Etiket Arşivleri: FE 211


FE 211 06.11.2004



Submitted By : Mutlu DEMİREL Group: 3


NaCI soln , nitric acid , ferric alum , , nitrobenzene , ammonium thiocyonate , potassiyum chromate , distilled water , burette , pipette , beaker , florence flask , graduated cylinder


Our aim was to determine percentage of the chloride ion in the sample by Volhard and Mohr methods.


Volhard method , Mohr method


Volumetric methods based upon the formation sparingly soluble silver salts are among the oldest known ; these procedures were and still are routinely employed for the analysis for silver as well as for the determination of such ions as chloride , bromide , visually detectable change in the solution.

For example , the formation of a second precipitate of distinctive color is the basis for end-point detection with the Mohr method. The formation of colored complex is used for the end-point detection with the Volhard method . Adsorption of some organic dyes can be used in end-point detection with fejans method.

Most application of precipitation titrations are based upon the use of a standart silver nitrate solution and are sometimes called argentometric methods.


Volhard Method :

10 mL of NaCI sample was added to 100 mL distilled water. Then 1 mL nitric acid , 5 mL ferric alum solution , 10 mL 0.1M and 15 mL nitrobenzene added to this solution respectively. After that the sample solution was titrated with ammonium thiocyanate until red color persist.

Mohr Method:

10 mL of NaCI sample was added to 50 mL distilled water. The 2 mL potassium chromate solution was added to it . After that the sample solution was titrated with 0.1M .



The experiment is based upon the determination of the chloride ion with the precipitate titrations. In the precipitate titrations end –point is detected by chemical indicators which procedures a visual detectable change , usually of color or turbity in the solution. The three of types of visual indicator are Mohr , volhard and fejans methods. In the experiment Mohr method and Indirect Volhard method is applied.

In the indirect Volhard method an ions such as are precipitated with excess silver ion , then the is back titrated with a potassium thiocyanate solution. The precipitates of AgI , AgBr and AgSCN are need to be removed from the solution before back titration because AgCI is more soluble than AgSCN , AgCI should be removed by filtration before back titration. But in order to remove AgCI we used nitrobenzene to mask the silver chloride.

In Mohr method chromate ion is used as the indicator in the titration of chloride with silver . In Mohr method , adjusement of concentration of chromate ion and the pH of the solution are important factors to care about. The pH must be about 7-8 .If it is lower chromate ion will change into dichromate and use more then needed silver ion to reach equivalent point. The concentration of chromate should bu usually about 0.005 M . If the conc. of chromate is high the yellow color of chromate ion cancels the appearences of the red colored silver chromate precipitate.

In the experiment gauss errors are possible. While reading the menicus personal judgement is important

Experiment-7 , 8

FE 211 23.11.2004



Submitted By : Mutlu DEMİREL Group: 3


Distilled water , , , , Volumetric Flask , Erlenmayer Flask , KI , HCI , , , KSCN , , starch , burette , graduated cylinder , florence flask , pipette


Determination of copper in a sample by iodometry and determination of oxygen in water by iodometry.


Iodometry method.


• Iodometry: starch added only immediately before the equivalence point. Titration is then continued until the blue disappears

• Iodimetry: starch added at the beginning

The iodine (triiodide) – iodide redox system, I3 oxidizing agent. On the other hand, iodide ion can be a strong reducing agent. Analytical reactions using iodine as the oxidizing agent are called iodimetry while procedures using iodide ion as the reducing agent are called iodometry. Since triiodide is such a weak reducing agent, there are few iodimetric determinations of analytical interest. Iodide ion is a strong enough a reducing agent that many oxidizing agents can react completely with the iodide ion resulting in many useful iodometric processes. The usual procedure involves the addition of an excess of iodide ion to the oxidizing agent analyte which produces iodine, which can be titrated with standard sodium thiosulfate solution. The iodine-thiosulfate reaction is quite fast and the equilibrium is far to the product side. Iodine is slightly soluble in water (0.00134 mol/L at 25 oC but is soluble in solutions containing iodide ion. Iodine forms the triiodide complex with iodide, I2 + II3 with an equilibrium constant greater than 500 at 25 oC. Normally, excess potassium iodide is added to the reaction mixture to increase the solubility of iodine and to decrease its volatility. It is possible to prepare primary standard iodine solutions by direct weighing.

