Sample Volume on Dissolved Oxygen During a BOD Test
For any chosen dilution, you can calculate the range of BOD values that give valid tests. The upper BOD value can be calculated from assuming the starting D.O. to be a typical saturation value at room temperature (say, 8 mg/L) and setting the final D.O. equal to 1 mg/L. The lower BOD value can be calculated by setting the �D.O. value to be 2 mg/L. A table of BOD ranges for different dilution factors is shown below in Table 1. (Can you duplicate these calculations?)
0
2
4
6
8
10
0 1 2 3 4 5 6 7 8
time (day)
D O
(m g/
L)
Too much organic mat'l
Not enough organic mat'l
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Table 1 BOD Measurable with Different Sample Dilutions (Adapted from Sawyer, et. al.) Volume of Sample in a 300-mL Bottle
(mL) Range of BOD Measurable
(mg/L) 0.02 30,000 - 105,000 0.05 12,000 - 42,000 0.1 6,000 - 21,000 0.2 3,000 - 10,500 0.5 1,200 - 4,200 0.7 857 - 3,000 0.8 750 - 2,625 1 600 - 2,100
1.25 480 - 1,680 1.5 400 - 1,400 1.75 343 - 1,200
2 300 - 1,050 5 120 - 420 10 60 - 210 50 12 - 42
100 6 - 21 300 2 - 7
Based on past experience, we've chosen sample volumes for the wastewater to be tested. To cover possible fluctuations from past experience, we will choose several dilutions to cover a wide range of possible BOD5 values. Some bottles will exceed the threshold values discussed above, but if we've chosen wisely, at least one bottle will result in a valid test. Length of Test The standard BOD test is read after 5 days. For the sake of convenience, we'll read ours after 7 days during the normal lab period.
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BOD Equations Expressing the word definition of BOD in the form of an equation yields:
VolumeSample VolumeBottleSampleConc.)DOFinal-Conc.DO(Initial
SampleofVolume usedDOofMass
BOD ´
==
The shorted version is:
P DO-DO
BOD t0t =
where: BODt = Biochemical oxygen demand at time t, mg/L DO0 = Initial dissolved oxygen concentration in the BOD bottle, mg/L DOt = Final dissolved oxygen concentration in the BOD bottle, mg/L P = Ratio of the sample volume to the bottle volume (Vsample/Vbottle) Vbottle = 300 mL for standard bottle
CONCLUSION Biochemical Oxygen Demand (BOD) is the amount of oxygen used in the metabolism of biodegradable organics. The design of treatment facilities relies on the information concerning the BOD of wastes. It factors into the choice of treatment methods and is used to determine the size of treatment units, particularly trickling filters and activated-sludge units. The BOD test is used to evaluate the efficiency of various processes during operation at treatment plants. Typically, BOD tests are conducted over a five day period. However, for convenience, this experiment will be conducted over a seven day period. The objectives of this lab is to measure the seven-day Biochemical Oxygen Demand (BOD7) of a dairy wastewater sample and to learn the importance of dilution.
Second Week of BOD Pre-Lab, Final Report & Presentation The material below applies to (1) the pre-lab for the second week of the BOD lab, (2) the final report and (3) the presentation. Define and discuss the terms and topics listed below in your own words (paraphrase the works of others). CITE YOUR SOURCES (provide references).
1. A. Ultimate BOD (BODU) B. Carbonaceous BOD (CBOD) C. Nitrogenous BOD (NBOD) D. Addition of nutrients E. Addition of seed microorganisms. F. Modeling BOD Remaining (BODR(t)) and BOD Exerted (BOD(t))*.
2. Theoretical Oxygen Demand (ThOD) 3. Chemical Oxygen Demand (COD)
* See additional notes (located after the procedures below).
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LAB 3: PROCEDURES Objective: The objectives of this lab are to measure the seven-day Biochemical Oxygen Demand (BOD7)
of a dairy wastewater sample, to learn how BOD is related to the amount of organic substances present the wastewater, and to learn the importance of dilution.
Bottle Set-up and Incubation 1. Each group will be given a water sample to test. The instructor will brief you on the source of the
sample to be tested. 2. We don't know in advance the BOD values of the different waters. Consequently, we will need to
make several dilutions so that the change in dissolved oxygen (DO) is neither too great nor too small. Each group will set up several BOD bottles, but because some tests are likely to fail, we'll combine everyone's data at the end.
3. Your instructor will direct you on how many bottles to set up, and what volumes of sample to test. Use a pipet for volumes less than 50 mL, and a graduated cylinder for larger volumes.
4. Without excessive agitation fill the remaining volume of each BOD bottle with dilution water provided by your instructor.
5. Measure the dissolved oxygen concentration using a DO meter and probe. Be very careful that there are no air bubbles in the bottle when you do this measurement. a. Remove the probe shield (remove it carefully to avoid damaging the plastic probe cap). b. When the probe shield is off, be careful to protect the probe tip from being scraped. c. Record the bottle numbers, their contents, and the DO concentrations (after dilution).
6. Seal each bottle with a glass plug. Be very careful that there are no air bubbles in the bottle. After sealing, make sure that there is water in the "wet well" at the top of the bottle. If not, remove the glass stopper, add more dilution water, and reseal. Finally, place a plastic cap over the "wet well" to prevent the seal from evaporating.
7. Place all of the bottles in the constant temperature room as directed by your instructor. Dissolved Oxygen Measurement 8. After seven days (i.e., the next lab period), take your samples from the constant temperature room.
Before opening each bottle, note (and record) if the water seal in the wet well is not intact and if any bubbles are visible in the bottle. (Bubbles may be CO2, but they may also indicate a malfunctioning seal.)
9. Measure the DO concentrations and record your data. 10. Rinse the DO probe with DI water; carefully replace the probe shield (CAUTION – do not tighten
shield while it is resting on the plastic probe cap; move it back before tightening.) 11. Wash the glassware, caps, etc…. and set them to dry on the glassware rack. Data Analysis 1. Check your data against the threshold values (i.e., the minimum allowable DO after seven days and
the minimum allowable change in DO). Throw out any invalid test results. 2. Calculate the BOD7 values associated with your valid test data and record those values. (Don't forget
to provide sample calculations.) 3. Collect and record the valid BOD7 values from the other groups. 4. Summarize all of the valid test results (BOD values from the entire class) the data statistically by
reporting the range, mean and standard deviation values.
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LAB 3: ADDITIONAL NOTES ON BOD
Let's think about how one might measure organic material. If the material was a single substance, such as sucrose (table sugar), it would be possible to employ one of several chemical techniques to isolate and measure the amount of sucrose. But suppose you wanted to measure a mixture of organics. It might be possible to isolate each and measure it through chemical means, but this wouldn't be practical or useful. In wastewater or even natural waters like rivers and lakes, there may be hundreds of organic substances. Analyzing for all of these compounds would be a huge task, and wouldn't tell us what we really want to know anyway. We could use a "lumped" parameter like Total Organic Carbon (TOC), but this doesn't tell us anything about the biodegradability of the organic substances or the rates of conversion. Biodegradability is a key issue.
All natural environments teem with microbial life (e.g., a typical agricultural soil can contain as many as a million bacteria per gram!). For these microorganisms, organic waste is a source of food. Like virtually all higher forms of life (i.e., you and me), many of these microbes use oxygen to metabolize their food. As microbes degrade organic substances, they pull oxygen from the local environment in direct proportion to the amount of organic material metabolized. There's lots of oxygen in the air, but the amount dissolved in water is limited (recall the gas transfer lab). If enough organics are present, microbes will use up all of the available oxygen and fish will die. This is the reason engineers are interested in measuring organic materials.
The Biochemical Oxygen Demand (BOD) test was developed in Britain in the early part of this century. It is a simplified physical model of what would happen if an organic waste is added to a stream. In the test, a liquid containing organic wastes (either full strength or diluted) is placed in a bottle. The DO is measured and the bottle is sealed. After some incubation time (usually 5 days), the DO of the sample is measured again. Because the bottle is sealed, the difference in the DO values represents the oxygen used by microbes in degrading the waste. The change in DO is an indirect measure of the organic substances in the bottle. Through this process we see that biodegradable organic materials have associated with them a potential demand for oxygen when they are degraded (hence the name biochemical oxygen "demand").
The calculation of the BOD uses the following equation:
Where
DO0, DOt = dissolved oxygen concentrations at t=0 and t=t P = dilution factor = Volume of sample � Volume of the BOD bottle (300 mL)
If the waste is strong, the oxygen in the bottle will run out before the end of the incubation period. This is why we dilute the waste. If we over-dilute it, the change of DO will be too small to be statistically reliable.
