INORGANIC AND ORGANIC CHEMISTRY LABORATORY REPORT

Subject : Inorganic and Organic Chemistry LaboratoryLecturer : Dr. rer. nat. Filiana Santoso, Mr. Hery Susanto M.SiInstructor : Mr. Tabligh Permana, Mr.Hery Sutanto M.Si, Ms. Sylvia Yusri,

S.SiFaculty/Class : Life Science/LS 2ADate of Experiment : 25 March 2014Date of Lab. Report : 8 April 2014Semester : 2Time of Experiment : 14.00 – 17.00 p.m

Experiment: Water Analysis

Name: Kristania HadhiwaluyoChita Sakina PutriantiElias Harmanto

S W I S S G E R M A N U N I V E R S I T Y

I. Objectives To determine which of these ions, Ag+ and Ba2+ and also NH3,that is

present in the tap water sample. To determine the hardness of water by measuring the concentration of

Ca2+ ion.

II. Theoretical Background Water has been known as one of the most important thing that is needed for human to survive. Human needs water for almost every activity they do even their body mostly contain of water, hence this very important role of water shows the importance of maintaining the quality of water especially when it is used for drinking. In the case of this experiment, tap water is analyzed and checked whether these ions, Ag+ and Ba2+ and also NH3,are present in the tap water sample. Checking the presence of these substances is important as a part of water analysis because tap water is treated with a large number of chemicals in order to kill bacteria and other microorganisms. In addition, it may contain other undesirable contaminants like toxic metal salts, hormones and pesticides, or it may become contaminated by chemicals or microbes within pipes (e.g. lead, bacteria, protozoa) where it is being distributed. Tap water usually contains chlorine, fluorine compounds, trihalomethanes (THMs), hormones, nitrates, pesticides, and salts of; arsenic, radium, aluminium, copper, lead, mercury, cadmium, barium.

There are three main substances that will be checked from this experiment, firstly the presence of Chloride. Chlorine is added to drinking water, often by addition of hypochlorous acid in order to control disease causing organisms and provide protection until water is consumed. Several water properties and constituents affect the efficiency of chlorination, and therefore affect the amount of chlorine that must be added to achieve levels of chlorine compounds that have disinfecting capability (hypochlorous acid, hypochlorite and inorganic chloramine). Secondly, asides from chlorine, ammonia, ammonium, organics, and reducing agents usually added to the water. However, although both of the substances have disinfectant capability, they are slow reacting and able to add

Campus BSD CityBumi Serpong DamaiTangerang 15321 – Indonesia

Tel.+62 21 537 6221Fax. +62 21 537 6201

sgu.info@sgu.ac.id

www.sgu.ac.id

unpleasant taste and odor at high levels which is undesirable in water used for consumption. Another substance that will be checked in the tap water sample is sulfate ions. Sulfate is a substance that occurs naturally in drinking water. The presence of this ion may not be bad for human health however health concerns regarding sulfate in drinking water have been raised because of reports that diarrhea may be associated with the ingestion of water containing high levels of sulfate, more over the high level of sulfate in drinking water can also be bad for pipes as it may lead to corrosion. The presence of these substances will be checked through selective precipitation by carefully selecting reagents and conditions for reactions (for example shown in the diagram below), to separate groups of ions by its reactivity and by taking advantages of their different solubilities.

According to the figure shown above, there are 5 different groups of cations that are able to be separated by different kinds of reagents forming precipitates, in the form of solid that is separated from the solution. This proven

the theory of the solubility rules of ions in a solution, by following the solubility rules of ionic compound in the solution as follows various reagent can be determined to separate certain ions from the solution, as the targeted ion (ion that will be separated from the solution) will form precipitate when the reagent was added.

Soluble Compounds1. All compounds of the alkali metals (Group 1A) are soluble2. All salts containing NH4+, NO3-, ClO4-, ClO3-, and C2H5O2- are soluble.3. All chlorides, bromides, and iodides (salts containing, Cl-, Br-, or I-)

are soluble except when combined with Ag+, Pb2+, and Hg22+.4. All sulfates (salt containing SO42-) are soluble except those which

containing Pb2+, Ca2+, Ba2+, Sr2+, Ag+, and Hg22+. Insoluble Compounds

1. “Sulfides (S2-) are usually insoluble. Exceptions include Na2S, K2S, (NH4)2S, MgS, CaS, SrS, and BaS.

