EDTA titration lab - Lecture notes 7 PDF

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Determination of Ca 2+ and Mg2+ in Water by EDTA Titration What are Lewis Acid-Base Reactions? In your study of acid-base chemistry, you learned of different models for describing the behavior of acids and bases. In the Arrhenius model, acids are substances that increase the concentration of hydronium ion (H3O+) in aqueous solutions, while bases increase the concentration of hydroxide ion (OH) in aqueous solutions. In the Bronsted-Lowry model, acids are proton donors and bases are proton acceptors. Finally, the Lewis model considers acids to be electron pair acceptors and bases to be electron pair donors. In this lab, you will use a Lewis acid-base reaction to determine the concentration of Ca2+ and Mg2+ in a water sample. These ions and can inhibit the action of soaps and cause precipitates called limescale. Water with high concentrations of these ions is said to be “hard”. Because, these ions have a positive charge and vacant orbitals, they are good electron pair acceptors and therefore Lewis acids. They can accept electron pairs from a donor that has an unshared pair of electrons (i.e., a Lewis base). When a Lewis base donates its electrons to a metal ion to form a complex ion, it is called a ligand. An example of a Lewis acid-base reaction involving the cyanide ion as a ligand is shown below.

The reaction of the silver ion with the cyanide ion produces a complex ion in a process called complexation. Complex ions are similar to polyatomic ions in that they are groups of covalently bonded atoms that carry an overall charge. Complexation reactions are reversible and are characterized by an equilibrium constant called a formation constant, Kf. Values of Kf are generally large, which indicates that these equilibria strongly favor products, leaving a relatively low concentration of free metal ions in solution. In fact, complexation reactions are commonly used to remove unwanted metal ions from

solutions. For example, EDTA is a complexing agent commonly added to food packaging in order to complex metal ions that catalyze reactions that cause food to spoil.

What is EDTA? EDTA is an example of a multidentate (many-toothed) ligand, which can bind metal ions through multiple atoms. Multidentate ligands are also called chelates, which comes from a Greek word meaning claw. Multidentate ligands, or chelating agents, wrap themselves around metal ions like a claw. The figure below shows EDTA chelating a metal ion.

EDTA Chelating a metal ion

A more complete structure and a shorthand line structure of EDTA are shown below.

A line structure is a shorthand way of depicting an organic molecule. The lines in these structures represent covalent bonds. Because all organic compounds contain carbon and hydrogen, these atoms are not shown. A carbon atom is assumed to be present at the intersection of two lines. Because carbon typically forms four bonds, the number of hydrogen atoms bonded to each carbon atom can be inferred. What are the different forms of EDTA? EDTA is an amphiprotic substance, which means that it can both donate and accept protons. The four hydrogen atoms shown in the above line structure are acidic, and because of this, the formula of EDTA is often abbreviated H4Y, where H4 represents the four acidic hydrogen atoms and Y represents the remaining structure. The Ka values for the sequential loss of these protons are 1, 0.032, 0.01, and 0.0022. The lone pairs of electrons on the nitrogen atoms in EDTA are capable of accepting protons. The Kb values for the protonation of these nitrogen atoms are 1.74 × 10-4 and 1.45 × 10-8. EDTA can therefore exist in many different forms depending on the pH of the solution. At very low pH, EDTA will be present in its completely protonated form H6Y2+. At very high pH, EDTA will be present in its completely deprotonated form, Y4-. At intermediate pH, EDTA will be present in one of its intermediate forms. The following structures depict the various ionized forms of EDTA, along with their abbreviations and the pH at which each predominates.

The negative log of an equilibrium constant is called a pK value. The pKa value, the negative log of the Ka value, for an acid gives the pH at which there is an equal concentration of an acid and its conjugate base. This can be seen by solving the Henderson-Hasselbalch equation for the pH at which an equal concentration of an acid and its conjugate base would be present in a solution. The Ka and Kb values can be used to determine the predominant form of an amphiprotic substance at different values of pH. The following table was constructed using the pKa and pKb values for EDTA. Form of EDTA

pH

H6Y2+

H5H+

0.0

H4HY

1.5

H3Y-

2.0

H2Y2-

2.66

HY3-

6.16

Y4-

10.24

Which form of EDTA do we start with? EDTA is most frequently purchased as the dihdrate salt of its H2Y2- form, Na2H2Y∙2H2O (structure shown below). Recall that hyrdrates are ionic compounds with loosely bound water molecules in their crystal structure.

