answer prelab question

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Abundant evidence of consistency and individual work.Uses Manual and resources posted on the course’s Canvas.All questions are addressed, and responses are given in a clear, complete manner that reveals the reasoning process (DESCRIBE or EXPLAIN).All calculations show dimensional analysis.Responses to SDSs questions are summarized in a manner that reveals understanding of relevance and context.

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ENV 3001L
Environmental Engineering
Laboratory Manual
Spring 2024
Prepared by
Anna R. Bernardo-Bricker, Ph.D.
Department of Civil and Environmental Engineering
Florida International University
Table of Content
Preface
………………………………………………………………………………………………………………………..3
I.
Accuracy, Precision, Errors, and Uncertainty …………………………………………………………4
II.
Tolerances of Laboratory Volumetric Measuring Tools …………………………………………..6
III.
Error Propagation in Arithmetic Calculations…………………………………………………………8
IV.
Categories of substances most used in ENV 3001L ……………………………………………….11
EXPERIMENT #1:
Solutions and Concentration – concepts and practical aspects ………………………………12
EXPERIMENT #2:
pH definition and measurement ………………………………………………………………………..16
EXPERIMENT #3:
Acid-Base Neutralization Titration ……………………………………………………………………..19
EXPERIMENT #4:
Determination of Total Hardness in water by complexometric titration with EDTA ….25
EXPERIMENT #5:
Determination of different forms of alkalinity in water samples by titration with a
standard acid solution ………………………………………………………………………………………30
EXPERIMENT #6:
Conductivity as a measure of Total Dissolved Solids in water ………………………………..34
EXPERIMENT #7:
Biochemical Oxygen Demand (BOD) Test ……………………………………………………………40
EXPERIMENT #8:
Turbidity and Surface Water Treatment: coagulation flocculation-sedimentation
treatment processes ………………………………………………………………………………………..44
EXPERIMENT #9:
Adsorption of organic substances from aqueous solutions onto various types of
activated carbon. ……………………………………………………………………………………………..50
Appendix for EXPERIMENT #1: An experiment designed to study the relationships of concentration to
solute and solvent amounts. ……………………………………………………………………………..54
ENV 3001L: Environmental Engineering Laboratory: Spring 2024
FLORIDA INTERNATIONAL UNIVERSITY
Page | 2
Preface
What knowledge, skills, and attitudes will students develop in this course?
Laboratory courses are a central aspect of education in science and engineering disciplines. These learning
environments promote hands-on and experiential knowledge. The goals of the ENV 3001L lab include
furthering students’ development of skills in four domains:
1. Field-of-study: The focus of the ENV 3001L laboratory course is to support learning of
environmental chemistry and environmental analysis concepts and procedures.
2. Practical skills: Handling equipment and chemicals, learning safe laboratory practices, measuring
accurately, and learning specific techniques and standard procedures for analysis of environmental
samples.
3. Logical skills: Observing carefully, practicing deduction and interpretation, learning how to plan
experiments that offer insights into a given experiment or concept.
4. Soft skills: Team working, reporting, discussing, time management, and problem-solving.
The laboratory experiments support the co-requisite course ENV 3001 not only by illustrating ideas and
concepts but also by helping students develop these four skills as aspects of critical thinking. The
instructional approach used is a combination of guided-inquiry and expository practices. Therefore, this
manual is not designed to be used as a cookbook.
What this Manual is not
This Manual is not intended to be used as a “cookbook”. For practicality, procedures for most of the
experiments are given; however, students are required to engage in guided-inquiry activities, and be aware
of procedure’s modifications specified by the Instructor during in-person and online class meetings.
What this Manual contains
Four sections:
Content in the first three sections should be applied to the analysis of data from each Experiment.
I. Accuracy, Precision, Errors, and Uncertainty
II. Tolerances of Laboratory Volumetric Measuring Tools
III. Error Propagation in Arithmetic Calculations
Next, Experiments description:
Each Experiment description starts with a list of concepts and definitions and extensive background that
covers main aspects of the lab experiment as well as equipment.