2. Oxides (O2-) are usually insoluble. Exceptions include Na2O, K2O, SrO, and BaO, which are soluble, and CaO, which is slightly soluble.

3. Hydroxides (OH-) are usually insoluble. Exceptions include NaOH, KOH, Sr(OH)2, and Ba(OH)2, which are soluble, and Ca(OH)2, which is slightly soluble.

4. Chromates (CrO42-) are usually insoluble. Exceptions include Na2CrO4, K2CrO4, (NH4)2CrO4, and MgCrO4.

5. Phosphates (PO43-) and carbonates (CO32-) are usually insoluble. 6. Exceptions include salts of the Na+, K+, and NH4+ ions” (Bordner

Group, n.d.).

The second thing that will be checked from this investigation is water hardness. Primarily, in terms of its mineral content water can be categorized into two categories which are hard and soft water. Just like its name, hard water is the water which contain high level of mineral content, whereas soft water is water that has low mineral content. Water is an excellent solvent and readily dissolves minerals it comes in contact with. As water moves through soil and rock, it dissolves very small amounts of minerals and holds them in solution. Calcium and magnesium dissolved in water are the two most common minerals that make water "hard." This is formed when water percolates through deposits of calcium and magnesium-containing minerals such as limestone, chalk and

dolomite. The hardness of water is referred to by three types of measurements: grains per gallon, milligrams per liter (mg/L), or parts per million (ppm).

The method that is being used to calculate the concentration of calcium ion (Ca2+) in the sample is by complexometric titration with EDTA (Ethylene Diamino Tetra Acetate). Basically, complexometric titration is a form of volumetric analysis in which the formation of a colored complex is used to indicate the end point of a titration. Complexometric titrations are particularly useful for the determination of a mixture of different metal ions in solution in this case the tap water sample. It is because some organic molecule can form stable complex with the metal ion inside the sample. EDTA is an example of a multidentate (many-toothed) ligand that is used in this experiment. EDTA (Ethylene Diamine Tetra acetate), has four carboxyl groups and two amine groups that can act as electron pair donors, or Lewis bases. The ability of EDTA to potentially donate its six lone pairs of electrons for the formation of coordinate covalent bonds to metal cations makes EDTA a hexadentate ligand. However, in practice EDTA is usually only partially ionized, and thus forms fewer than six coordinate covalent bonds with metal cations. In the case of this experiment, EDTA is used to bind calcium ion (Ca2+) in the tap water sample through multiple atoms according to the reaction shown below

III. Equipment and Materials Equipment:

- Erlenmeyer flask, 3, 100 cm3

- Test tube, 5- Rack of test tube- Volumetric pipette, 10 cm3

- Volumetric flask, 50 cm3

- Dropping Pipette, 2- Graduated pipette, 50 cm3

- Graduated Cylinder, 50 cm3

Ca2+ + EDTA4- [CaEDTA]2-

- Glass Beaker, 2, 100 cm3

- Funnel- Spatula, 3- Burette clamp - Burette, 50 cm3

Materials:- H2O(l), distilled water, 20 cm3

- H2O(l), tap water, 100 cm3

- EDTA, Ethylene Diamino Tetra acetate, 5 x 10-3 M, 50 cm3

- AgNO3(s), Silver Nitrate- BaCl2(s), Barium Chloride- NaOH(s), Sodium Hydroxide- Eriochrome Black T (EBT) Indicator- Ammonia buffer

IV. Procedures 1. Step 1: Qualitative Water Analysis

1. 100 cm3 tap water was poured into the 100 cm3 glass beaker. 2. 20 cm3 distilled water was poured into the 50 cm3 graduated

cylinder.3. The 20 cm3 distilled water in the graduated cylinder was poured

into the 100 cm3 glass beaker.4. Dropping pipette was used to take little amount of 20 cm3 distilled

water in the graduated cylinder into the test tube which then was labeled with a label.

5. A new dropping pipette was used to take little amount of 100 cm3

tap water inside the 100 cm3 glass beaker into three different test tubes which then were labeled with a label; Test tube 1, 2, and 3.