P DODO
BOD t -
= 0
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A typical graph of DO in the BOD bottle as a function of time might look like Figure 1 below:
Figure 1. Decay of the DO in the BOD Bottle
As shown in the graph, if we wait a long time, the bacteria consume all of the organic material, and the DO stops dropping. The total DO drop at this point represents the "ultimate BOD" (also called BODu or sometimes BODL for "limiting BOD"). Twenty to thirty days is a long time to wait, so usually the bottles are opened and measured after 5 days, and the test results are reported as the "5-day BOD" or BOD5. Originally, 5 days was chosen as the incubation time because that is the longest travel time of any river in Britain. We continue to use this value today, although the choice of incubation time is really arbitrary. To convert between BOD5 and BODu, we need to know something about the kinetics of the BOD test. Suppose we set up 7 BOD bottles on day 0. Periodically over the next month, we open a bottle and measure the DO. If we use these DO values in the BOD equation shown above, we would see an increasing BOD result. In other words, BOD1 > BOD2 > BOD3 … because DO1 < DO2 < DO3 …. If we plot the BOD values against time, we get a curve like that shown in Figure 2.
Figure 2 Increase of BOD values with time
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Notice that the measured values of BOD start at 0 and end up at the ultimate value. When you compare Figure 2 with Figure 1, you see that Figure 2 is actually Figure 1 flipped over. While it is not always true, very often we can fit a first order exponential model to these data. The equation of the model is: BODt = BODu (1 - e-kt)
If you plug "5" in for "t", you can convert between BOD5 and BODu using this equation.
It is important that you understand the difference between BODu and BODt. BODu is proportional to the total biodegradable organic content of the water. If you add more biodegradable organic material to water, BODu will increase. BOD5 is the amount of BOD exerted in the 5-day test and is proportional to the amount of organic material degraded in 5 days. It is normally assumed that BOD5 is proportional to the BODu but this may not always be true. In Figure 3 for instance, three samples all exerted about the same BOD5, yet they all had different BODu's. How might this happen? Different organic substances degrade at different rates. Most waters contain a mixture of organic substances. In two samples, if the amounts of organic substances degradable in 5 days are equal, then the BOD5 values will be equal. The two samples could have different amounts of organics that degrade more slowly, so the BOD u values will be different.
Figure 3. Measured BOD variation over time for different organic substances
Please note that the first order model does not always fit BOD data. Don't be surprised if in real life, you get data that doesn't plot out in such neat curves. Like any mathematical model, use the first order equation with care.
Finally, there is one more kinetic issue to consider. Suppose we put a water sample with biodegradable organics into a batch reactor and supply it with oxygen in the form of air bubbles. Now suppose we take a sample out each day and test the sample in a BOD5 test. We would see that the BOD5 of the original water decline over time as shown below:
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Figure 4 BOD decay over time
What we see here is that the BOD5 of the water declines over time. If the k value is constant, then the relationship between BOD5 and BODu will be constant. So if the measured BOD5 declines over time, the BODu, representing the organic material in the water, also declines over time. In other words, the organic material remaining in the water is being converted to something else -- CO2 and other nondegradable substances -- by the bacteria. Often, a first order decay reaction can be fit to these data. If the organic material in the water (expressed in BOD units) is called "L" (as it is in your book), then
L = L0 e-kt
Usually L is measured as BODu. We will see this relationship often in solving mass balance problems for the degradation of organic materials in streams, lakes, or treatment plants. Be warned, however, that not all organics decay according to a first order model. Don't be surprised by real-life data following a different pattern.
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LAB 4: ADSORPTION Objectives The purpose of this exercise is to demonstrate the phenomenon of adsorption and compare the adsorptive behaviors of different materials. Batch adsorption tests will be conducted using methylene blue (MB), a cationic organic dye, and several different adsorbents (activated carbon, clay, and sand). Background Adsorption of solutes originating from the bulk of one phase on to the surface of an adjacent phase occurs throughout nature. This mass transfer process is used to remove contaminants in water treatment operations. The Freundlich and Langmuir isotherms are two mathematical models commonly used to describe adsorption phenomenon. The isotherm is an empirical relation between the concentration of a solute on the surface of an adsorbent (mass/mass) to the concentration of the solute in the liquid with which it is in contact (mass/volume or mass/mass). Another model is the B.E.T. The Freundlich isotherm equation is as follows:
q = x/m = KCe 1/n
where: K and n are empirical coefficients; K represents adsorption capacity at unit concentration 1/n represents strength of adsorption.
x = mass of MB adsorbed (mg) = (mass of MB in the liquid at the start) - (mass of MB in the liquid at the end) = (Ci*VMB)start - (Ce*VMB)end = (Ci - Ce)VMB Ci = initial concentration of MB in the solution (mg/L) Ce = final concentration of MB in the solution when at equilibrium (mg/L) VMB = volume of MB solution (L) m = mass of adsorbent (the material that adsorbs MB) (mg)
Procedures Calculating the methylene blue concentration Methylene blue can be measured using a spectrophotometer. The intensity of a ray of monochromatic light decreases exponentially as the concentration of light-absorbing material in the medium (water in our case) increases (Beer's Law). A spectrophotometer can sense the absorbance of light at specified wavelengths.
Absorbance = log(I0/I) = proportional to concentration where I0 and I are the intensities of the light entering and leaving the test sample, respectively. Light absorbance at a wavelength of 655 nm has been found to work well for detecting methylene blue. The absorbance can be related to the concentration of the light-absorbing substance through use of a linear calibration curve. 1. Your instructor will provide a series of dilutions of a stock methylene blue solution. Use a pipet to
remove about 10 mL of the solution and place it into a small test cell designed for the spectrophotometer (we will use the square 10 mL sample cells).
2. Zero the spectrophotometer unit using deionized water, then measure absorbance values of the standards your instructor gives you and report those values to the class to be recorded in your lab notebook.
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3. Plot concentration of standard (y-axis) against absorbance (x-axis) and use the trendline function on the computer (in Excel) or on your calculator to fit a linear curve through the data. Be sure to include the equation for the curve on your plot. What we want is a calibration curve equation that yields concentration values when experimentally-found values of absorbance are plugged in to it.
4. Use this equation to calculate methylene blue concentrations from absorbance values in the rest of the lab.
Adsorption Test 1. Place 50 ml of methylene blue solution into four100-ml beakers. 2. Weigh out the adsorbent material assigned to your group and prepare to add it to the beakers with the
dye in the next step. It's OK if you don't hit the specified amount exactly; just record the mass you use.
Mass of Adsorbent: 0.01, 0.05, 0.10, & 0.50 grams 3. When your instructor gives the signal, add all of your adsorbent masses to the flasks containing MB
and swirl or stir the beakers for 15 minutes. 4. Transfer approximately 15 mL of solution from the top of your beaker (avoid the solids) into
centrifuge tubes and give it to your instructor for centrifugation. 5. After centrifugation, remove your tube. Use a pipet to remove about 10 mL of the supernatant, and
place it into a small test cell. Do this step very carefully. Suspended solid material will cause you to get an inaccurate absorbance reading.
6. Put the test tube into the spectrophotometer and read the absorbance.
Data Analysis 1. Calculate the mass of MB adsorbed per mass of adsorbent, q = x/m 2. Using your data, plot either:
a. log(x/m) as a function of log(Ce) on arithmetic scales. A linear trend line is used to obtain the Freundlich coefficients as follows: K = 10y-Intercept and 1/n = the slope.
b. x/m as a function of Ce on log-log scales (as shown in Figure 1 below). A power function trend line yields the Freundlich equation directly.
3. Exchange data with your classmates; Compare the results; In the final report explain why the adsorption capacities of the materials are different.
4. For the best adsorbent use the Freundlich isotherm equation to generate a smooth plot and plot values of x/m as a function of Ce on arithmetic scales.
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Figure 1. Example of Freundlich Isotherm-generated data plotted on log-log scales. When plotted in this manner a power function trend line yields the Freundlich equation directly. [Note: When plotting Log(q) vs Log(Ce) on arithmetic scales a linear trend line is used to obtain the Freundlich coefficients as follows: K = 10y-Intercept and 1/n = the slope.]
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LAB 5: SOLIDS IN WATER
BACKGROUND
Tests will be performed to measure solid materials present in wastewater. These tests are of great use in assessing the reuse potential of a wastewater and in determining which types of treatment processes are best for the wastewater. Four types of tests will be performed to determine total solids (TS), total suspended solids (TSS), volatile suspended solids (VSS), and turbidity. The measure of solid materials within water supplies is of particular importance to engineering practice. While the effects of such measures are of primary concern to environmental and biological engineers, they are not independent of other engineering applications. Water’s abundance and its significant role in life processes has made it into one of the many commonalties shared by all engineering disciplines. As such, the effects that solid materials have on water’s chemical and physical properties are of considerable importance to the practicing engineer. With this in mind, this lab was developed to familiarize one with four different measures of solid materials in water and to provide a means to obtain such measures.