A detailed procedure is given for some experiments, while for others a simple experimental design is
part of the learning objectives.
ENV 3001L: Environmental Engineering Laboratory: Spring 2024
FLORIDA INTERNATIONAL UNIVERSITY
Page | 3
I.
Accuracy, Precision, Errors, and Uncertainty
These terms are often used interchangeably; however, they have different meanings. The
uncertainty is a quantitative indication of the quality of a result. All measurements of physical quantities
are subject to uncertainties in the measurements. At the very least governed by what can be achieved with
the quality of the equipment or instrument used to make the measurement; for example, see Section II in
this Manual for the uncertainty of glassware. Additionally, other sources of variability, called errors, can
affect a measurement.
Let us first start by defining the terms accuracy and precision. The bull’s-eye analogy helps visualize
the meaning of these two terms. Accuracy is how close you get to the bull’s-eye, whereas precision is how
close the repetitive shots are to one another.
Accuracy is formally defined as the degree of agreement or difference of a measured value
compared to the “true,” “expected,” “accepted”, or “reference” value of the parameter. We’ll call these
the reference value. The “relative difference” between measured and actual values as defined by:
=
−

=
Equation I.1
| − |
Equation I.2

These differences can also be expressed as a percentage:
=
−
× 100

=
| − |

Equation I.3
× 100
Equation I.4
Precision is the term used to refer to the degree of mutual agreement among individual
measurements (x1, x2,… xi) as a result of repeated applications under the same condition. Precision
measures the variation among measurements and may be expressed in different ways. If one has three or
more measurements (n≥3), the standard deviation can be used to communicate precision of the data. For
n 7
(See Equation 3.3)
Weak Base + Strong Acid Salt (subject to hydrolysis) + Water
Salt pH < 7 Weak Base + Weak Acid Salt (subject to hydrolysis) + Water can’t predict; depends on Ka and Kb That is, if the salt AB is the product of the reaction between a weak acid and a strong base then the conjugate base of the salt, A-, will react with water to regenerate the acid according to Equation 3.3. − −( ) + 2 ( ) ⇔ ( ) + ( ) ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Equation 3.3 (if A- is from a weak acid) Page | 21 Whether the acid is strong or weak, in a titration analysis, the balanced chemical equation (Equation 3.1) establishes that at the equivalence point, when enough moles of BOH has been added to completely consume the HA initially present in the sample analyzed, Equation 3.4 must be valid. Further, the molar concentration of the analyte can be calculated by dividing the moles of the analyte by the volume of solution in which those moles were originally contained (Equation 3.5). × ℎ × ( ) • 1 1 = = Equation 3.4 Equation 3.5 Finding the equivalence point in a neutralization titration curve The Standard Method APHA 2310 (APHA, 2005) describes using the first derivative of the titration curve to identify the equivalence points. The first derivative of a function is the slope of the tangent line for any point on the function. This analysis clearly shows when the function is increasing, decreasing, or where it has a horizontal tangent. For the titration curve function, the first derivative is a plot of ΔpH/ΔVtitrant in the y-axis vs. Vtitrant in the x-axis. The equivalence points of a titration are revealed as local maximum in the first derivative plot. • Stoichiometry of the titration and the Henderson-Hasselbach approximation This section explains the strategy that students will use to find the pKa of the monoprotic weak acid. This strategy consists of two overall steps: stoichiometric calculations of the titration, followed by applying the Henderson-Hasselbach equation. Here we start