Lab Manual Quantitative Analytical Method, Deepak Pant [epub ebook reader .TXT] 📗

are:
Ag+ + Cl- AgCl (s) white
SCN- + Ag+ AgSCN (s) white
Fe3+ + SCN- Fe(SCN)2+ red
The most important application of the Volhard method is for determing the presence of halide ions. An excess of silver nitrate is added to the sample and back-titrated with a standard thiocyanate solution.
EXPERIMENT: THE VOLUMETRIC DETERMINATION OF CHLORIDE
Precipitate-forming titrations are not common. However, they are widely used for determination of the halides, chloride (Cl-), bromide (Br-) and iodide (I-), using silver nitrate (AgNO3) as the titrant. In this part of the experiment, chloride will be determined by titration with silver nitrate using dichlorofluorescein as an indicator. This is commonly known as Fajan's Method.
THEORY
Chloride present in a sample is quantitatively insoluble in a solution containing excess of silver ion:
Ag+(aq) + Cl- (aq) AgCl(s) Ksp = 1.82 x 10 10
The same determination may be accomplished volumetrically if a standard solution of Ag+ is available. The end point of the reaction may be determined through the use of an adsorption indicator, dichlorofluorescein. Its function may be described as follows:
If the reaction is run in neutral or basic solution some of the indicator will dissolve to form the dichlorofluorescinate anion, which is represented as In . Before the equivalence point, with Ag+ as titrant, excess Cl is present in solution. The excess Cl- is adsorbed onto the precipitate particles formed and the indicator anion is repelled by the negatively-charged precipitate.
Ag+(aq) + 2Cl-(aq) AgCl:Cl (s)
At the equivalence point, there is little or no excess Cl , and just beyond the equivalence point Ag+ is in excess and becomes the primary adsorbed ion. The charge on the precipitate changes from negative to positive and the indicator anion is adsorbed.
AgCl:Cl (s) + Ag+(aq) AgCl:Ag+(s) + Cl (aq)
AgCl:Ag+(s) + In (aq) AgCl:Ag+In-(s)
yellow rose pink
The color change is:
yellow  rose pink

It is believed that the indicator anion (yellow) forms a complex ion with Ag+, adsorbed on the silver chloride precipitate, which alters its light absorbing properties, and hence its color. The indicator function is critically dependent on the availability of a large precipitate surface area to allow adsorption. The greatest surface area results from a precipitate comprised of very small particles (colloidal). Stabilization of these colloidal particles (recall that a colloid has a very high surface-to-volume ratio) is accomplished by adding a protective colloid, such as dextrin.
It is important that the titration be conducted quickly in diffuse light as photodecomposition of the silver chloride renders the solution purple, making it difficult to discern the pale pink end point signal from the purple background.
PROCEDURE
Preparation of a Standard Silver Nitrate Solution
Obtain from the instructor approximately 8.5 g of AgNO3 in a clean, dry, weighing bottle. Grind the contents to a fine powder with an agate mortar and pestle. Return the AgNO3 to the weighing bottle, dry in a oven at medium power, and accurately weigh, by difference, the contents into a 500-mL volumetric flask. Dissolve with deionized water and dilute to the mark. Store in a dark area and calculate the molarity of the solution.
Preparation of the Sample for Analysis
Using the unknown chloride sample from part 1, accurately weigh three samples of the unknown into three 250-mL Erlenmeyer flasks (the sample size should be approximately 0.3 g). Add approximately 50 mL of deionized water to each flask and swirl to dissolve. Add 10 mL of a 2% dextrin suspension, 5 drops of dichlorofluorescein indicator, and titrate each sample to the rose pink end point.
CALCULATION OF RESULTS
Calculate your results as the % Cl in your sample. Be certain to report each individual value, the mean value, the absolute deviation from the mean of each value, the relative average deviation in parts per thousand, the standard deviation and the confidence interval at an appropriate confidence level.
4. Complex metric titration
The technique involves titrating metal ions with a complexing agent or chelating agent (Ligand) and is commonly referred to as complexometric titration. This method represents the analytical application of a complexation reaction. In this method, a simple ion is transformed into a complex ion and the equivalence point is determined by using metal indicators or electrometrically. Various other names such as chilometric titrations, chilometry, chilatometric titrations and EDTA titrations have been used to describe this method. All these terms refer to same analytical method and they have resulted from the use of EDTA (Ethylene diamine tetra acetic acid) and other chilons. These chilons react with metal ions to form a special type of complex known as chelate.
Objective
1. To perform water analysis regarding the presence of chloride, sulphate, and ammonium ions.
2. To measure water hardness caused by dissolved calcium and magnesium ions by complexometric titration

Equipment
 Beaker
 Test tube
 pH paper
 Erlenmeyer flask
 Burette
 Stands
 Pipette