TOTAL SOLIDS
The first measure, total solids (TS) deals with both undissolved and dissolved solids. Undissolved solids are usually referred to as suspended solids (discussed further below). Total solids, as the name implies, is a measure of all residue left after evaporation of the water phase, usually at around 103�C. This test consists of placing a known volume of a sample in a pre-weighed evaporating dish, allowing all water to evaporate, and then determining the mass of dried solids as the difference between the dish plus dried solids and the dish alone. Total solids is calculated as mass of dried solids divided by the sample volume as follows:
Note: A & B are defined differently here than in the eqn.’s that follow.
where: TS = Total Solids (mg/L) A = Final mass of container with sample after dry (mg) B = Initial mass of empty container (mg) V = Sample volume (mL)
There are many reasons for determining the total amount of solids present in water. One reason is for determining if water is suitable for domestic use, such as for drinking water. Water that contains more than 500 mg/L of solids often has a laxative effect on humans, especially those who are not accustomed to it (Schroeder 1987). It is also extremely valuable in analyzing raw and digested sludges, which is important when designing and operating sludge-digesters, vacuum-filters, and incinerators for wastewater treatment facilities. Another use of the total solids test is in detecting changes in the water density of waterways to determine if there is possible contamination from wastewater.
TOTAL SUSPENDED SOLIDS
The second measure concerns total suspended solids (TSS). Suspended solids are the undissolved solids present in water. Analysis of suspended solids is useful when the turbidity measurements do not provide adequate information and is very valuable in analyzing polluted waters. Measurement of suspended solids aids in evaluating the strength of wastewaters and helps when evaluating the efficiency of treatment facilities. While total solids measured all solid material, total suspended solids is a measure of only those solids that are larger than approximately 1.2 µm in size and typically held in suspension within the sample. In other words, total suspended solids are those that are “undissolved”. The measurement of TSS is accomplished by means of a filter that removes all suspended solids within a given sample. Total suspended solids can be classified into two main categories, those that are volatile and those that are not.
V 1000 xB)-(A
TS =
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Note: A & B are defined differently here than in the other equations.
where: TSS = Total Suspended Solids (mg/L) A = Final mass of container with filter and sample after dry (mg) B = Initial mass of container with filter (mg) V = Sample volume (mL)
VOLATILE SUSPENDED SOLIDS
Volatile suspended solids (VSS), also called organic solids, are solids that are combustible at 550�C. The volatile solids test is often used in estimating the organic characteristics of solids present in a wastewater and thus is an indirect measurement of the organic matter present. VSS is measured by igniting the total suspended solids and measuring whatever is left. The mass that is left behind after combustion is a measure of the non-volatile solids while the mass lost during incineration comprises the amount of volatile solids.
V 1000 xB)-(A
VSS = Note: A & B are defined differently here than in the other equations.
where: VSS = Volatile Suspended Solids (mg/L) A = Mass of container with filter and sample after drying and before ignition (mg) B = Mass of container with filter and sample after ignition (mg) V = Sample volume (mL)
TURBIDITY
The final measure of solid material within water is called turbidity. The turbidity test is useful in water treatment for determining the quality of drinking water. When colloidal (clay-size) particles are present in the water (mud puddle water for our purposes), they cause incoming light to become scattered, thus causing the water to appear cloudy or turbid. The portion of light that scatters at an angle of 90 degrees from the direction of the source light is measured using photo-sensitive cells. A turbidimeter is an instrument capable of measuring turbidity. Turbidity is often measured using a special the unit NTU.
CONCLUSION
The measure of solid materials within water supplies is of particular importance to engineering practice. While the effects of such measures are of primary concern to environmental and biological engineers, they are not independent of other engineering applications. Water’s abundance and its significant role in life processes has made it into one of the many commonalties shared by all engineering disciplines. As such, the effects that solid materials have on water’s chemical and physical properties are of considerable importance to the practicing engineer. In this Lab, tests will be performed to measure solid materials present in wastewater. Four types of tests will be performed to determine total solids (TS), total suspended solids (TSS), volatile suspended solids (VSS), and turbidity.
V 1000 xB)-(A
TSS =
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LAB 5: PROCEDURES (Measuring solids in water)
Purpose of Lab The purpose of this lab exercise is for the student to become familiar with four different measures of solid materials in water: (1) total solids (TS), (2) total suspended solids (TSS), (3) volatile suspended solids (VSS), and (4) turbidity. Procedures Total Solids 1. Weigh two (2)* clean evaporating dishes on the precision balance (the one with a glass enclosure
with sliding sides). Record those values. 2. Add a known volume of representative sample (about 10 ml) to the dish and place it in the drying
oven (at 105 � C) until the water is completely evaporated away. 3. Weigh the evaporating dish plus sample. Record that value. 4. Remove the dish from the oven and cool in a dessicator. 5. Reweigh the dish (with solids) and record the mass. * Normally three or more samples would be analyzed. Do you know why? We are using less due to the
limited time available during the lab period. Total and Volatile Suspended Solids (TSS and VSS) 1. Obtain two (2) filters and aluminum boats from your instructor. These filters will have been
previously washed and dried. Because you are weighing things on an analytical balance, which can read 1/10,000 gram, you want to be careful to avoid contaminating the sample with dust or dirt from the bench top, oils from your fingers, etc. So, handle the filters and boats with tongs or tweezers at all times, and rest them only on clean surfaces. One suggestion is keep the boats on a designated clean piece of paper.
2. Identify the boat with a physical mark (indentation). This is your identification mark, so take care that each filter stays with the same boat.
3. Weigh the aluminum boats with the filters in them on the analytical balance. Record those masses. 4. Using tweezers, place a filter on the suction apparatus as directed by the instructor. Be sure the filter
flask is clean. 5. Filter 25 to 50 mL of representative sample (or another volume as directed by the instructor). You
want to put as much sample through the filter as you can without causing it to clog. Stop adding sample when the flow rate through the filter slows noticeably. Approximately 15 ml of filtrate will be needed for turbidity measurement.** Record the amount of sample filtered, and take care that your sample is well-mixed at all times. ** Dilution, while not desirable, may be necessary for water containing significant amounts of particulate
matter in order to obtain 15 ml of filtrate. If done, the resulting TSS and VSS values must be multiplied by the appropriate dilution factor and diluted sample must be used for the initial turbidity reading.
6. Turn off the vacuum, and without disassembling the filter holder and funnel, remove them from the filter flask. Pour the filtrate into a labeled beaker for later turbidity analysis. Replace the filter holder and funnel and restart the vacuum.
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7. Wash the filter with three successive washings of deionized water (about 10 mL each) while continuing suction. Wait 3 minutes between washings. If some solids have been caught in your glassware (graduated cylinder, funnel), rinse them out with these washings.
8. Turn off the suction and, using tweezers, place the filter in its boat. 9. Repeat steps 4-8 for each sample as directed by your instructor.
10. Place the all the aluminum boats and filters in the drying oven (at 105 � C) for at least one hour. 11. Remove the aluminum boats and filters from the oven and place in the desiccator for at least 15
minutes or until they are at room temperature. 12. Weigh the aluminum boats and filters. Record those values.
13. Place the filters and boats into the muffle furnace (at 550 � C) for 20 to 30 minutes. The muffle furnace is hot, so be careful!
14. Remove the filters from the muffle furnace and place them into the desiccator for at least 15 minutes or until they are at room temperature.
15. Weigh the filters and record the masses again. Look again at the filters. How do their appearances
compare with their appearances before combustion at 550 � C? Turbidity 1. Turbidity is measured on an instrument. Your instructor will describe how to use the instrument. 2. Measure and record the turbidities of the available water samples as directed by your instructor. 3. Measure and record the turbidities of the filtrates for those samples you used in the TSS test. Data Analysis Calculate the TS values for the samples you analyzed. TS is the mass of dry solids on the dish (after drying at 105 �C) divided by the volume of sample passed through the filter. Express your results in mg/L (milligrams of dry particles per liter of suspension). Show all your work in your lab notebook. Calculate the TSS values for the samples you filtered. TSS is the mass of dry solids on the filter (after drying at 105 �C) divided by the volume of sample passed through the filter. Express your results in mg/L. Show all your work in your lab notebook. For the samples you combusted at 550 � C, calculate the VSS. VSS is the mass of dry solids that volatilized (left) the filter after combustion at 550 �C divided by the volume of sample passed through the filter. Express the VSS in mg/L and as a percentage of the TSS. For the samples you filtered, calculate the percentage of solids captured on the filter based on the percentage change in the turbidity. How does this value compare with the ratio of TSS to TS based on mass measurements? If there is a difference, explain why.