Sodium Thlosulfate

Sodium thiosulfate is commonly used as the pentahydrate, Na2S2O3•5 H2O, and

+ 2 e- 3 I- – , is a weak

solutions are standardized against a primary standard. Thiosulfate solutions are not

stable over long periods of time, and frequently borax or sodium carbonate is added

as a preservative. Iodine oxidizes thiosulfate to the tetrathionate ion:

I2 + 2 S2O3

The reaction of iodine with thiosulfate is fast and goes to completion. The equivalent weight of Na2S2O3•5 H2O is the molecular weight, 248.17, since one electron per molecule is lost. Iodine is rarely used as a primary standard for thiosulfate because it presents a problem in weighing and maintaining its solution concentration. Both potassium bromate and potassium iodate oxidize iodide quantitatively to iodine in acid solution:

+ 5 I- IO3 – + 6 I- BrO3 –

The bromate reaction is slow but can be speeded up by increasing the hydrogen ion concentration. The iodate reaction is rapid and it requires only a slight excess of hydrogen ions for complete reaction. The main disadvantage of these two salts being used as primary standards is that their equivalent weights are quite small. In both cases the equivalent weight is one-sixth of the molecular weight. Pure copper is used as a primary standard for sodium thiosulfate and is recommended when the thiosulfate is to be used for the determination of copper. Upon addition of excess iodide to a solution of Cu(II), a precipitate of CuI is formed along with I2. The liberated iodine is then titrated with standard sodium thiosulfate.

+ 4 I- 2 Cu2+ I2 + 2 S2O3

Formation of the precipitate and the addition of excess iodide forces the equilibrium to the right. The pH of the solution must be maintained between pH 3 and 4 by the addition of a buffer. It has been found that iodine is adsorbed onto the surface of the copper(I) iodide precipitate and must be displaced to obtain correct results. Potassium thiocyanate is usually added just before the end point is reached to displace the adsorbed iodine. Note that the net stoichiometry of the react is 1:1 since 2 moles of copper requires 2 moles of thiosulfate.


15 mL of Cu containing sample was diluted to 50 mL. Then 1 mL and 4 gr KI was added to it. The solution was titrated with thiosulfate .During the titration 2 gr KSCN and 5 mL starch was added to solution. Then the titration was continued untill the blue color was disappeared. 1 mL and 1 mL alkaline KI were added to 100 mL water. The formation of precipitate was allowed. When the brown precipitate was obtained 1.5 mL was added to the solution and and mixed. The solution was titrated with

2 I- -2 + S4O6 -2 + 6 H+ 3 I2 + 3 H2O + 6 H+ 3 I2 + Br- + 3 H2O 2 CuI(s) + I2 2 I- -2 + S4O6 -2

thiosulfate adding starch. İndicator just before the end point.


Determination of Copper by iodometry:

Volume of used 16.2 mL

Molarity of 0.1 mL =

= moles

Determination of oxygen in water:

Volume of used 3.0 mL

weight of =

mg per L=


In this experiment the oxygen in water was determined by iodometry. To determine oxygen in water polarographic methods an be used as dropping Mercury electrode or simply by using dissolved oxygen meter. But the determination is done by Winkler method that is based on the oxidation of by oxygen which is reduced by iodide in an acidic medium.Later the iodine formed s determined by titration with standart thiosulfate.

In this experiment the amount Cu in a sample was found by iodometric method. In acid solution practically all oxidizing agents will oxidize iodide ion to iodine quantitatively. The iodine formed in the reaction can then be titrated by means of a standard sodium thiosulfate solution. This type of indirect titration is given the general name of iodometry. Iodometric methods of analysis have a wide applicability for the following reasons:

1. Potassium iodide, KI, is readily available in high purity.

2. A good indicator, starch, is available to signal the equivalence point in the reaction between iodine and thiosulfate. Starch turns blue-black in the presence of iodine. Therefore, when the blue-black color disappears, the iodine has been completely reduced to the iodide ion.

3. Iodometric reactions are rapid and quantitative.

4. A precise and stable reducing agent, sodium thiosulfate (Na2S2O3), is available to react with the iodine. The amount of iodine liberated in the reaction between iodide ion and an oxidizing agent is a measure of the quantity of oxidizing agent originally present in the solution. The amount of standard sodium thiosulfate solution required to titrate the liberated iodine is then equivalent to the amount of oxidizing agent. Iodometric methods can be used for the quantitative determination of strong oxidizing agents such as potassium dichromate, permanganate, hydrogen peroxide, cupric ion and oxygen.