6. A spatula was used to take some solid AgNO3(s), Silver Nitrate into the first test tube which was previously filled with tap water.

7. The changes that was formed due to the addition of the substance was recorded in the data table shown below.

Water Analysis Part 1: QUALITATIVE DATA

No. Test Tube Observation

1 Cloudy precipitate which indicates the presence of Cl-

(Chloride ion) in the tap water sample

2Cloudy precipitate but less cloudy as the precipitate in the first test tube. This indicates the presence of SO42- (Sulphate ion) in the tap water sample

3 No precipitation and strong, pungent smell which indicates the absence of NH3 (ammonia) in the tap water sample

Image

2. Step 2: Quantitative Water Analysis 1. 50 cm3 EDTA, Ethylene Diamino Tetra acetate with concentration of

5 x 10-3 M was diluted into 50 cm3 of 5 x 10-4 M in a by 10 cm3 volumetric flask by adding 10 cm3 of 5 x 10-3 M EDTA which was taken by 10 cm3 volumetric pipette and 90 cm3 of distilled water, H2O(l).

2. The titration set up was set according to the figure below.

3. Funnel was being attached to the end of the burette to let the solution to flow easily.

4. 50 cm3 of 5 x 10-4 M EDTA solution was poured into the burette until it hits the 50 cm3 mark present on the burette.

5. This was recorded as the “initial burette reading” spot in the data table as shown below

Water Analysis Part 2: QUANITATIVE DATATRIALS

1 2 3Initial burette reading (mL) ± 0.05 mL 50.00 31.00

End Point /Final burette reading (mL) ± 0.05 mL 31.00 15.20Volume of EDTA (mL) ± 0.05 mL 19.00 15.80

Figures

6. The burette was slowly attached back to the clamp. 8. 25 cm3 distilled water was taken from the 100 cm3 glass beaker by

using 50 cm3 graduated pipette, and then was poured into the 50 cm3 titration flask (Erlenmeyer flask) labeled as flask 1.

9. 20 drops of Ammonia buffer with a pH value of 10 and 7 drops of EBT indicator were added into the Erlenmeyer flask used in the 8th step causing the solution to turn color into purple.

10.The 8th and the 9th step of the procedure were repeated twice to two other 50 cm3 titration flask (Erlenmeyer flaks), which were labeled as flask 2 and 3.

7. The first 50 cm3 titration flask (Erlenmeyer flask 1) was placed underneath the burette above a white paper.

8. 50 cm3 of 5 x 10-4 M EDTA solution was added slowly from the burette to the tap water in the Erlenmeyer flask.

9. The Erlenmeyer flask was swirled in order to efficiently mixed the chemical, when EDTA was added into he flask.

10.The 8th step was conducted continuously until the purple color of the solution started to change into blue.

11.50 cm3 of 5 x 10-4 M EDTA solution was added drop-wise from the burette to the tap water in the Erlenmeyer flask until the blue solution shown below was formed.

12.The volume of 5 x 10-4 M EDTA left in the burette was recorded in the blank “final burette reading” spot on the data table of the 1st trial and the “initial burette reading” spot in the 2nd trial column.

13.The seventh up to the twelfth step of the titration procedure was repeated but the flask used was changed in to the “Flask 2” and the “final burette reading” data was recorded in the 2nd trial column.

14.The volume of the titrant left in the burette was recorded in the blank “final burette reading” spot on the data table and the “initial burette reading” spot in the 3rd trial column

15.The seventh up to the twelfth step of the titration procedure was repeated once more but the flask used was changed in to the “Flask 3” and the “final burette reading” data was recorded in the 3rd trial column

16.The volume of EDTA used in titration including the color that was produced from the three trials of the titration procedure was compared, and the closest value and color was chosen as the two data used in the experiment.