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Discussion Questions
1. Calculate the total volume of the samples used in the TS test based on a calculation of the sample mass from measurements made on the analytical balance and on an assumption that sample density is 1.00 kg/Liter. Compare the gravimetric-based calculated volume with the volumetric measurement. Which method is more accurate? Why? Under what circumstances would you use one method rather than the other?
2. Knowing what you do about the source of the water sample tested, and the nature of the particles in the sample, do your lab results make sense? Explain why or why not. (For instance, are the relative TSS and VSS values what you expected, or were there any surprises?)
3. What fraction of the turbidity was removed on the glass fiber filter? If that fraction is less than 100%, discuss why all the turbidity wasn't removed. (Hint: What was in the filtrate?)
4. Do the TSS and turbidity tests measure the same thing? When would you use one type of measurement and when the other? (Second hint: Think about mass vs. number.)
5. Would you expect the analytical results to be higher than, lower than, or the same as the true value under following conditions, and why?
a. Weighing a warm crucible? b. Estimate the organic content by volatile-solids analysis of a sample containing a large
quantity of organic materials having high vapor pressure. c. Estimate the organic content of a sample by combustion at 800oC
Lab Write-up 1. Report your data and the results of your calculations in your lab notebook.
1. Provide sample calculations for every different kind of calculation done. Be sure to check for appropriate significant figures.
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Example data table format:
Water Sample Identification Table Sample ID Description of Sample
Total Solids (TS) data and calculated results. Sample
No. Mass
boat (g) Mass boat + sample
before drying (g) Mass boat + solids
after drying (g) Vol. of
sample (mL) Mass of
sample (b) Total Solids
(mg/L)
Total Suspended Solids (TSS) and Volatile Suspended Solids (VSS) data 1 Sample
No. ID Number
on Boat Mass of clean
filter and boat (g) Mass of filter, boat, and solids after drying (g)
Mass of filter, boat, and solids after combustion (g)
Total Susp. Solids (TSS) and Volatile Susp. Solids (VSS) data 2 and calculated results. Sample
No. Volume of sample
filtered (mL) Total Suspended Solids (mg/L)
Volatile Suspended Solids (mg/L)
VSS/TSS ratio (expressed as %)
Turbidity data and calculated results Sample
No. Turbidity of sample
(NTU)
Turbidity of filtrate
(NTU)
Percent of turbidity removed by filtering (%)
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LAB 6: CHEMICAL COAGULATION JAR TESTING INTRODUCTION Coagulation is accomplished through neutralization of negatively charged colloidal (clay size) particles. This is necessary since the repelling forces between the particles, due to their negative surface charge, is greater than the attractive body forces when the particles are separated. For this reason, without the neutralization process, particles do not collide, stick together and grow into larger particles that would settle to the bottom, thus, the small particles remain in suspension.
The purpose of chemical coagulation is to condition the water such that small particle growth is promoted, thus increasing particle size and density. The physical process in which the particles actually group together forming “flocs”, is known as flocculation. This increase in size and density of the particles facilitates removal in gravity sedimentation and water filtration processes. Coagulation and flocculation are two separate stages. They play a vital role in improving the aesthetics of drinking water and also remove microorganisms and other contaminants that can be harmful to our health.
BACKGROUND
Coagulation refers to the chemical alteration of the suspended particles that allows the particles stick together in larger clumps called flocs. Flocculation refers to the physical process that collects these flocs into even larger clumps that easily settle out by means of gravity. Particles in water can be classified into two different size groups: colloidal and suspended solids. Colloidal particles, also referred to as colloids range in size from 1 micrometer to 5 nanometers while suspended solids generally consist of particles larger than 0.5 micrometers. These particles can also be differentiated by the nature of how they react with water molecules. Hydrophobic particles are defined by having a low attraction to water and are generally inorganic in origin. Hydrophilic particles react more readily with water and often consist of organisms both living and dead such as bacteria, algae and viruses. Regardless of which group these particles may fit into it is their chemical properties and not their size that dictates the way in which they will react during the coagulation and flocculation processes.
Almost all colloidal particles have a negative charge, meaning that they will be attracted to particles holding a positive charge and repel other particles that are also negative. This tendency to repel other negatively charged ions slows the process of gravity induced settling considerably. Due to this slow rate of settling it is necessary to artificially expedite the rate at which the particles can be separated.
When a coagulant such as alum is put into solution in water it is rapidly dissolved into aluminum ions and sulfate ions. These molecules carry a positive charge and combine with the negatively charged colloids to form various ions of aluminum oxides and hydroxides. The ions that are produced depend on many variables such as temperature, pH and the speed and method of mixture. Two types of reactions then occur between the ions and the colloids, both of which result in coagulation. One of these types is known as charge neutralization and occurs when the positive charges from the aluminum ions counteract the negative charges of the colloids. The other type is called bridging and occurs when aluminum hydroxide ions stick together several negative colloid particles, acting as a sort of chemical glue between like charges.
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Coagulants, in our case aluminum sulfate, are used to neutralize the charges. When the aluminum sulfate is added to water it dissolves and forms aluminum ions and sulfate ions. The aluminum ions are of particular importance in this process since they are unstable and form stringy charged species of aluminum oxides and hydroxides that adsorb to particles, thereby neutralizing the particles which creates conditions where they can form flocs.
Water treatment is a process used to remove impurities within domestic water supplies for their use within domestic, commercial, and industrial applications. These processes have become particularly important within the last 100 years with the growth of both urban and industrial development. As demands for water purification have increased, so to has the responsibility of engineers to develop more efficient and cost effective applications to water treatment. Conventional water treatment usually involves two processes: clarification and disinfection. Clarification removes most of the turbidity within the water sample while disinfection destroys any pathogenic microbes that might be present. In addition to clarification and disinfection, processes of softening, aeration, and fluoridation are often used in public water sources.
Figure 1 Process Diagram of a Typical Water Treatment Plant
The process of clarification is shown above and is most often comprised of four distinct stages: coagulation, flocculation sedimentation, and filtration.
1. Coagulation is a chemical process used to destabilize colloidal-sized particles. 2. Flocculation is a physical process that increases particle-particle collision frequency and thereby
causes growth of particles into larger sizes. Through the addition of coagulants and subsequent flocculation step, non-settling particles are brought together to form larger, heavier particles called floc.
3. Sedimentation is then used to allow the floc to settle to the bottom of large storage tanks where it is then removed.
4. The water is then passed through several layers of filters to remove any of the remaining suspended solids.
flocculation
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Aluminum sulfate, or alum, is the most common coagulant used. When added to water, aluminum sulfate initially dissolves into aluminum cations and sulfate anions. Some of these aluminum ions then form various types of stringy aluminum oxides and hydroxides that adsorb to particles. This accomplishes one of two things. First, it reduces the effective charge of the suspended solids within the water sample thereby allowing the particles to collide and stick together. Second, the aluminum oxides and hydroxides bond to the negatively charged solid particles, thereby increasing their effective mass and size.
The net effect of this ion and molecular compound formation is the development of large, massive particles. This process, called flocculation, is enhanced through the use of large paddles that continuously mix the water. Mixing is necessary so that the velocities of the individual water (and solid) molecules are kept in continuous fluctuation, thereby allowing the particles to collide and stick together. As the particles become more and more massive, their density increases. Eventually the density of the floc becomes greater than the surrounding water and particles begin to settle to the bottom. This last process, known as settling, occurs only when the velocity of the water is kept to a minimum. In water treatment plants this is usually accomplished through the use of large storage tanks where the movement of the water can be closely monitored. After the floc has settled to the bottom of the tank, it can be removed by means of large scrapers. These scrapers funnel the “sludge” to the bottom of the tank where it is siphoned out.
CONCLUSION
Coagulation is accomplished through neutralization of negatively charged colloidal (clay size) particles. The purpose of chemical coagulation and physical flocculation is to promote the growth of small particles into larger particles called flocs, thus increasing particle size and density. The physical process in which the particles actually group together forming “flocs”, is known as flocculation. Coagulation and flocculation are two separate stages in the process of separating and removing suspended particulate matter from water. They play a vital role in improving the aesthetics of drinking water and also remove microorganisms and other contaminants that can be harmful to our health.
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LAB 6 PROCEDURES (Coagulation) Objectives
The objective of this lab is to become familiar with the procedures that are to be followed while conducting jar tests using a chemical coagulant. We will determine the optimum amount of alum and the mixing speed that produces the best flocculation in a water sample.
Approach Laboratory jar testing will be used to determine the optimum dosage of alum and optimum degree of mixing for coagulating and flocculating a water sample supplied by the instructor.