As has been mentioned above, the endpoint in a titration of iodine with thiosulfate is signaled by the color change of the starch indicator. When starch is heated in water, various decomposition products are formed, among which is beta-amylose which forms a deep blue-black complex with iodine. The sensitivity of the indicator is increased by the presence of iodide ion in solution. However, if the starch indicator solution is added in the presence of a high concentration of iodine, the disappearance of the blue-black color is very gradual. For use in indirect methods, the indicator is therefore added at a point when virtually all of the iodine has been reduced to iodide ion, causing the disappearance of the color to be more rapid and sudden. The starch indicator solution must be freshly prepared since it will decompose and its sensitivity is decreased. However, a properly prepared solution will keep for a period of a few weeks. A preservative such as a small amount of mercuric ions may be added to inhibit the decomposition. Solutions of sodium thiosulfate are made up to an approximate concentration by dissolving the sodium salt in water that has previously been boiled. Boiling the water is necessary to destroy micro-organisms which metabolize the thiosulfate ion. A small amount of Na2CO3 is added to the solution in order to bring the pH to about 9.

The solution is standardized by taking a known amount of oxidizing agent, treating it with excess iodide ion and then titrating the liberated iodine with the solution to be standardized. Oxidizing agents such as potassium dichromate, bromate, iodate or cupric ion can be employed for this procedure. You will be using potassium iodate, KIO3, as your primary standard. The reaction between IO3 6H++IO3 Iodometric methods depend on the following equilibrium: and I- – is given as — >3I2+3H2O -+5I- I2 + I-< ===> I3 –

Since the solubility of I2 in water is quite low, the formation of the tri-iodide ion, I3 allows us to obtain useful concentrations of I2 in aqueous solutions. The equilibrium constant for this reaction is approximately 700. For this reason iodometric methods are carried out in the presence of excess iodide ion. The reaction between iodine and the thiosulfate ion is:

I2 + 2S2O3 < ===> 2I- 2- + S4O6 2-

This reaction proceeds quantitatively in neutral or slightly acidic solutions. In strongly alkaline or acidic solutions the oxidation of the thiosulfate does not proceed by a single reaction. In the former, the thiosulfate ion is oxidized to sulfate as well as to the tetrathionate. In the latter, the thiosulfuric acid formed undergoes an internal oxidation-reduction reaction to sulfurous acid and sulfur. Both of these reactions lead to errors since the stoichiometry of the reactions differs from that shown above for the thiosulfate as a reducing agent. The control of pH is clearly important. In many cases the liberated iodine is titrated in the mildly acidic solution employed for the reaction of a strong oxidizing agent and iodide ion. In these cases the titration of the liberated iodine must be completed quickly in order to eliminate undue exposure to the atmosphere since an acid medium constitutes an optimum condition for atmospheric oxidation of the excess iodide ion.

The basic reaction in the determination of copper using the iodometric method is represented by the equation: 4I- 2Cu2+ < ===> 2CuI(s) + I2

This is a rapid, quantitative reaction in slightly acidic solutions, if there is a large excess of iodide ion present and if the copper is in the form of a simple ion rather than a complex one. The iodine that is liberated can be titrated in the usual manner with standard thiosulfate solution. The reaction involving cupric ion and iodide takes place quantitatively since the cuprous ion formed as result of the reduction is removed from the solution as a precipitate of cuprous iodide. Iron interferes since iron(III) ions will oxidize iodide. Since the iron will be found in the +3 oxidation state as a result of the dissolution of the brass sample, a means of preventing this interference is necessary. This can be accomplished by converting the iron(III) to a soluble iron(III) phosphate complex using phosphoric acid. At a pH of 3.0-4.0 the iron phosphate complex is not reduced by iodide ion. If arsenic and antimony are present, they will provide no interference at this pH if they are in their higher oxidation states.

The following are the most important sources of error in the iodometric method:

1. Loss of iodine by evaporation from the solution. This can be minimized by having a large excess of iodide in order to keep the iodine tied up as tri-iodide ion. It should also be apparent that the titrations involving iodine must be made in cold solutions in order to minimize loss through evaporation.

2. Atmospheric oxidation of iodide ion in acidic solution. In acid solution, prompt titration of the liberated iodine is necessary in order to prevent oxidation.

3. Starch solutions that are no longer fresh or improperly prepared. The indicator will then not behave properly at the endpoint and a quantitative determination is not possible.