by explaining first the Henderson-Hasselbach equation. The Henderson-Hasselbach equation (Equation 3.9) is derived starting from the equilibrium of the weak acid in water, Equation 3.6 as follows: + ( ) ⇔ −( ) + ( ) ≅ Rewriting Equation 3.7: Taking log10 on both sides: Equation 3.6 + � − ( ) �� ( ) � Equation 3.7 � ( ) � + [ ( ) ]≅ ×� ( ) � � − ( ) � � ( ) � + − 10 � ( ) � ≅ − 10 − 10 � ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY � − ( ) � � Page | 22 Assuming formal concentrations of species in are equal to their equilibrium concentration [ ]: ( ) + − 10 � ( ) � ≅ − 10 − 10 � = − 10 ( ) Defining: − ( ) � Equation 3.8 ( ) We arrive at the Henderson-Hasselbach equation: ≅ − 10 � − ( ) � Equation 3.9 Students can apply Equation 3.9 to the titration data in two different ways to figure out Ka for the weak acid. First; however, −( ) and ( ) need to be determined. Before the equivalence point of the titration, the concentrations of the remaining acid, ( ) , and the produced salt, −( ) , in units of mol/L (M) after each single addition of an amount of titrant, BOH, can be found by applying the following stoichiometric relationships, all based on Equation 3.1: = × × 1 1 = − = ( ) − = × × − = − ( ) 1 − 1 Equation 3.10 Equation 3.11 Equation 3.12 Equation 3.13 Equation 3.14 Materials and Procedures Reagents: • • • Secondary standard Sodium Hydroxide (NaOH) or Potassium Hydroxide (KOH) solution, commercial Strong acid solution of unknown concentration Acetic Acid (CH3COOH) solution of unknown concentration Equipment (partial list): • • • Magnetic stirring plate Micro stir bar (~7 mm L x 2 mm Dia) pH meter ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 23 Figure 3.1 Sketch of a set up for a pH-meter monitored Titration (potentiometric titration) Procedure: 1. Fill a burette with the standardized strong base solution (titrant). 2. Using a graduated cylinder, measure a specified volume of the acid and place it into a 250 mL beaker. 3. Arrange the beaker/stirring plate/burette/pH meter in a titration setup as shown in Figure 3.1. 4. Place the micro stir bar inside the beaker and set the magnetic stirring to a speed such that a constant mixing is maintained, but splashing the sample or knocking the electrode is avoided. 5. Record the initial pH of the sample. 6. Add 1-3 drops of the appropriate indicator, and record color. 7. Record the initial pH of the sample plus the indicator. 8. Start to record the data required to produce the titration curve. Record both the burette reading and resulting pH after each addition of titrant. Take enough measurements at appropriate increments of the titrant volume. Appropriate volume increments are those that allow drawing a well-defined curve. When the recorded data shows the pH value beginning to drop more quickly, add the titrant at the smallest possible increments. NOTE: As you carry out the titration, make a note of the point when the indicator changes color, and the final color of the indicator. 9. When the pH values start to taper off, again add larger increments of the strong base titrant. Continue for two or three more additions or until your data clearly shows that the pH has even out. Literature Cited: 1) American Public Health Association (APHA), 2005. Method 2310, in “Standard method for examination of water and wastewater”, 21st Edition. APHA, AWWA, WPCF, Washington, DC. 2) Masters, Gilbert M., Ella, Wendell P. 2008. Introduction to Environmental Engineering and Science. Prentice Hall, Upper Saddle River, NJ, U.S.A. ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 24 EXPERIMENT #4: Determination of Total Hardness in water by complexometric titration with EDTA Learning Objectives: • • • • Determine the total hardness of prepared solutions containing Ca+2 and Mg+2 and in various samples of water (maybe drinking water or environmental water samples). Compare the total hardness in the analyzed samples. Study the effect of one of the independent or control variables on the determination of hardness. Explain the concept of formation and stability of metal-complexes and its application for the quantitative determination of hardness in water. Key Concepts and Definitions Recall the concepts and definitions from pages 19 and 20: • Analyte • Standard Solution • Primary standards • Titration • Equivalence point of a titration • Endpoint of a titration Total Hardness (TH): Defined in the Method APHA 2340 as the sum of dissolved calcium (Ca+2) and magnesium (Mg+2) concentrations, both expressed as calcium carbonate (CaCO3); TH in mg/L as CaCO3. This titration Method of analysis does not distinguish between Ca2+ and Mg2+. Coordinate bond: a covalent bond that consists of a pair of electrons supplied (donated) by only one of the two atoms it joins. The anions or molecules supplying the electrons are called “ligands”. Simple ligands •• •• •• −1 include water ( 2 ), ammonia ( 3 ) and chloride ions (� •• •• � ). •• Coordination metal complex (chelate complex): A chemical compound formed by a coordination bond between a metal ion and a molecular or ionic substance, called ligand, that contains at least one atom with an unshared pair of electrons. These complexes can be neutral or charged. When the complex is charged, it is stabilized by neighboring counter-ions. Coordination metal complexes are also called chelate complexes (chela=pincer-like claw). Theoretical Background • Principle of the EDTA Titrimetric Method Ethylenediaminetetraacetic acid (EDTA) has a total of six binding places to form metal complexes. The fully protonated form of EDTA, Figure 4.1 a, is a hexaprotic weak acid with successive pKa values of pKa1=0.0, pKa2=1.5, pKa3=2.0, pKa4=2.66, pKa5=6.16, and pKa6=10.24. The first four values are for the carboxylic acid protons (-OH) and the last two values are for the ammonium protons (-NH) (Harvey, D., 2019). This means that only at pHs ≥ 10.24 the EDTA molecule would be fully deprotonated and, thus, be able to form chelate complexes with metals divalent cations, M+2, according to the stoichiometry shown in ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 25 Equation 4.1, and illustrated in Figure 4.1 b. The number of “chelate” coordination sites depends on the size of the metal ion, however, all metal–EDTA complexes have a 1:1 stoichiometry. +2 −2 ( ) + −4 ( ) ⇔ [ − ]( ) Equation 4.1 In the determination of hardness with EDTA several competing equilibriums are involved. The sample solution is buffered at a pH of 10.0 ± 0.3 as a compromise between chelate stability and the need to prevent precipitation of metal ions being analyzed. In order to prevent precipitation of metal ions (as CaCO3(s) or Mg(OH)2(s)), an ammonia buffer is specifically used in this analysis since ammonia can also act as ligand and form weak metal complexes with Ca+2 and Mg+2. a) b) Figure 4.1 Structure of the ethylenediaminetetracetic acid (EDTA) chelating agent showing a) the six coordination sites (4 carboxylic acids and 2 Nitrogen), and b) in a six-coordinate metal–EDTA complex with a divalent metal ion. Kf is the value of the formation or stability constant for a complex. Equation 4.2a and Equation 4.2b indicate that the [Ca-EDTA]-2 complex is 100 times more stable than the [Mg-EDTA]-2 complex: +2 −2 ( ) + −4 ( ) ⇔ [ − ]( ) −2 [ − ]( ) = [ +2 ][ −4 ] = 10+10.7 ( ) ( ) −2 +2 ( ) + −4 ( ) ⇔ [ − ]( ) −2 [ − ]( ) = [ +2 ][ −4 ] = 10+8.7 ( ) ( ) Equation 4.2a Equation 4.2b Aqueous solutions of EDTA as well as the complexes [Ca-EDTA]-2 and [Mg-EDTA]-2 are colorless. The identification of the end point of the complexation analysis depends on the use of a complexometric indicator. Eriochrome Black-T (EBT), used here as complexometric indicator, functions according to the following dynamics: At a pH≅10, the ionized form of Eriochrome Black T, [EBT(aq)]-4, displays a prominent ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 