Materials Phipps and Bird multiple stirring apparatus 4, 1.5-L square Plexiglas “beakers” 4, 50-mL glass beakers (one per large beaker) Graduated pipets Graduated cylinders (various sizes) Hach nephelometric turbidimeter
Stock solution of alum: 2,000 mg/L of Al2(SO4)3�14H2O for doses > 10 mg/L
200 mg/L of Al2(SO4)3�14H2O for doses � 10 mg/L
Procedures A. Jar test to determine optimum dosage at a single mixing speed
1. Measure and record the mixer paddle diameter and the temperature of the test water prepared by your instructor.
2. Be sure the water in the carboy is well-mixed. Measure out 4 samples in the 1.5-L square beakers and a small sample (20 – 30 mL) for measuring initial turbidity. The samples should be between 700 and 800 mL each. (The exact volume is not important as long as the same volumes are used in all four beakers and you record what the volume is.).
3. Measure the initial turbidity of the test water, using the turbidimeter, as per your instructor's direction.
4. Locate the beakers at the center of the mixer shaft on the multiple-paddle mixer. (Running the mixer at a very slow speed will help you do this.)
5. Calculate the volumes of stock solution needed to achieve the target dosages (Cjar). Based on conservation of mass:
CstockVstock = CjarVjar ; Vstock = (CjarVjar)/Cstock The target alum concentrations are as follows: Groups1 & 3: 0, 5, 20, & 40 mg/L
Groups 2 & 4: 0, 10, 30, & 50 mg/L 6. Measure out the appropriate volumes of alum stock solution into small beakers but do not add
them to the water samples yet. 7. Turn the mixer light on. Start the mixers on the highest possible speed. When you are ready to
start timing, add the pre-measured alum stock to all the jars simultaneously. After 10 seconds,
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reduce the mixer's speed to 30 rpm and flocculate for 15 minutes. Observe and record the time required for flocs to form in each sample.
8. After flocculating for 15 minutes, turn off the mixer. Immediately, observe and record the relative average floc sizes for the samples, using a 1 to 4 scale (1 = smallest; 4 = largest).
9. Allow the samples to settle for 20 minutes. 10. Decant or pipette about 50 mL of supernatant from each beaker, taking care not to disturb the
settled material, or catch any fine particles floating on the surface. Measure the turbidities of these samples, then discard the contents of all of the beakers.
11. Select the two "best" flocculation dosages based upon the first appearance of floc, floc size, final turbidity (most significant factor), and coagulant dose.
12. Share the resulting data from Part A with the other groups. Side note: The turbidity standard for drinking water is 0.5 NTU. Is it necessary to achieve this after settling? No. In a conventional treatment plant, settling is followed by filtration. Conventional filters are capable of removing up to 10 NTU of turbidity reliably. Therefore, your goal in flocculation is to produce water that can be treated by sand filtration at the lowest cost. Discuss your choices with your instructor before proceeding.
B. Jar test to determine optimal speed In Part B, take your two water samples from the same carboy that you obtained samples from in Part A. Be sure the water in the carboy is well-mixed. If you have any doubt, measure the starting turbidity again. If it is higher, dilute the sample with tap water until you get the same turbidity as in Part A. 1. Repeat the test at a higher speed (60 rpm) using the two "best" alum dosages as determined in the
30 rpm test. It isn't necessary to record the time to first floc formation, or the relative floc sizes, but do record the final turbidities. Recommendations: a) While the beakers are undergoing slow mix (i.e., flocculation), have group members measure
out water samples and coagulant doses for the next part of the lab (15 rpm). b) Remove the 60 rpm samples from the mixer at the end of the flocculation period and set them
aside for the 20-minute settling period.. This lets you begin the next part of the procedure and save time.
2. Repeat the test as described above at a lower speed (15 rpm). Data Analysis Create a spreadsheet to accomplish the following computations, and then answer the questions: 1. Plot the data from the 30 rpm test using coagulant dosage as the abscissa (x) and final turbidity as the
ordinate (y). Include Part A data obtained from other groups so that the data series contains seven data points. Be sure your graph is properly labeled. What flocculation mechanism(s) are evident? What were the two best doses? Why?
2. Calculate the velocity gradients at the three mixing speeds. Show a sample hand calculation (see the equations provided below). For each of your best two doses, plot (x-y scatter) the data from all three tests using velocity gradient (G) as the abscissa (x) and final turbidity as the ordinate (y). Use only the data that your group collected. You will have two data series on the graph, one for each of your best two doses. Each series will consist of three data points -- a final turbidity value for each G value. Use straight or curved lines to connect data markers. Based on your results, what value of mean velocity gradient is optimal for flocculation? Why? Include in your answer a discussion how floc growth rate (or size) is impacted when G (mixing speed) is increased from a value of zero to a very large value.
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Velocity gradient and power equations
G = (P/��)1/2 where: G = velocity gradient (s-1) P = power dissipated in the water (W or ft-lb/s)
� = viscosity (N-s/m2 or lb-s/ft2)
� = volume (m3 or ft3)
If turbulent conditions are assumed (a good assumption in our case):
P = (KT n3 D5 � ) / g where: P = power (ft-lb/s) n = rotational speed (rev/s) D = diameter of paddle (ft)
� = specific weight of water (lb/ft3) g = acceleration due to gravity (32.17 ft/s2) KT = empirical constant = 2.4 for Phipps & Bird mixer (calculated from data provided by Phipps and Bird.)
Works Cited
1. Culp, Gordon L., Russell L. Culp. New Concepts in Water Purification; Litton Educational Publishing Inc., 1974
2. Hudson, Herbert E. Water Clarification Processes, Practical Design and Evaluation; Litton Educational Publishing Inc., 1981
3. Sawyer, Clair M., Perry L. McCarty, Gene F. Parkin. Chemistry for Environmental Engineering; McGraw-hill, 1994
References for G calculations:
1. McCabe, W. L., J. C. Smith, and P. Harriott, Unit Operations of Chemical Engineering, 5th ed., McGraw-Hill, 1993.
2. Rushton, J. H., "Mixing of Liquids in Chemical Processing", Industrial and Engineering Chemistry, American Chemical Society, v 44, n 2, p 2931, Dec 1952.
3. Rushton, J. H. and J.Y. Oldshue, "Mixing -- Present Theory and Practice", Chemical Engineering Progress, v 46, n 4, p 161, April 1953.
4. Reynolds, T. D. and P.A. Richards, Unit Operations and Processes in Environmental Engineering, PWS Publishing Company, 1996.
5. Wagner, E.G., "Jar Test Instructions, Conduct of Jar Tests and the Important Information Obtained", booklet from Phipps and Bird, 8741 Landmark Rd., Richmond, VA 23228, July 1993.
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LAB 7: CHEMICAL EQUILIBRIUM, BUFFERING, AND ALKALINITY
INTRODUCTION
Natural bodies of water are an important part of a carbon cycle. Carbon moves from the atmosphere into the tissues of plants and animals and into water bodies. Inorganic carbon complexes contained within water bodies affect numerous important aquatic chemistries, including alkalinity, acidity, CO2, pH, total inorganic carbon, and hardness. For example the amount of CO2 affects pH and photosynthesis. Alkalinity, acidity, and pH are explored in this lab. Water hardness will be explored in a subsequent lab.
pH
The pH of water is critical to the survival of most aquatic plants and animals. It is a chemical parameter that increases our understanding of the water’s health. pH may vary during the day or season, so it is important to take many pH readings throughout the day to obtain diurnal variations, and take many readings throughout the year at the same time of day in order to get an accurate long-term representation. pH measurement is quick and easy and establishes a baseline so that unexpected water quality changes can be understood. Some basic characteristics of pH are as follows:
· pH is a measure of how acidic or basic (alkaline) a solution is. pH < 7 is acidic; pH > 7 is basic; and pH = 7 is neutral.
· pH measures the hydrogen ion (H+) activity in a solution, and is expressed as a negative
logarithm. pH = -Log{H+} where { } denotes activity. For fresh waters, { } � [ ] where [ ] denotes molar concentration.
· pH measurements are typically given on a scale of 0.0-14.0 (www.epa.gov), but values outside of this range can exist.
· For each 1 unit change in pH, the alkalinity or acidity changes by a factor of 10 (i.e., pH of 5 is 10 times more acidic than 6)
Many species of life have trouble surviving if the pH is lower than 5 or higher than 9. Small shifts in the pH of water can affect the solubility of certain metals such as iron and copper, which indirectly affects the aquatic organisms by releasing metals that were present in the lake’s sediments. Acids, such as from acid rain can diminish the survival of fish eggs.
Some other factors that can change the pH in a body of water as follows:
· Sewage Overflows · Chemicals in runoff · Bacterial Activity · Acid drainage from coal mines, accidental spills, and acid rain. · Water turbulence (mixing chemicals from the sediment into the water column) · Photosynthesis by aquatic plants. · Algal blooms caused by an overload of nutrients
The objectives of this experiment are to: 1) determine the strength of an acid by a "conservation of equivalents" method, 2) measure the alkalinities of several water samples and to see how they vary, and 3) develop the titration curves for several water samples to examine the buffering characteristics of these
waters and see the relationship between buffering and alkalinity.