Chemicals and Materials
 Aquadest
 Tap water
 AgNO3 (silver nitrate)
 BaCl2 (barium chloride)
 NaOH (sodium hydroxide)
 1 M Ammonium buffer (1M NH4OH)+ 1 M NH4Cl)
 5x10-3 M EDTA solution (ethylenediaminetetraacetic acid)
 Eriochrome Black T
Procedure
2 beakers were prepared and filled with the following fluid:
- Beaker 1: 20 ml distilled water
- Beaker 2: 20 ml tap water
2. The appearance and smell of each sample were checked and the pH values were measured using pH paper.
3. A few drops of 1 n AgNO3 were put into each beaker.
4. Observation was done to see whether precipitation occurred. When precipitation occurred, chloride ions were present.
5. Some crystals of BaCl2 were dissolved in 5 ml of distilled water.
6. Another set of beakers were prepared and filled with appropriate fluids
7. A few drops of BaCl2 solution were put into each beaker.
8. Observation was done to see whether precipitation occurred. When precipitation occurred, sulphate ions were present.
9. 2 test tubes were prepared and filled with the following fluid:
- Test tube 1: distilled water
- Test tube 2: tap water
10. One flake of NaOH was put inside each test tube.
11. The appearance and smell (pungent smell of ammonia) were observed for presence of ammonium chloride.

Measurement of water hardness with EDTA titration:
1. The 1500 ml of 5 x 10-4 EDTA solution was prepared by mixed the 50 ml of 5 x 10-3 M EDTA with solution of 50 ml NH4OH 1 M and 50 ml NH4Cl.
2. The 25 ml tap water was put into an Erlenmeyer flask and it was added with 20 drops of ammonia buffer and 10 drops of Eriochrome Black T indicator. This solution color was red purplish.
3. The burette was clamped to the ring stand and the EDTA solution was put into the burette until it reached 50 ml.
4. Then, the solution in the Erlenmeyer flask was titrated with EDTA solution until its color reached blue color.
5. The total volume of EDTA solution that being used to titrate the solution in the Erlenmeyer flask was recorded.
6. The second procedure until fifth procedure was repeated once.

I. Observation and Measurement Result

Chemical check of water

CHARACTERISTIC DISTILLED WATER TAP WATER
Appearance Transparent Transparent
Smell No smell No smell
pH 5.5 6.5
Table 1: Characteristics of distilled and tap water

CONDITION DISTILLED WATER TAP WATER
Addition of AgNO3 Colorless, no precipitation Turbid
Addition of BaCl2 Colorless, no precipitation Colorless, no precipitation
Addition of NaOH Warm, air bubbles form a bit, no smell Warm, air bubbles form a bit, no smell
Table 2: The reactions of each type of water after addition of compounds.

Measurement of water hardness with EDTA titration:
Distilled water
After it was added with the Eriochrome Black T Indicator, the solution color turned to dark blue.

Tap water
After it was added with the Eriochrome Black T Indicator, the solution color turned to purplish red. After it was titrated with EDTA, the solution color turned to blue.
The volume of EDTA used in the first titration = 16.6 ml.
The volume of EDTA used in the second titration = 15.7 ml.


Figure: end point of titration
II. Evaluation and Calculation

Chemical check of water
In this experiment, tap water and distilled water was checked to observe the chemical compound contained in each of the water. Unfortunately, sample water from the river was not given for checking. Therefore, not much comparison could be made.
Distilled water should only contain pure water (hydrogen and oxygen molecules) with no addition of gasses, metal ions, and other compounds. This was because distilled water was made through the process of distillation which purifies the water from contamination. The appearance of distilled water is transparent and no smell was detected. This was because there were no contaminants that give out odor.
The tap water observed in our experiment was taken from the laboratory. The appearance of tap water was transparent and may contain minerals, disinfectant, and contaminants. Chloride and sulphate ions may also be present in tap water. However, the substances contained in tap water were odorless. Therefore, no smell was detected. The tap water’s pH is 6.5.
When each of the water samples were added with AgNO3, the AgNO3 ions dissociates:
AgNO3(aq) → Ag+(aq) + NO3-(aq)
Then, the Ag+ ions reacted in the presence of Cl- ions:
Ag+(aq) + Cl-(aq) → AgCl(s)

AgCl is insoluble in water and it precipitated to become solid. The color of AgCl solid is white. When AgNO3 was put into distilled water, no precipitate was found. However, when it was put into tap water, the water became turbid. This shows that the tap water contains chloride ions.
When each of the water samples were added with BaCl2, the ions dissociates:
BaCl2(aq) → Ba2+(aq) + 2Cl-(aq)
Then, the Ba2+ ions reacted in the presence of SO42- ions:
Ba2+ (aq) + SO42- (aq) → BaSO4(s)
BaSO4 is insoluble in water and it precipitated to become solid. The color of BaSO4 solid is white. When BaCl2 was put into distilled water and tap water, no precipitate was found. This shows that the tap water and distilled water does not contain sulphate ions.
When each of the water samples were added with NaOH, the ions dissociates:
NaOH(s) → Na+(aq) + OH-(aq)
Then, the OH- ions reacted in the presence of NH4+ ions:
NH4+ (aq) + OH- (aq) → NH4OH(aq)
NH4OH(aq) → NH3(g) + H2O(l)
NH3 is a gaseous compound and can be detected by its pungent smell. After the addition of NaOH to both samples, the test tube becomes warm. This shows that the reaction is exothermic. Also, water bubbles were observed at both test tubes. However, the smell of ammonia was not detected. This shows that ammonia gas may formed but the amount is too small to be detected. The formation of ammonia gas shows that ammonium chloride is present.
Measurement of water hardness with EDTA titration:
Water hardness is an expression for the sum of the