Which form of EDTA forms the most stable complex ions? Very stable complex ions are formed between metal ions and EDTA in its completely deprotonated form, Y4-. In order to get a significant portion of EDTA into this form, solutions of EDTA used in titrations are typically buffered at high pH. In this experiment, you will use a buffer of NH3 and NH4Cl to maintain a pH of 10.

The completely deprotonated form of EDTA, Y4-, can form bonds through the lone pairs of electrons on the 4 oxygen atom and 2 nitrogen atoms for a total of 6 binding sites. Because of the free rotation about single (sigma) bonds EDTA is able to wrap around metal ions and orient these six binding sites toward the metal ion. These six binding sites are arranged in the shape of an octahedron around the metal ion as is shown in the above figure for the EDTA chelation of a metal ion. How is EDTA used to determine the concentration of metal ions in a solution? EDTA is able to form stable 1:1 complex ions with many different metal ions, and can therefore be used in titrations to determine the concentrations of these metal ions in a solution. Because the reactions used in these titrations involve the formation of complex ions, they are called complexiometric titrations. How do indicators in complexiometric titrations work? Indicators used compexiometric titrations are complexing agents themselves. Unlike EDTA, however, they are colored, and their color is different in their free and complexed forms. For example, the color of Eriochrome black T, one of the indicators used in this experiment, is blue in its free form and wine red when it is complexed with a metal ion. The color of these indicators is also affected by pH, which is another reason that buffers are often used to regulate pH during complexiometric titrations. When the indicator is added to water containing metal ions it will rapidly form a complex ion, and the solution will take on the color of the complexed form of the indicator. As EDTA is added to the solution it will react with the free metal ions in the solution first. After enough EDTA has been added to react with all of the free metal ions, additional EDTA will displace the metal ions in the indicator-metal ion complex. This occurs because the formation constant for the EDTA-metal ion complex is larger than that of the indicator-metal ion complex. In other words, EDTA binds more strongly to the metal ions. As the indicator loses its metal ions, the solution will become the color of the indicator in its free form. The following summary and reactions illustrate this process. Initially, indicator is added to a solution containing metal ions. The indicator reacts with the metal ions in solution to form complex ions, and the solution takes on color of indicator-metal complex ion.

Prior to equivalence point, added EDTA reacts with free metal ions in solution to form complex ions. The solution color does not change during this part of the titration.

As the equivalence point is approached, added EDTA displaces metal ions from the indicator-metal complex ion, and the solution takes on color of the free indicator.

How can the individual concentrations of Ca2+ and Mg2+ be determined? In this lab, you will use EDTA to determine the concentration of metal ions in a water sample. The most common positively charged metal ions in natural waters are Ca2+ and Mg2+. For this lab, you will assume that these are the only metal ions present in your water. You will perform two sets of titrations. One set of titrations will be used to determine the total concentration of Ca2+ and Mg2+ present in your water sample. After selectively precipitating the Mg2+ as Mg(OH)2, you will perform a second set of titrations to determine the concentration of Ca2+ in your water. The concentration of Mg2+ in your water sample can then be determined by difference. Materials EDTA: Na2H2EDTA∙2H2O (molar mass 372.24 g/mol), 0.6 g per student pH 10 Buffer: Add 142 mL of 28 % (by mass) aqueous NH3 to 17.5 g of NH4Cl and dilute to 250 mL with deionized water. Eriochrome black T indicator: Dissolve 0.2 g of the solid indicator in 15 mL of triethanolamine plus 5 mL of absolute ethanol 50 % (by mass) aqueous NaOH Water Standard: Evian® bottled water or other brand if calcium and magnesium concentrations are known. Unknowns: Collect a water sample from a stream, lake, or ocean. Use a plastic bottle and fill it completely to the top so that no air will be present in the sealed bottle. This will minimize the growth of bacteria in your water. Procedure Solution preparation: If the EDTA has not been dried for you, start by drying about 1 g of Na2H2EDTA∙2H2O (molar mass 372.24 g/mol) at 80°C for 1 h and cool in a desiccator. Weigh out ~0.6 g of the dry EDTA and record the mass to the nearest 0.0001 g. Add the EDTA to a 500 mL volumetric flask and add about 400 mL of deionized water. Dissolve the EDTA with swirling and heating if necessary. If you use heating, you will need to allow your EDTA to cool to room temperature before diluting to the 500-mL mark on the neck of the volumetric flask. Use a squirt bottle of di H2O to careful bring the volume of the solution to 500 mL.