V. Observation (Data)

Water Analysis Part 1: QUALITATIVE DATANo. Test

Tube Observation

1 Cloudy precipitate which indicates the presence of Cl- (Chloride ion) in the tap water sample

2Cloudy precipitate but less cloudy as the precipitate in the first test tube. This indicates the presence of SO42- (Sulphate ion) in the tap water sample

3 No precipitation and strong, pungent smell which indicates the absence of NH3 (ammonia) in the tap water sample

Image

Water Analysis Part 2: QUANITATIVE DATATRIALS

1 2Initial burette reading (mL) ± 0.05 mL

50.00 31.00

End Point /Final burette reading (mL) ± 0.05 mL

31.00 15.20

Volume of EDTA (mL) ± 0.05 mL 19.00 15.80

Figures

VI. Discussion

QUALITATIVE ANALYSIS

As you can see from the figure on the previous page, it is shown that there are 5 different test tubes that are used in this experiment, 4 of them were filled with tap water while the other one is filled with distilled water. Three of the test tubes were tested with different types of substances; test tube 1 with solid Silver Nitrate (AgNO3(s)), test tube 2 with solid Barium Chloride (BaCl2(s)), and finally test tube 3 with flakes of Sodium Hydroxide (NaOH(s)).

Firstly, when the first test tube was tested with solid Silver Nitrate (AgNO3(s)), it immediately forms a cloudy precipitate, which makes the water murky. The presence of these precipitate indicates the presence of Cl- (Chloride) ion in the form of of rigid, white precipitate of Silver Chloride (AgCl(s)). Silver Chloride (AgCl(s)) will exist as an insoluble solid (precipitate) because it is belong to the insoluble group of salts, like what is stated by the solubility rules, “all chlorides, bromides, and iodides (salts containing, Cl-, Br-, or I-) are soluble except when combined with Ag+, Pb2+, and Hg22+.” Its limited ability to be dissolved completely in a solution causes this compound to be undissolved, forming tiny white particles on the bottom of the test tube. In fact Silver Chloride (AgCl(s))

AgCl + NO3- AgNO3 + Cl-

does dissolved in a solution, but since its solubility value is so small and neglectable value which only reaches.0.002 g/L, Silver Chloride (AgCl(s)) will still be considered as an insoluble salt as a saturated solution will contains only about 1.3 x 10-5 moles of Silver Chloride (AgCl(s)) per liter of water.

Secondly, the second test tube containing tap water was tested with Barium Chloride (BaCl2(s)). The result of this experiment that is presented in the data table shows that when Barium Chloride (BaCl2(s)) was added into the tap water sample it does form precipitate, however, in a lesser amount than the first test tube which was tested with solid Silver Nitrate (AgNO3(s)), Barium Sulfate (BaSO4

(s)) was undissolved in the solution, thus forming precipitate due to two main reasons. According to the solubility rules of soluble and insoluble compound, it is stated, “all sulfates (salt containing SO42-) are soluble except those which containing Pb2+, Ca2+, Ba2+, Sr2+, Ag+, and Hg22+.” Barium Sulfate (BaSO4 (s)) is included in this exceptional condition due to the fact that solubility of sulfates (particularly those containing ion from Group II) decreases down the group hence making this compound to has a very small value of solubility compared to other elements belong to the same group.

Finally the third test tube was tested with some flakes of Sodium Hydroxide (NaOH(s)). However, it is seen from the result of the experiment that the addition of this substance did not form any precipitate to the tap water sample, which would mean the absence of NH3 in the water sample. Qualitatively, the absence of NH3 in the water sample can also be predicted from the absence of pungent smell that formed when some flakes of Sodium Hydroxide (NaOH(s)) was added.

QUANTITATIVE ANALYSIS - Determining the Hardness of Water SampleMany industrial and domestic water users are concerned about the hardness of their water. Hard water is water that has high mineral content, whereas soft water is water that has low mineral content. Water is an excellent solvent and readily dissolves minerals it comes in contact with. As water moves through soil and rock, it dissolves very small amounts of minerals and holds them in solution. Calcium and magnesium dissolved in water are the two most common

BaSO4 + Cl- BaCl2 + SO42-

NH4Cl + NaOH NaCl + H2O + NH3

minerals that make water "hard." This is formed when water percolates through deposits of calcium and magnesium-containing minerals such as limestone, chalk and dolomite. The hardness of water is referred to by three types of measurements: grains per gallon, milligrams per liter (mg/L), or parts per million (ppm). In the case of this experiment, Calcium ion (Ca2+) was the targeted ion to be calculated in order to determine the hardness of the water sample. The table below is provided as a reference.