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Chemical Equilibrium, Carbonate System, and Alkalinity
Chemical equilibrium is the condition within the course of a reversible reaction whereupon there is no net change in the amounts of reactants or products. Reversible reactions are reactions where the products, as soon as they have formed, react to produce the original reactants. Therefore, in chemical equilibrium two opposing reactions occur at equal rates. Such equilibrium may be represented by the qualitative formulation,
DCBA +Û+
While the reaction rates between A+B and C+D are equal during chemical equilibrium, the concentrations may or may not be. This concentration differential may be accounted for by a coefficient called the equilibrium constant (Ke). Consider the following reaction,
iA + jB � mC + nD
where the differential concentrations of each reactant are represented by the superscripts i, j, m and n. The reaction constant may be determined by the following relationship,
eji
nm
K BA DC
=
This implies that for a given reaction, the equilibrium constant is proportional to the amount of products and reactants present at equilibrium. This constant, however, is not the same for all reactions and is dependent upon the pressure and temperature according to LeChatelier’s principle. LeChatelier’s principle states that if a system in a balanced, or equilibrium, state is disturbed, the system will readjust itself so as to neutralize the disturbance and restore equilibrium. Consider the reaction in which the stoichiometric coefficients are all equal to 1 and the chemical system has reached equilibrium,
A + B � C + D
If more of the reactant C is added to the balanced system, the opportunity for the reaction of C and D increases. As the reaction between C and D increases, the more of A and B that is produced. This increase in the reaction of C and D offsets the original condition that disturbed the equilibrium by removing the additional C and producing more of A and B. In order that Ke remain constant, D must decline while A and B increase. This process continues until equilibrium in the system is reached.
Closely related to chemical equilibrium is the process of buffering. Buffering describes the process by which pH changes very little when an acid or base is added to water. Despite the addition of more H+ (acid) the concentration of acid in the sample remains constant. This process is accomplished by the addition of a chemical agent known as a buffer. A solution of acetic acid (CH3COOH) and sodium acetate is a commonly used buffering agent.
In natural waters, the reaction responsible for buffering is,
-+ += 332 HCOHCOH
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The bicarbonate ion can further dissociate into a carbonate and hydrogen ion,
-+- += 233 COHHCO
When acid is added to the water most of the H+ reacts with HCO3- to make H2CO3. If the concentration of CO32- remains high in comparison to added H+, the pH will remain nearly constant.
The amount of acid that may be added to a solution without the pH value dropping below 4.5 is known as the alkalinity. Alkalinity is most often referred to as the ability of solution to neutralize acid and may be determined by summing the molar concentrations of the agents that react with the H+. For the example reaction above,
Alkalinity = (1)[HCO3-] + (2)[CO32-] + (1)[OH-] – (1)[H+]
where the numbers in () are the numbers of H+ that can be reacted with each ion and the numbers in [ ] are the molar concentrations. Using the units only,
Alkalinity = (1 or 2 eq/mole) x [mole/L] = eq/L = N (normality)
Alkalinity is often expressed in the units of meq/L (milliequivalents per liter) or mg/L as CaCO3.
Alkalinity as CaCO3 = (Alkalinity in meq/L) x (50 mg as CaCO3 / meq )
The calculation of alkalinity in the lab may be accomplished by using an equivalents balance,
(Nalkalinity)(Vwater) = (Nacid)(Vacid added)
CONCLUSION The pH of water is critical to the survival of most aquatic plants and animals. It is a chemical parameter that increases our understanding of the water’s health. It is quick and easy and establishes a baseline so that unexpected water quality changes can be understood. The amount of acid that may be added to a solution without the pH dropping below 4.5 is known as the alkalinity. Alkalinity is most often referred to as the ability of solution to neutralize acid and may be determined by summing the molar concentrations of the agents that react with the H+. Alkalinity is expressed in the units of meq/L (milliequivalents per liter) or mg/L of CaCO3. The calculation of alkalinity in the lab may be accomplished by using an equivalents balance.
The objective of this experiment is to (1) determine the strength of an acid by a "conservation of equivalents" method, (2) measure the alkalinities of several water samples and to see how they vary, and (3) develop the titration curves for several water samples to examine the buffering characteristics of these waters and see the relationship between buffering and alkalinity.
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LAB 7: PROCEDURES (Alkalinity and Buffering) Objectives 1. To determine the strength of an acid by a "conservation of equivalents" method. 2. To measure the alkalinities of several water samples and to see how they vary. 3. To develop the titration curves for several water samples to examine the buffering characteristics of
these waters and see the relationship between buffering and alkalinity.
Procedures Part A: Determining the Strength of an Acid Equipment and Materials:
q Primary standard solution (provided by the instructor) q Sulfuric acid solution (made by instructor) q Buret and stand (Please be careful with these; they are expensive.) q 250-mL beaker or Erlenmeyer flask q Methyl orange indicator q 50-mL graduated cylinder q 25-mL volumetric pipet
1. Obtain from your instructor 20 mL of a base solution made with sodium carbonate (Na2CO3). This is called a "primary standard".
2. Place 2 drops of Methyl-Orange indicator (pH 4.5) solution into the sample. 3. Fill your buret with the acid (titrant) using a small beaker. (You don't have to fill the buret to any
particular mark.) 4. Note the beginning mark on the buret. (Remember to read the bottom of the meniscus.) 5. Carefully, add titrant to the sample while continuously swirling. When the color shifts from a light
yellow to a light orangish-pink, the endpoint has been reached. Stop adding acid at that point. 6. Note the ending mark on the buret. Calculate the volume of titrant used. 7. Calculate the normality of your acid from: (NV)standard = (NV)acid. Include this in your sample
calculations. 8. Collect the Nacid values from each lab group. Calculate the average and use that value for the
alkalinity testing.
Sample format for recording primary standard data in your lab notebook:
Normality of the primary standard solution (meq/L)
Titration Data for Your Group: Volume of
Primary Standard solution (mL)
Initial Buret Marking (mL)
Final Buret Marking (mL)
Titrant Volume (mL)
Normality of acid (meq/L)
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Class Data (cross out any values not used in the calculation of the average): Group Acid Normality (meq/L)
1
2
3
4
5
Average
Part B: Measuring Total Alkalinity 1. Measure the pH's of the water samples provided by the instructor. 2. Put 20 or 25 mL of sample into a beaker or Erlenmeyer flask. (Measure this volume carefully.) 3. Place two drops of Methyl-Orange indicator (pH 4.5) solution into the sample. 4. Titrate with your standardized acid (in the buret) until the methyl orange changes color. 5. Note the ending mark on the buret. Calculate the volume of titrant used. 6. Calculate the total alkalinity of each sample from: (NV)Titrant = (alk)(VSample). Express your answer in
both meq/L and mg/L as CaCO3. Include one of these calculations in your sample calculations.
Sample format for recording data in your lab notebook: Water Samples
Label Description DI Deionized water SW Surface water GW Groundwater DI + BS Deionized water + baking soda (NaHCO3). Amount of NaHCO3 added/ L: g/L SP Soda pop
Titration Data
Sample Initial pH Sample
Volume (mL) Initial Titrant Buret Mark
(mL)
Final Titrant Buret Mark
(mL)
Titrant Volume (mL)
DI SW GW DI + BS SP
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Alkalinity Calculations
Sample Sample
Volume (mL) Titrant
Volume (mL) Titrant
Normality (meq/L)
Total Alkalinity (meq/L)
Total Alkalinity (mg/L as CaCO3)
DI SW GW DI + BS SP
Part C: Titration Curves 1. Your instructor will set up a pH meter, buret and water sample. 2. Add acid to 100 mL of sample in very small increments, recording the pH and the amount of titrant
added until the pH is less than 4.5. 3. Repeat with other water samples as directed by the instructor. 4. Return unused titrant in the buret to the container that it came from. Data Analysis Part B: Calculate the alkalinities of the water samples tested and express your answers in both meq/L and
mg/L as CaCO3 units. Summarize these in a table. Discuss the meaning of the results; are they what one would expect? Why?
Part C: Using a spreadsheet, graph pH on the y-axis vs. cumulative meq. of acid added per liter of water sample on the x-axis. Plot all titration curves on the same graph. Properly title and label your graphs. Discuss the meaning of the results; are they what one would expect? Why? Make sure that you include a discussion of each water’s buffering capacity.