Water standard: In order to test your reagents and the accuracy of the titration, you may first wish to titrate a water sample with a known concentration of calcium and magnesium. Bottled mineral water for which these concentrations are known can serve as a standard. Unknown water samples: If you choose to analyze unknown water samples, you will need to determine the appropriate sample size. The ability of titrations to precisely to determine unknown concentrations relies on the ability to precisely deliver volumes with a buret. In order to deliver a volume to 4 significant digits using a buret, you must deliver at least 10 mL. Precision increases with the delivery of higher volumes. Refilling the buret however will increase the number of volume readings that must be made and therefore increase the uncertainty in determining the total volume delivered. The goal is to deliver as much titrant as possible without having to refill the buret. Using 20-40 mL of titrant is a reasonable goal. Because the concentrations of Mg2+ and Ca2+ in your unknown water sample are unknown, it is impossible to know what sample volume will require 20-40 mL of titrant. Therefore, you will need to perform some quick titrations (using the procedure below) with different sample volumes to determine the appropriate sample size. Start with a 50-mL sample size and then adjust accordingly. If the appropriate sample volume is less than 50 mL, bring the total volume up to ~50 mL by adding sufficient deionized water. Once you have selected an appropriate sample size, perform careful titrations on at least three samples of your unknown water. Total Mg2+ and Ca2+ determination: Pipet a 50-mL sample of water into a 250-mL Erlenmyer flask. To the sample add 3 mL of the pH 10 buffer and 6 drops of Eriochrome black T indicator. Rinse and fill a clean 50-mL buret with your EDTA solution and record the initial volume. Titrate your water sample until the color changes from wine red to blue. Perform a quick titration to determine the approximate volume needed to reach the endpoint. In order to accurately identify the endpoint, where the last trace of red has been removed from the solution, you may want to add small volumes of mineral water to your titrated sample and practice reaching the endpoint. Once you have confidently identified the endpoint, save this sample as a reference for subsequent titrations. Carefully titrate 3 more samples of the water. Blank titration: Titrate a 50-mL sample of the lab water that was used to prepare your solutions to determine whether it contains any calcium or magnesium ions. If the lab water requires the addition of EDTA to reach the endpoint, you will need to subtract this volume from each of your titration volumes. Ca2+ determination: If you wish to determine the individual concentration of calcium and magnesium in your water sample, use the following procedures for precipitating the Mg2+ and titrating the Ca2+ that remains in solution. Pipet four water samples into separate Erlenmyer flasks. To precipitate the Mg2+, add 30 drops of 50% NaOH to each sample and swirl for 2 min. The precipitate may not be visible. Because Eriochrome black T does not work well at this elevated pH, you will use hydroxynaphthol blue as the indicator in these titrations. Add ~0.1 g of hydroxynaphthol blue to each sample. Perform one quick titration to determine the approximate endpoint, and practice finding the endpoint if necessary. To accurately determine the endpoint of each of the remaining samples, titrate to the blue endpoint and then allow the sample to sit for 5 min with occasional swirling. This will allow any Ca(OH)2 that may have precipitated to redissolve. If the solution has turned back to red, add additional titrant to reach

the blue endpoint. Following the same procedure, perform a blank titration on a 50-mL sample of lab water and make any volume corrections necessary. Calculations: From the precise mass of EDTA that you weighed out, calculate the concentration of your EDTA. Use the volume of EDTA needed to reach the endpoint in each of your titrations to determine either the total concentration of Ca2+ and Mg2+ or the concentration of Ca2+ alone. Recall that EDTA reacts with either of these ions in a 1:1 mole ratio. If you performed the Ca2+ titration, determine the Mg2+ concentration by difference from the total of Ca2+ and Mg2+. Calculate the relative standard deviation of replicate titrations. If bottled water was titrated, calculated the relative difference of your Ca2+ and Mg2+ concentrations to those listed by the manufacturer.

Pre-lab Questions

Determination of Ca2+ and Mg2+ in Water by EDTA Titration

Name:_______________________ Instructor:____________________

1. If a student prepared an EDTA solution for a complexiometric titration by dissolving 0.5946 g of Na2H2EDTA∙2H2O (molar mass 372.24 g/mol) in enough water to bring the total solution volume to 500.0 mL, what would the molar concentration of EDTA in the solution be? Clearly show any required calculations with proper units and significant digits.

2. A student titrates a 50.00-mL sample of water with a 0.003125 M EDTA solution. If the titration requires 34.64 mL of EDTA to reach the endpoint, what is the total concentration of Mg2+ and Ca2+ in the water sample? Clearly show any required calculations with proper units and significant digits.