26 blue color while the complexes [Mg-EBT]-2 and [Ca-EBT]-2 are wine-red color. When a tiny amount of EBT is added to a sample of water containing hardness at a pH≅10, it combines with Ca2+ and Mg2+ ions to form complexes which shows a wine-red color. Equation 4.3a and Equation 4.3b indicate that [Mg-EBT]-2 is ~40 times more stable than [Ca-EBT]-2. +2 −4 -2 ( ) + ( ) ⇔ [ − ]( ) −2 [ − ]( ) = [ +2 ][ −4 ] = 10+5.4 ( ) Equation 4.3a ( ) −4 -2 +2 ( ) + ( ) ⇔ [ − ]( ) −2 [ − ]( ) = [ +2 ][ −4 ] = 10+3.8 ( ) Equation 4.3b ( ) Examining the Kf values shown in the two sets of Equations 4.2 (a and b) and Equations 4.3 (a and b), the corresponding [M-EDTA]-2 complexes are significantly more stable than the [M-EBT]-2 complexes, especially for Ca+2. The values of Kf for the EDTA complexes are larger than those for the EBT complexes by factors of 7.94x106 and 2,000 times larger for Ca2+ and Mg2+, respectively. The color of the sample ready for analysis is wine-red, once the titration starts the EDTA displaces EBT, replacing the [M-EBT]-2 complexes with [M-EDTA]-2 complexes favoring Ca+2 first and then Mg+2. The END of the titration analysis is indicated by the displacement of the last bit of wine-red [Mg-EBT]-2 by the colorless [Mg-EDTA]-2 and, therefore, the end point is indicated by the prominent blue color of the free [EBT(aq)]-4. The previous information is the chemistry behind the equivalence point for the complexometric analysis of Ca+2 and Mg+2 Hardness using EDTA as complexation agent and the detection of the end point using EBT as indicator. To determine the Hardness based on the volume of titrant consumed to reach the equivalence/end point, an analogous process to that shown for EXPERIMENT #3 is used. Using the same theory applied for the derivation of Equation 3.4 and Equation 3.5 (page 22), in the case of Hardness (“H”= Total Hardness as defined in page 25 ) we have that: × ℎ × 1 " " 1 = " " Equation 4.4 If CaCO3 is chosen as the surrogate substance representing “H”, then we can convert the moles of “H” to mass as CaCO3 by multiplying the moles by the Molecular Weight of CaCO3 (100 mg/mmol): " " × 100 mg CaCO3 3 ( ) ⟶ 3 Equation 4.5 For proper dimensional analysis note that the Molarity should be expressed as mmol/mL, the volume of titrant in mL and the volume of the sample in L. ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 27 Materials and Procedures Reagents: • • • Primary standard Ethylenediaminetetraacetic Acid (EDTA) solution, commercial. Water Hardness Buffer, with added Magnesium, for EDTA “hardness” titration (Commercial buffer contains: Magnesium Chloride Hexahydrate, Ammonium Chloride, Ammonium Hydroxide, Deionized Water) Freshly prepared Eriochrome Black-T (EBT) indicator solution or a very small amount of solid powder Samples for analysis: The following is a list of possible samples analyzed by each group. Duplicated analyses and additional experimental tests may be required at the time of the laboratory session. • • • • • Prepared solutions of known concentrations of Ca+2 and Mg+2 Tap water Samples of bottled drinking water Samples collected from local surface water lakes All groups should perform a blank test by titrating a specified volume of deionized water (defined in page 12) under the same conditions described in the Procedure. Figure 4.2 Setup for the analysis of water hardness using the EDTA titration method. The procedure for this analysis follows closely the APHA Method 2340C (APHA, 2005). This standard method highlights three important considerations:    Magnesium ion must be present to yield a satisfactory end point. To insure this, a small amount of complexometrically neutral magnesium salt of EDTA is added to the buffer (in our case commercial buffer); this automatically introduces sufficient magnesium and obviates the need for a blank correction. A limit of maximum 5 min is set for the duration of the titration to minimize the tendency toward CaCO3 precipitation. Ammonia buffer and EBT indicator must be added to each sample not any earlier than immediately before titration analysis. ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 28 Procedure: 1. Fill a burette with the standard EDTA solution (titrant). 2. Using a graduated cylinder or pipette, measure a specified volume of each of the sample for analysis and place them into separate 250 mL Erlenmeyer flask. Immediately before starting the titration of each sample: 3. Add 2-3 mL of buffer solution and verify that the pH is 10.0 ± 0.3. 4. Add a few grains of the EBT indicator solid powder indicator, enough to produce a distinct wine-red color but that leaves the solution transparent. 5. Titrate the sample to a blue color end-point. 6. Measure the pH at the end point. Literature Cited 1) American Public Health Association (APHA), 2005. Method 2340C, in “Standard method for examination of water and wastewater”, 21st Edition. APHA, AWWA, WPCF, Washington, DC. 2) Harvey, D., 2019. LibreText Book: Analytical Chemistry 2.0 (Harvey). 9.3: Complexation Titrations. Last updated June 29, 2019. ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 29 EXPERIMENT #5: Determination of different forms of alkalinity in water samples by titration with a standard acid solution Learning Objectives: • Determine the concentrations of the different forms of alkalinity in water samples by using the twoendpoint titration method with a standard acid solution. • Describe the reactions that are involved in the titration. • Discuss the association between the pH of the sample and the anions contributing to alkalinity. Key Concepts and Definitions Recall the concepts and definitions from pages 19 and 20: • Analyte • Standard Solution • Primary standards • Titration • Equivalence point of a titration • Endpoint of a titration Alkalinity: Is a chemical property of water. The Standard Method APHA 2320 defines “alkalinity of a water is its acid-neutralizing capacity. It is the sum of all the titratable bases. The measured value may vary significantly with the end-point pH used. Alkalinity is a measure of an aggregate property of water and can be interpreted in terms of specific substances only when the chemical composition of the sample is known (APHA, 2005).” By convention, the alkalinity of each of the three anionic species predominantly contributing to the alkalinity of water, hydroxide (OH-), bicarbonate (HCO3-) and carbonate (CO32-), are reported in concentration units of mg/L as calcium carbonate (CaCO3). Theoretical Background • Principle of the Acid Titration Method Given that alkalinity is the “acid-neutralizing capacity” of water, it is fitting that a titration analysis would be based on the principle of acid-base neutralization and involve a strong acid as a titrant. Thus, at its most fundamental level the alkalinity of a water sample can be determined by measuring the volume of standard acid that must be added to a given volume of the sample to lower its pH to a specified value. To be able to determine the alkalinity of a water sample, the titrant must have an accurately known concentration. Either sulfuric acid (H2SO4) or Hydrochloric acid (HCl) are typically used as titrants sources of protons, H+. Waters with alkalinity have a pH on the basic side of the pH-scale, and as the titrant acid is added the pH decreases. The Method APHA 2320 involves identifying and recording two endpoints known as Palkalinity and T-alkalinity defined as follows. In these formulas from the APHA 2320 method, the square bracket symbols ([ ]) can be understood as either equivalent volume of titrant or concentrations. Concentrations can be expressed in three different units: mmol/L, meq/L, and more often as mg/L as calcium carbonate (CaCO3). ENV 3001L: Environmental Engineering Laboratory: Spring 2024 FLORIDA INTERNATIONAL UNIVERSITY Page | 30 Phenolphthalein alkalinity (P) is