Water Hardness Scale

Grains Per Gallon

Milligrams Per Liter (mg/L) or Parts Per

Million (ppm)Classification

less than 1.0 less than 17.1 Soft1.0 - 3.5 17.1 - 60 Slightly Hard3.5 - 7.0 60 - 120 Moderately Hard

7.0 - 10.5 120 - 180 Hardover 10.5 over 180 Very Hard

The method that is being used to calculate the concentration of calcium ion (Ca2+) in the sample is by complexometric titration with EDTA (Ethylene Diamino Tetra Acetate). EDTA is an example of a multidentate (many-toothed) ligand. It is used in this experiment because it can bind metal ions through multiple atoms which form a very stable complex. In the case of this experiment, EDTA is used to bind calcium ion (Ca2+) in the tap water sample according to the reaction shown below

The equation shown above is used as the basic of the titration used in this experiment, and in order to find the concentration of calcium ion (Ca2+) in the sample 25 cm3 of the tap water sample is titrated with 50 cm3 EDTA that is dropped slowly from the burette. However since neither EDTA4-, nor Ca+2 nor its complex, Ca(EDTA)2- is colored a pH indicator, Eriochrome Black T (EBT) is used to indicate the equivalence point of the titration and reacted together in the reaction shown below

Ca2+ + EDTA4-

[CaEDTA]2-

Calcium and EDTA complex structure

EBT is basically a pH indicator with a blue color, which is used in this experiment to detect the presence of calcium ion (Ca2+) in the 25 cm3 of the tap water sample. In this titration process it is seen from the reaction as well as the reaction above that before the end point the solution containing EBT indicator and tap water sample is purple in color however the color changes into blue during the equivalence point. The solution changes from purple into blue color at the equivalence point because all the calcium ion (Ca2+) has been reacted therefore causing them to be tied up with EDTA forming Ca(EDTA)2- complex. This causes EBT molecules that also have the ability to form complex with calcium ion (Ca2+) to be free and unreacted, which is represented with the blue color of the solution. Whereas at first, when EDTA has not been added to the sample calcium ion (Ca2+) are free and only being tied to the EBT molecules forming Ca-EBT complex that is represented as the purple color of he solution. However since Ca-EBT complex is much less stable than Ca-EDTA complex, when EDTA is added to the sample Ca(EDTA)2- complex will form instead of Ca-EBT complex.

Before equivalence point

At equivalence pointColor changes from

purple (right) Blue (left)

In addition, ammonia buffer solution with pH 10 is also added to the solution because buffer maintains the pH of the solution through out the reaction. To maintain high alkaline medium ammonia buffer is added to EDTA in analysis of hard water. It is necessary to keep the pH at about 10 for two reasons, firstly all reactions between metal ions and EDTA are pH dependent, which required the

InCa(aq)  In2-(aq)  +  Ca2+(aq) (Purple) (Blue Endpoint)

solutions must be kept basic (and buffered) for the reaction to go to completion as EDTA molecule can only form a very stable complex with calcium ion (Ca2+) in the sample when its structure is completely deprotonated with a charge of 4- Secondly, EBT indicator requires a pH of 8 to 10 for it to function as an indicator during this titration.

Concentration of The Calcium ion (Ca2+) in the Tap Water Sample

Volume of EDTA = 19.00+15.802 = 17.40 ml

Concentration of EDTA = 5.00 x 10-4 molL

Moles of EDTA = ¿)×17.40ml = 0.0087 mmol

Reaction: Ca2+ + EDTA4- CaEDTA2-

Based on the equation above, Moles of EDTA = Moles of Ca2+

Moles of Ca2+ = 0.0087 mmolMass of Ca2+ = 0.0087mmol ×40 g

mol = 0.348 mg

Concentration of Ca2+ = 0.348mg25×10−3 L

= 13.92 ppm

So, the concentration of the calcium ion (Ca2+) in the tap water sample is 13.92 ppm

The concentration of the calcium ion (Ca2+) in the tap water sample is 13.92 ppm. In this case, the hardness of the tap water sample used in this experiment is considered as soft water, which did not contain high mineral content or high amount of calcium ion (Ca2+). The category of this tap water that is not included as hard water concludes it as a safe water sources to be used in both industrial and household. It is because the lower the hardness level of a water would mean the smaller possibilities for these mineral deposits that are found inside the water to form insoluble precipitate which can clog the inside of appliances, such as dishwashers, and pipes and also forming clogs in appliances which reduce the heat-transfer efficiency, hinder the flow of water, and therefore decrease the quality of water.