FE 211 09.11.2004



Submitted By : Mutlu DEMİREL Group: 3


Ca containing sample , beaker , pipette , buret , bunsen burner , florence flask , HCI , , Methyl Orange , funnel , thermometer , ,


Determination of Calcium in a sample.


Volumetric method.


Calcium dissolves out of almost all rocks and is, consequently, detected in many waters. Waters associated with granite or siliceous sand will usually contain less than 10 mg of calcium per litre. Many waters from limestone areas may contain 30-100 mg l-1 and those associated with gypsiferous shale may contain several hundred milligrams per litre. Calcium contributes to the total hardness of water. On heating, calcium salts precipitate to cause boiler scale. Some calcium carbonate is desirable for domestic waters because it provides a coating in the pipes which protects them against corrosion.


5 mL Ca containing sample and 10 mL of HCI were poured in a beaker and diluted to 150 mL . Then the solution was heated to remove . After that 50 mL warm and 3-4 drops methyl orange were put into solution. We waited for hours to observe the precipitate. Then the solution was filtered and washed to remove the precipitate. The precipitate was put into flask with filter paper. 150 mL distilled water and 50 mL 6M were added to flask. Soon the solution was heated to 60 to dissolve ppt. Finally the hot solution was titrated with 0.02M


Volume = 21.3 mL

Molarity of =10 g/L

Volume of sample = 5 mL

Density of sample = 10 g/L

Number of mole = mol

Number mole of


Weight =

weight of sample =

Percentage of =


In this experiment calcium is determined by an indirect volumetric method . The method is based on precipitating ions with oxalate as . With this manner was added the solution and calcium oxalate is formed : There were some important factors to apply during the procedure. Addition of HCI must be done slowly to avoid loss by splattering due to the evalution. Our aim was to drive off the while we heat the solution. Because is acidic and may effect our result. While filtrering the solution Watmann 42 paper was used . This filter paper ha narrow pores. So the loss of precipitate was prevented. Then the precipitate was put into a beaker with filter paper , and 6M was added to it. This was done to coprecipitate the again. Finally our solution was titrated with 0.02M until the pink color was observed. The used was used for our calculation. During the procedure HCI was not used to solve our Ca containing sample. Also all titration procedures have an error becaus of the differance between the volume of equivalence point and the end point.


FE 211 01.11.2004



Submitted By : Mutlu DEMİREL Group: 3


HCl , , , , , , Zimmermann-Reinhardt reagent , Distilled water , Buret , Erlenmayer Flask , Graduated Cylinder , Pipet , Bunsen Burner.


Our purpose is to determine the iron by permanganete titration.


Gravimetric Analysis


Iron is an abundant element in the earth’s crust, but exists generally in minor concentrations in natural water systems. The form and solubility of iron in natural waters are strongly dependent upon the pH and the oxidation- reduction potential of the water. Iron is found in the +2 and +3 oxidation states. In a reducing environment, ferrous (+2) iron is relatively soluble. An increase in the oxidation-reduction potential of the water readily converts ferrous ions to ferric (+3) and allows ferric iron to hydrolyse and precipitate as hydrated ferric oxide. The precipitate is highly insoluble. Consequently, ferric iron is found in solution only at a pH of less than 3. The presence of inorganic or organic complex- forming ions in the natural water system can enhance the solubility of both ferrous and ferric iron.


Standardization of Permanganate Solution with Sodium Oxalate: 10 ml of sodium oxalate ( ) was put into a flask and 25 mL was added to it. Then solution was heated nearly to boiling. Then the solution was titrated with permanganate until the pink color was observed. Finally the molarity of permanganate solution was calculated.

Iron Determination:

5 mL of HCI was added to 10 mL sample in a flask. Then the solution was heated. After , stannous chloride ( ) was added to the hot solution drop by drop until the yellow color changes to light green and two drops of excess was added. Then the solution diluted to 100 mL and allowed to cool at room temperature. After that 10 mL was added to this solution rapidly with stirring. When the white ppt was observed. The solution diluted to 300 mL and 25 mL Zimmermann – Reinhardt reagent was added to it. Finally the solution was titrated with permanganate ( ) .Untill the ping color was observed and %Fe was calculated in the sample.