always determined by titration to a pH=8.3. At this pH the phenolphthalein indicator changes from pink to colorless. The volume of acid used represents the total concentration of hydroxide [OH-] ions and one half of the carbonate [CO32-] ions present in the sample. P= [OH-] + ½[CO3-2] Equation 5.1 Total alkalinity (T) is determined by titration to a specified acidic pH of 5.1, 4.8, 4.5 or 4.0 depending upon the amount of carbon dioxide present (Standard Method 2320 A). In this laboratory Experiment Total Alkalinity is determined by titration to endpoint detected by using as indicators either methyl orange or mixed bromocresol green-methyl red. The observed changes of color and associated pHs are: for methyl orange from yellow to salmon pH~4, and for mixed bromocresol green-methyl red from blue-green to violet-gray (light pink) pH=4.5. T=[OH-] +[CO3-2]+[HCO3-] Equation 5.2 The results obtained from the titration, P-alkalinity and T-alkalinity volumes, provide a way to classify and quantify the major forms of alkalinity present in the sample. In Method 2320 the entire alkalinity is attributed to the anions hydroxide (OH-), carbonate (CO32-) and bicarbonate (HCO3-), assuming the absence of other (weak) inorganic or organic acids, such as silicic, phosphoric, and boric acids. It further assumes the incompatibility of hydroxide (OH-), and bicarbonate (HCO3-) alkalinities. Classifications are made on a per-volume basis using the specifications given in Table 5.1. Each of these volumes are then converted to concentration units. The concentration of the alkalinity components and of P-alkalinity and Talkalinity can be reported as milliequivalents per liter (meq/L) or mg/L as CaCO3, and, less frequently, as millimoles per liter (mmol/L). Table 5.1 Relations used to distribute the volume of titrant consumed in the two end-point titration among hydroxide (OH-), carbonate (CO3-2) and bicarbonate (HCO3-) ions Result of Titration P=0 2P < T Hydroxide (OH-) 0 0 Carbonate (CO3-2) 0 2P Bicarbonate (HCO3-) T T-2P 2P = T 2P > T
P=T
0
2P-T
P
2P
2(T-P)
0
0
0
0
If CaCO3 is chosen as the surrogate substance representing, in each case, OH-, HCO3-, and CO3-2 then
a direct relationship must be established between the moles of titrant consumed to reach the point
equivalent to each analyte and the CaCO3 surrogate analyte (MW of CaCO3 = 100 mg/mmol). If sulfuric acid
(H2SO4) is used as titrant, the chemical reaction between the titrant and the surrogate analyte is shown in
Equation 5.3a
2 4 ( ) + 3 ( ) → 4 ( ) + 2 3 ( )
ENV 3001L: Environmental Engineering Laboratory: Spring 2024
FLORIDA INTERNATIONAL UNIVERSITY
Equation 5.3a
Page | 31
and using an analogous process to that shown for EXPERIMENT #3 and EXPERIEMNT #4, one arrives at the
following expression for the concentration of each analyte:
1 3 100 mgCaCO3
×
1 2 4 3
2 4 × 2 4 ×

⟶

3
Equation 5.3b
If hydrochloric acid (HCl) is used as titrant, the chemical reaction between the titrant and the surrogate
analyte is shown in Equation 5.4a
2 ( ) + 3 ( ) → 2 ( ) + 2 3 ( )
Equation 5.4a
and using an analogous process to that shown for Experiment #3, one arrives at the following expression
for the concentration of each analyte:
× ×
1 3 100 mgCaCO3
×
3
2 Cl

⟶

3
Equation 5.4b
In Equation 5.3b and Equation 5.4b, a proper dimensional analysis requires that the volumes of titrant,
distributed by applying Table 5.1, should be in milliliters (mL), the molarity (M) of the titrant should be
expressed in mmol/mL (numerically identical to mol/L), and the volume of the sample in liter (L).
Watch the Instructor’s presentation posted on Canvas to learn how to express the concentrations of the
individual analytes in meq/L and mmol/L units.
Materials and Procedures
Reagents:
•
•
•
•
Standard Acid titrant solution.
Phenolphthalein indicator 0.5% alcoholic solution
Methyl orange indicator 0.5% aqueous solution (*)
Bromocresol Green-Methyl Red commercial mixture for Alkali