3. The same student then adds NaOH to another 50.00 mL sample of the same water and then titrates again with EDTA. This time the titration requires 23.67 mL of EDTA solution. Determine the individual concentrations of Ca2+ and Mg2+ in the water sample. Clearly show any required calculations with proper units and significant digits.

Determination of Ca2+ and Mg2+ in Water by EDTA Titration

Report Sheet

Name:_______________________ Instructor:____________________

Mass of Na2H2EDTA∙2H2O used to prepare EDTA solution (g) _______________ Concentration of EDTA solution (M) ______________________ Water standard Total Mg2+ and Ca2+ determination:

Initial Volume (mL)

Final Volume (mL)

Volume of EDTA titrant Volume Volume for Delivered (mL) blank (mL)

Corrected Volume (mL)

Rough Trial Trial 1 Trial 2 Trial 3 Blank Total concentration of Ca2+ and Mg2+ (M) Trial 1 Trial 2 Trial 3 Mean Standard Deviation Relative Standard Deviation (%)

Ca2+ determination:

Initial Volume (mL) Rough Trial Trial 1 Trial 2 Trial 3 Blank

Volume of EDTA titrant Final Volume Volume for Volume (mL) Delivered (mL) blank (mL)

Corrected Volume (mL)

Concentration of Ca2+ (M) Trial 1 Trial 2 Trial 3 Mean Standard Deviation Relative Standard Deviation (%)

Total concentration of Ca2+ and Mg2+ (M)

Mean Concentration of Ca2+ (M)

Concentration of Mg2+ by difference(M)

Comparison of manufacturer-reported and experimentally-determined Mg2+ and Ca2+ concentrations Ca2+ concentration (M) Reported by Manufacturer Experimental Relative Difference (%)

Sample calculations (use proper units and significant digits): EDTA concentration:

Total concentration of Ca2+ and Mg2+:

Concentration of Ca2+:

Concentration of Mg2+:

Mg2+ concentration (M)

Determination of Ca2+ and Mg2+ in Water by EDTA Titration

Report Sheet

Name:_______________________ Instructor:____________________

Unknown water samples Volume of water sample used for titration (mL)__________________________ Total Mg2+ and Ca2+ determination:

Initial Volume (mL)

Volume of EDTA titrant Final Volume Volume for Volume (mL) Delivered (mL) blank (mL)

Corrected Volume (mL)

Rough Trial Trial 1 Trial 2 Trial 3 Blank

Total concentration of Ca2+ and Mg2+ (M) Trial 1 Trial 2 Trial 3 Mean Standard Deviation Relative Standard Deviation (%)

Ca2+ determination:

Initial Volume (mL) Rough Trial Trial 1 Trial 2 Trial 3 Blank

Volume of EDTA titrant Final Volume Volume for Volume (mL) Delivered (mL) blank (mL)

Corrected Volume (mL)

Concentration of Ca2+ (M) Trial 1 Trial 2 Trial 3 Mean Standard Deviation Relative Standard Deviation (%)

Total concentration of Ca2+ and Mg2+ (M)

Mean Concentration of Ca2+ (M)

Sample calculations (use proper units and significant digits): EDTA concentration:

Total concentration of Ca2+ and Mg2+:

Concentration of Ca2+:

Concentration of Mg2+:

Concentration of Mg2+ by difference(M)

Post-lab questions

Determination of Ca2+ and Mg2+ in Water by EDTA Titration

Name:_______________________ Instructor:____________________

1. Write out the net ionic equation for the reaction that was used to remove Mg2+ from your water sample so that Ca2+ alone could be determined.

Why is it that Mg2+ is removed from solution but Ca2+ is not? The Ksp values for the hydroxides of Mg2+ and Ca2+ are 1.8 × 10-11 and 6.5 × 10-6 respectively.

2. Calcium Disodium Versenate (see structure below) is a form of EDTA that can be used to treat lead poisoning in humans. In this drug, EDTA is present as a complex ion with Ca2+. How is it possible for this drug to remove Pb2+ from a person’s blood if it is already complexed to Ca2+? The formation constants, Kf values, for the EDTA- Ca2+ and EDTA- Pb2+ complex ions are 4.9 × 1010 and 1.1 × 1018 respectively.

3. Two indicators, Funky Green (FG) and Infernal Orange (IO), are being considered for an EDTA titration for the d...