Standardization of Permanganate with Sodium Oxalate:

the volume of = 10 mL

the of = 20 g/L

the volume of permanganate = 38 mL

the molecular weight of =134 g

weight of =

Mole of =

Molarity =

Percentage of Fe ion:

the volume of sample = 5 mL

the of sample = 60 g/L

the volume of permanganate = 13 mL

the molarity of permanganate = 0.02 M

the mole of permanganate =

weight of =

weight of sample =


Volumetric determination of iron is basicly based on the dissolution of the sample reduction of to , addition of special reagent and titration with standart permanganate. But instead of permanganate we should use create or dichromate solution as titrant. Firstly we calculated the molarity of permanganate which we used as standart solution during the determination of Fe. Later we prepared the titrand. During the experiment we heated the sample solution but didn’t boil it. This is because the prevent the loss of as vopar. Then we added stannous chloride to the solution. Our aim was to reduce ions into .Because is more soluble form of . We added 1-2 drops more to minimaze the air oxidation of and to eleminate the interference of impurities. After that we added 10 mL to the solution to oxidize the excess according to reaction: A small quantity of silky white precipitate of was appeared. İf the precipitate was gray it means that the was to much and changed the to Hg This Hg ions reacts with with slow rate , consuming more than required to reach the end point .İt means you failed the experiment . For this reason after adding white precipitate should be observed. In this case of no precipitation formation , after addition of it is observed that the was not in excess and most probably reduction of to is incomplete. Therefore , the sample solution should be discarded. If either gray precipitate formed or no precipitation observed. In our experiment white precipitate was observed. Soon we added 25 mL Zimmermann-Reinhardt reagent this reagent eliminates the interference of chloride with titration , Finally we titrated the sample solution with permanganate solution untill the pink color observed.





(aq) + 5e-

+ 5e-

+ 8H+

(aq) Fe3+

+ 8H+

+ 5e-

(aq) 5Fe3+

+ 8H+

+ 5Fe2+




+ 4H2O

(aq) + e

+ 4H2O

+ 5Fe3+ + 4H2O


FE 211 25.10.2004



Submitted By : Mutlu DEMİREL Group: 3


HCl , HAc , Phenolphthalein , Methyl Orange , NaOh , Dilute Water ,Buret ,

Erlenmayer Flask ,


In this experiment we titrated HCI and HAc mixture with NaOH to determine

amounth of HCI , HAc and NaOH and the molarity of NaOH.


Gravimetric Analysis


The capacity of water to accept protons is called alkalinity. Alkalinity is important in water treatment and in the chemistry and biology of natural waters. Alkalinity serves as a pH buffer and reservoir for inorganic carbon. It helps to determine the ability of water to support algal growth and other aquatic life; thus, it can be used as a measure of water fertility. In natural waters, the species responsible for alkalinity are OH-, HCO3-, and CO32-. Alkalinity is equal to [OH-] + [HCO3-] + 2[CO32-]. The coefficient of “2” before [CO32-] is necessary because carbonate accepts two protons. However, in natural waters of pH from 7 – 9, the predominant species is bicarbonate.

Alkalinity can be measure by titration with a standard solution of HCl. The titration reaction for the bicarbonate ion is: HCO3- + H+ _ H2CO3 The endpoint can be determined with the methyl orange indicator or by measuring the pH of the solution throughout the titration. In the later case, analysis of the titration curve allows determination of the end point. Alkalinity is expressed as the number of moles of H+ required to titrate one liter of water sample or as mg CaCO3/L of water.


5 ml HCI and 5 ml HAc was poured into a erlenmayer flask. Then 2-3 drops of methyl orange indicator was added to this solution and titrated with standart solution until the color of the solution change. Then 2-3 drops of phenolphthalein was added to the solution and again titrated with NaOH until the end point was observed.


HCI = n NaOH fw of HCI = 36.5 g volume of NaOH = 8 mL

NaOH =

NaOH = n HCI = 0.0008 mol

fw of

volume of NaOH = 11 mL

NaOH = n

NaOH =

NaOH = n = 0.0011 mol


In this titration reaction we have a mixture of strong acid , HCI , and weak acid HAc , as primary solution. The mixture of acids were titrated with base , NaOh because the Ka for weak acid was abouth or less so two end point was observed. The first end point was for HCI , the second was for HAc . To observe the end points of this two different acids, two different indicators were used. Methyl orange for HCI and Phenolphthalein for HAc.