Discussion Questions (20% of the report total score) 1. (2.5%) Why didn't the soda pop have any alkalinity? (Consider the form of the carbonate species at
this pH.) 2. (2.5%) Why did adding baking soda to the deionized water increase its alkalinity? 3. (7.5%) Was the resulting alkalinity of the DI+BS sample about what you expected? Show
calculations for what you expected. 4. (7.5%) A factory wishes to discharge an acidic waste to a local stream. State regulators want to assure
that the local fish population is not harmed by too drastic a change in pH. Your job is to estimate how much wastewater can be discharged (in Liters/min) so that the pH of the stream downstream doesn't change by more than 1.0 pH units. The proposed discharge has a concentration of 5.50 N. The stream has a flow rate of 30 cfs and a titration curve like that found in the groundwater.
Lab Write-up Report your alkalinity data and the results of your calculations in your lab notebook. Provide sample calculations for every different kind of calculation done. Be sure to check for appropriate significant figures. Present your titration curve data and graphs on print-outs from a spreadsheet program. Provide sample calculations for the calculations done on spreadsheets also. Show the calculations. Answer the discussion questions.
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LAB 7: NOTES ON CHEMICAL EQUILIBRIUM, BUFFERING, AND ALKALINITY
Chemical reactions often go in two directions (reversible reactions):
A + B à C + D (forward), or A + B ß C + D (backward)
We express this on paper as: A + B Û C + D, or in shorthand: A + B = C + D. For example, 2SO2(g) + O2(g) à 2SO3(g). But some SO3(g) breaks down, so we have at the same time 2SO2(g) + O2(g) ß 2SO3(g)
When the rate of the forward reaction equals the rate of the reverse reaction, the system is in "dynamic equilibrium". The word "equilibrium" means the relative concentrations of reactants and products do not change with time. "Dynamic" reminds us that the reactions are still going on.
Equilibrium Constants Just because the reaction rates are equal, doesn't mean the concentrations are equal. In the example above, the equilibrium concentrations are:
[SO2] = 1.65 M (M = molarity = moles/L) [SO3] = 12.43 M
In general, for the reaction:
iA + jB = mC + nD
it has been empirically determined that the equilibrium concentrations are related by the following relationship:
eqji
nm
K BA DC
= ][][ ][][
(Remember: "products over reactants")
Note that Keq is usually determined by experiment, and is affected by various environmental factors, most notably, temperature.
LeChatelier's Principle or the Law of Mass Action If a chemical reaction at equilibrium is subjected to a change in conditions that moves it away from equilibrium, then the reaction rates adjust so that a new equilibrium state is reached. In other words, the system adjusts in a way that, at least partially, offsets the condition that disrupted the equilibrium. This adjustment continues until a new equilibrium is established. For example, suppose a system is described by the following equilibrium equation:
A + B = C + D
Now, suppose we add more "C" from some outside source (say, from a salt that contains C). The increased concentration of C increases the opportunities for C and D to react. The additional C will increase the rate of the reverse reaction and some of the additional C will be converted to A and B. (This is the action that offsets the condition that disturbed the original equilibrium. The increased rate of
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disappearance of C counters our action of adding C to the solution.) If more C and D react, the concentration of C and D decline and the concentrations of A and B increase. This process continues until the equilibrium equation is satisfied again. Mathematically, if we artificially increase C, then for the Keq equation to hold true, D must decline and A and B must increase. This coincides with the chemical changes that occur.
Buffering "Buffering" describes that phenomenon in which pH changes only a little when an acid or base is added to a water. In other words, despite adding H+ in the form of an acid, the concentration of acid in the water (measured by pH) doesn't change accordingly. It stays relatively constant. In most natural waters, the reaction responsible for this phenomenon is:
H2CO3 = H+ + HCO3- (where H2CO3 = carbonic acid and HCO3- = bicarbonate ion)
The bicarbonate ion can further dissociate into carbonate ion and H+ as follows:
HCO3- = H+ + CO32-
So, when acid is added, most of the H+ added reacts with HCO3- to make H2CO3. (The H+ would react with any CO32- to make HCO3-, but at the pH values of most natural waters, the amount of CO32- is very small compared to the concentration of HCO3-.) As a result of this reaction which "sucks up" the added H+, the concentration of H+ in the water doesn't change much, and the pH stays relatively constant. This is buffering.
Alkalinity In the system described above, pH stays relatively constant until all of the bicarbonate and carbonate ion are used up. At that point, the pH changes fairly quickly. The amount of acid that can be added until this point is called the alkalinity of the water. Alkalinity is usually defined as the ability of the water to neutralize acid. It can be calculated by summing the molar concentrations of the species that react with H+.
Alkalinity = (1)[ HCO3-] + (2)[ CO32-] + (1)[OH-] - (1)[H+]
The numbers in () are the numbers of H+ that can be reacted with each ion. The numbers in [ ] are the molar concentrations. If the equation is written out in units, we get:
Alkalinity = (1 or 2 eq/mole) x [mole/L] = eq/L = N (normality)
Alkalinity is expressed in units of meq/L (milliequivalents per liter) or mg/L as CaCO3.
Alkalinity as CaCO3 = (Alkalinity in meq/L) x (50 mg as CaCO3 / meq )
We will measure alkalinity in the lab by adding acid to a water sample until the pH starts to change (at about pH 4.5). The amount of acid added (in meq per L of sample water) is the amount that can be neutralized. This is analogous to measuring a hole in the ground by measuring the volume of water needed to fill it.
The calculation of the "amount" of acid is based on an "equivalents balance":
(Nalkalinity) x (Vwater sample) = (Nacid) x (Vacid added by titration) Units: (eq/L) x (L) = eq
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LAB 8: HARDNESS INTRODUCTION: Hardness is the combined concentration of multivalent cations in terms of equivalents per liter or milligrams per liter as calcium carbonate. Calcium (Ca+2) and magnesium (Mg+2) ions mostly contribute to the hardness in natural waters. Hardness also affects the way soap acts. In hard water, the cations form precipitates with some of the materials in natural soaps. This reduces the amount of lather and can actually feel gritty (hence the name hard water). Whereas, soft water often feels slimy to consumers. Other ways that hardness affects water are described further below.
The principle behind measuring water hardness is rather simple. An indicator, e.g., an organic dye called Eriochrome Black T (EBT) or alternative indicator is added to the water sample. Despite normally being blue, the EBT turns the water a red-purple (or pink) color. This is caused by a process known as complexation where the EBT binds to magnesium ions within the water. This provides us with a means to measure (indirectly) the amount of hardness forming ions in the water sample. This lab will demonstrate these measurements and determine the water hardness of several samples of water (measured in mg/L as CaCO3).
BACKGROUND: Water hardness can usually be defined as the concentration of calcium and magnesium ions found within a given water sample, usually expressed in terms of calcium carbonate (CaCO3). The hardness of a water may be calculated using the following equation,
)/2])([]([ 22 moleeqMgCaHardness ++ +=
Or, in terms of CaCO3, )/2])([]([ 22 moleeqMgCaHardness ++ += (50,000 mg CaCO3 per eq.) where, [Ca2+] = molar concentration of calcium
[Mg2+] = molar concentration of magnesium
The most frequently used standard classification of water hardness is shown below.
Water supply classification
Degree of Hardness
Concentration
meq/ L mg/L as CaCO3 grains/ gallon Soft Water 0.00 to 0.34 0.00 to 17.1 0.00 to 1.00
Slightly Hard Water 0.35 to 1.00 17.2 to 60.0 1.01 to 3.50
Moderately Hard Water 1.01 to 2.40 60.1 to 120 3.51 to 7.00
Hard Water 2.41 to 3.60 121 to 180 7.01 to 10.5
Very Hard Water Over 3.60 Over 180 over 10.5
While hardness is not generally considered a health concern, it does pose several everyday problems. These problems stem from the formation of calcium and magnesium precipitates, the process that produces cave formations. Such precipitates, though beautiful in their natural form, can wreak havoc on
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engineering systems. For example, the formation of these precipitates on the walls of pipes (often called scale) increases pipe friction and, in extreme cases, may completely block the pipe. In other cases, scale may prevent valves from sealing properly or may even prevent them from working at all.
Of particular concern is the formation of scale in hot water systems. Because calcium carbonate, CaCO3 (the precipitate form), becomes less soluble as temperature increases, hot water systems are especially prone to the formation of scale. Scale within a boiler can prevent the transfer of heat between the exchange tubes and the water by acting as an insulator. In some cases, scale may completely block the piping within the boiler, leading to expensive replacement costs.
Another problem associated with water hardness is the inability to form lather with soap, as discussed above. In hard water, cations form precipitates with some of the minerals of natural soaps. This reduces the amount of lather the soap will produce. Hard water may also produce a deposit called “soap curd” that may remain on the skin causing irritation or infection. “Soap curd” can also pick up dirt from dirty laundry water and hold it on clothes, making it more difficult to get them clean.