VII. Conclusion In our experiment, there are two main substances that is present in our sample. Firstly the presence of Cl- (Chloride) ion in the form of of rigid, white precipitate of Silver Chloride (AgCl(s)) is indicated by the precipitate that makes the water murky When solid Silver Nitrate (AgNO3(s)), was added. Secondly, the presence of sulfate ion SO42- in the form of Barium Sulfate (BaSO4 (s)) precipitate when Barium

Chloride (BaCl2(s)) was added into the tap water even though it is in a lesser amount than the first test tube which was tested with solid Silver Nitrate (AgNO3(s)). Lastly, we can conclude that ammonia (NH3) does not present in the tap water sample due to the absence of the pungent smell that is formed when some flakes of Sodium Hydroxide (NaOH(s)) was added. In terms of water hardness, our tap water sample is considered as soft water instead of hard. It is due to the low concentration of the calcium ion (Ca2+) in the tap water sample, 13.92 ppm which is way below the range of ppm for the slightly hard water, 17.1 ppm. Due to the low concentration of the calcium ion (Ca2+) we can conclude that it as a safe water sources to be used in both industrial and household. It is because the lower the hardness level of a water would mean the smaller possibilities for these mineral deposits that are found inside the water to form insoluble precipitate which can clog the inside of appliances, such as dishwashers, and pipes and also forming clogs in appliances which reduce the heat-transfer efficiency, hinder the flow of water, and therefore decrease the quality of water.

VIII. References

Sutanto, Hery, and TablighPermana. Inorganic and Organic Chemistry 2 Laboratory Manual.Tangerang: Swiss German University, 2013. Print.

"COMPLEXOMETRIC TITRATION OF CALCIUM." Stony Brook University. N.p., n.d. Web. 4 Dec. 2014. <http://www.ic.sunysb.edu/Class/che134/lectures/134p03.pdf>.

"Determination of Water Hardness By Complexometric Titration Class Notes." Determi-nation of Water Hardness By Complexometric Titration Class Notes. N.p., n.d.

Web. 11 Apr. 2014. <http://homepages.ius.edu/DSPURLOC/c121/week13.htm>."EDTA equations." EDTA equations. N.p., n.d. Web. 13 Apr. 2014. <http://genchem.rut-

gers.edu/edtaeq.html>."Explanation of Water Hardness." Explanation of Water Hardness. N.p., n.d. Web. 10

Apr. 2014. <https://www.fcwa.org/water/hardness.htm>."How does EDTA Chelation therapy work? (Part one)." Good Life Labs. N.p., n.d. Web.

13 Apr. 2014. <http://goodlifelabs.com/how-does-edta-chelation-work-1/>.Sawyer, John. "Surface Waters: Ammonium is Not Ammonia – Part Two." Ammonium

and Ammonia in Drinking Water. N.p., n.d. Web. 13 Apr. 2014. <http://www.ex-tension.iastate.edu/CropNews/2008/0502JohnSawyer.htm>.

"Sulfate in Drinking Water." Home. N.p., n.d. Web. 11 Apr. 2014. <http://water.epa.-gov/drink/contaminants/unregulated/sulfate.cfm>.

"Water Hardness." Water Hardness. N.p., n.d. Web. 10 Apr. 2014. <http://www.chem-istry.wustl.edu/~edudev/LabTutorials/Water/FreshWater/hardness.html

"What's in your drinking water?." Tap Water Content. N.p., n.d. Web. 11 Apr. 2014. <http://freshlysqueezedwater.org.uk/waterarticle_watercontent.php>.

"Why you use pH 10 buffer in EDTA titration?." WikiAnswers. Answers Corporation, n.d. Web. 13 Apr. 2014. <http://wiki.answers.com/Q/Why_you_use_pH_10_buffer_in_EDTA_titration?#slide=1>.