FE 211 18.10.2004



Submitted By : Mutlu DEMİREL Group: 3


Ashless Paper , , HcL , Water , , Bunsen Burner , Methyl Orange, , Erlenmayer Flask , , Porcelain Crucibles , Graduated Cylinder , Furnace


In this experiment we calculated the weight and % of the sample ,


Gravimetric Analysis


The gravimetric determination of iron in soluble salt is done by converting all iron to , precipitating as a ferrichydroxide and igniting it to . The solubility of ferrichydroxide is very low and therefore quantitative precipitation can be achieved even in acidic solutions of ph up to 4. However , the precipitate is gelatinous and can not be filtrated through sintered glass crucibles because the pores of the crucibles are easily clogged with these types of precipitate. Therefore the precipitate has to be filtered through a filter paper. If the air oxygen is not sufficent , the compound may change to or even metallic iron by carbon from filter paper or reducing gases thar may firm when washing the filter paper.


5 g sample was added 1 ml Hcl and diluted to 250 ml. Then pulled filter paper and 2-3 drops of was added to this solution. After heating just to boil , 3 drops of methyl orange and was added to the solution. After that the solution solution was allowed to precipitate. When the supernatant was become clear , precipitate was allowed to coalugate an filter pulp for 15-30 minutes. Later the solution filtrated and washed with hot wash liquid .After testing the fresh portion filtrate with the filter paper was folded and put into a crucible.Then the crucible was charred off . After that the crucible was ignited at 800-900 in the furnace and weighed. So the weight of was found.


The volume of sample : 5ml

The density of sample : 60 g/L

The molecular weight of Fe : 55.847 g

The molecular weight of : 159.699 g

The mass of crucible : 12.0746 g

The mass of crucible + sample : 12.15.95 g

The mass of sample : 0.0849 g


In this experiment we calculated the weight and % of iron (III) with gravimetric analysis. We used to oxidize ions into before precipitating the solution. Then we added to make the solution alkali and precipitate reaction is obseved as ; While filtering the solution we washed the precipitate with wash liquid and tested the fresh portion filtrate with to observe ions. Washing to completed when no turbitidy is observed. Then the crucible was put on the bunsen burner with low fire. Low fire because of the carbon from the filter paper may reduce to Fe The result of the experiment may have some errors because of the ions that couldn’t be eleminated.Also carbon atoms from the filter paper may reduce to Fe althoug being careful.So the determined weight may not be only pure .

Chapter – 2 – Problems

FE 211 – CHAPTER 2


Ex.1: Calculate the mean and median for each of the following sets of data.

Fe 211 I.Midterm Summary




Analytical Chemistry: involves separating , identifying and determining the relative amounts of the components making up a sample of matter.

Qualitative Analysis: which deals with the problem of what is in a substance

Quantitative Analysis: which handles the problem of how much of each constituent is present.

Central Value: The most commoncentral value used by the chemist is the MEAN.

Precision: is a measure of how close the experimental values are to each other , or term precision is used to describe the reproducibility of results.

Accuracy: is a measure of how close the the experimental results are to the accepted values and is expressed is term of error.

Gravimetric Factor: Stoichiometric relationship between the analyte and the product weighed.This constant is sometimes called the gravimetric factor.

Colloiadal Suspensions: whose tiny particles are invisible to the naked eye .

Crystalline Suspensions: The temporary dispersion of such particles in the liquid phase is called a crystalline suspension.

Density: of a substance measures its mass per unit volume.

Specific Gravity: of a substance is the ratio of its mass to the mass of an equal volume of water at

Volumetric or Titrimetric Method: A quantitative analysis based upon the measurement of volume is called a Volumetric or Titrimetric Method.

Titration: is a reaction carried out by the carefully controlled addition of one solution to another.

Titrant or Standard solution: is a reagent of known concentration that is used to carry out a volumetric analysis.

Titrand or Sample solution: which is often placed in an Erlenmeyer flask.

Equivalence Point: in a titration is reached when the amount of added titrant is chemically equivalent to the amount of analyte in the sample.

Back Titration: It is sometimes necessary to add an excess of the Standard titrant and than determine the excess amount of by back-titration with second Standard titrant.

End point: The point at which a physical change is observed at the sample is called the end point.

Titration error: The difference in volume between the equivalence point and end point is the titration error.

Indicator: A common method of detecting end points involves the introduction of a substance called an indicator.

Primary Standards: Every volumetric method is based upon a primary standard ,which is used directly or indirectly to establish the concentration of the standard solution.

Standard Solutions: Standard solutions play a central role in all volumetric methods of analysis.