Table 1 (below) shows the main cations that cause hardness in water and the major anions associated with them.
Table 1 Cations causing hardness and the major anions associated with them
Cations causing hardness Anions typically with them
(These do not cause hardness!) Ca2+ HCO3-
Mg2+ SO42-
Sr2+ Cl-
Fe2+ NO3-
Mn2+ SiO32-
CONCLUSION
Water hardness can usually be defined as the concentration of calcium and magnesium ions found within a given water sample, usually expressed in terms of calcium carbonate (CaCO3). Although hardness is not generally a health concern, it is important with regards to the formation of precipitates because of the problems that they can cause.
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LAB 8: PROCEDURES (Hardness)
Objectives The purpose of this lab is to measure the total hardness of several different waters. Determining the Normality of the EDTA Titrant Measure the normality of the EDTA titrant solution (in mg/L as CaCO3) using a standard CaCO3 solution (either 0.10 mg CaCO3 per mL or1.00 mg CaCO3 per mL, as directed by the instructor). 1. Using a pipet, place 50 mL of the standard 100 mg/L CaCO3 solution into a 100 or 250 mL
Erlenmeyer flask. 2. Add about 2 mL of Hach hardness 1 buffer solution, swirl to mix. 3. Then add ManVer 2 Hardness Indicator Powder Pillow, swirl to mix. 4. Titrate with EDTA titrant while swirling the flask until the solution color begins to change from
red-wine to blue. At that point add a few drops at 3 to 5 second intervals while swirling the flask. Note the EDTA titrant level on the buret when the titration is complete.
5. Calculate the normality of EDTA titrant as follows:
ா்ܰ = (0.100 ݉݃ ଷܱܥܽܥ ݎ݁ )(ܮ݉ ௦ܸ௧ௗௗ)
ா்ܸ ݀݁ݐܽݎݐ݅ݐ where: NEDTA = Normality of the EDTA titrant (mg CaCO3 per mL) VCa-Standard = Volume of the CaCO3 standard used (mL) VEDTA titrated = Volume of the EDTA titrant (mL)
6. Record your results. 7. Pool the class results and calculate an average value for the normality of EDTA titrant. Hardness Measurements 1. Using a graduated cylinder, measure out a 50-mL sample of one of the test waters. 2. Place the sample in a 100 or 250 Erlenmeyer flask. Add 2 mL of the Hach hardness 1 buffer solution. 3. Then add ManVer 2 Hardness Indicator Powder Pillow, swirl to mix. 4. Titrate to the blue endpoint of the EBT as you did in Part A. 5. Record your results.
6. Calculate hardness as follows:
where: H = Total hardness of the sample (mg/L as CaCO3) VEDTA = Volume of the EDTA titrant added (mL) NEDTA = Normality of the EDTA titrant (mg CaCO3 per mL) Vs = Sample volume (mL)
7. Repeat the measurement for other water samples provided by the instructor. 8. Return unused titrant in the buret to the container that it came from.
References
1. Sawyer, Clair N., McCarty, Perry L., and Gene F. Parkin. Chemistry for Environmental Engineering. 4th ed. McGraw-Hill, Inc., 1994.
2. Vesilind, P. Aarne. Introduction to Environmental Engineering. Boston: PWS Publishing Company, 1997.
s
EDTAEDTA
V
) mg/mL mg/L
)(1000)(N(V H =
California State University, Fresno Department of Civil & Geomatics Engineering
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LAB 8: NOTES ON HARDNESS Definition of Hardness
Hardness is the combined concentration of multivalent cations in terms of equivalents per liter or milligrams per liter as calcium carbonate. In natural waters, the vast majority of hardness is contributed by calcium (Ca+2) and magnesium (Mg+2) ions. One can calculate the hardness from the calcium and magnesium concentrations as follows:
Hardness (eq/L) = {[Ca2+] + [Mg2+]}x(2 eq/mole)
Hardness (mg/L as CaCO3) = {[Ca2+] + [Mg2+]}x(2 eq/mole)x(50,000 mg CaCO3 per eq)
where [Ca2+] and [Mg2+] are the molar concentrations of calcium and magnesium
Importance of Hardness
Hardness is not generally a health concern, except in very high concentrations. Calcium and magnesium do, however, form precipitates (CaCO3 and Mg(OH)2) relatively easily. Cave formations, for example, are mostly CaCO3. In pipe systems, these precipitates show up as "scale" on faucets and glass shower doors. Inside pipes, scale formation increases pipe friction and reduces the pipe cross section. In the extreme case, scale can completely block pipes. Scale formation in valves prevents them from sealing properly. Unlike most salts, CaCO3 becomes less soluble as temperature increases. Thus, scale formation is a particular problem in hot water systems. Scale formation inside boiler heat exchange tubes reduces the heat transfer efficiency by insulating the tube on the inside, and can block the tubes completely. Scale also forms in home hot water heaters. Besides temperature, pH is an important factor in scale formation. (Think about the carbonate system and what you know about solubility product.)
Besides scale, consumers experience hardness in the way soap acts. In hard water, the cations form precipitates with some of the materials in natural soaps. This reduces the amount of lather (suds) and can, in the extreme, actually feel "gritty" (hence the name "hard" water). On the other hand, water that is too soft feels "slimy" to consumers, who often feel that they haven't rinsed completely when the water is too soft.
Principle Behind the Total Hardness Measurement Technique
An indicator, the organic dye Eriochrome Black T (EBT), is added to a water sample with calcium and magnesium ions. Normally, EBT is a blue color. In the water sample, magnesium ions (part of the hardness) bind to the EBT indicator in a process known as complexation. The Mg-EBT complex is a red- purple color. Now, the water sample is titrated with a solution containing the organic compound EDTA (Ethylenediaminetetraacetic acid). EDTA is a stronger complexing agent than EBT. The EDTA will first form complexes with all the free multivalent cations in the sample. When they are gone, the EDTA will then "steal" Mg ions from the Mg-EBT complexes. When this happens, the EBT reverts to its original blue color. This color change indicates the endpoint of the titration. The equivalents of EDTA added equals the equivalents of multivalent cations in the water.
California State University, Fresno Department of Civil & Geomatics Engineering
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LAB 9: CADILLAC DESERT: WATER & TRANSFORMATION OF NATURE Part I: Mulhollands’s Dream
Focus Issue: Growth and Sustainability Video: View the video at “Video Links” on the course Blackboard site main menu.
Mulholland’s Dream demonstrates how our desire to live in certain areas may conflict with the sustainable amount of natural resources in that area. When faced with the problem, we have devised many ways to transport water to where we want to live. But at what cost?
Discussion Points: · What are the consequences of, and solutions to, regional population growth and the increasing
competition for water resources? Should we limit development in certain areas, like Central Valley? Make connection to the current drought and state water emergency situation.
· Should water-rich communities share their resources with water poor communities? How would you feel as a member of a water-rich community? What would it be like for you if you were part of the other community?
· Is it ethical to move water from one location to provide for needs of a different location? Or is it ethical to limit the water for its original location? Who should determine when we should take such an action: Local communities, state agencies, or the federal government?
Research Activities: · Research and locate the source of water for your city or county. Is it rivers, lakes or aquifers? Is
the water brought in from somewhere else? Has that source changed over the years? What is the status of the water supply?
· Research how much water your city or county uses per year, and how much is used per person or per family, per day or per year. Is that average going up or down? If the population is rising in your area, what does that mean of future water supplies?
· Research the background of the water crisis at west side of Central Valley, what caused it, and what is the possible solution?
· Research the advantages and disadvantages of a large water project such as dams and aqueducts. What are the benefits of leaving a river in its natural state, as opposed to diverting the water for agriculture and urban development? Is compromise possible? If so, how?
Deliverables: A research paper with a minimum of 1500 words covering one or two discussion points listed above. Use research to support the arguments.
References: Cadillac Desert: Water and the Transformation of Nature. A discussion and Viewers Guide to the PBS Series. CSUF Maddan Library, Music & Media Library. Media HD 1694 A5 C33 1997
Reisner, M. Cadillac Desert: The American West and Its Disappearing Water. New York: Penguin Books, 1997.
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GO BULLDOGS!!!
- LAB 1: WATER CHEMISTRY CALCULATIONS
- LAB 2: GAS TRANSFER
- Procedures
- LAB 3: BIOCHEMICAL OXYGEN DEMAND
- LAB 4: ADSORPTION
- LAB 5: SOLIDS IN WATER
- Purpose of Lab
- LAB 6: CHEMICAL COAGULATION JAR TESTING
- LAB 7: CHEMICAL EQUILIBRIUM, BUFFERING, AND ALKALINITY
- LAB 8: HARDNESS
- References
- LAB 9: CADILLAC DESERT: WATER & TRANSFORMATION OF NATURE
