LABORATORY WORK PRECIPITATION TITRATION. MOHRS METHODS

15

LABORATORY WORK
PRECIPITATION TITRATION.
MOHRS METHODS

Objective: Determination of chloride in solid and liquid samples by the Mohr Method

Learning Outcome:
Students understand the terms volumetric analysis, morarity, molality normality and redox titration.
Students acquire the skill to prepare standard solutions of silver nitrate and sodium chloride.
Students understand the apparatus used for a titration.
Students acquire the skill to perform the precicpitation-titration in the real lab after understanding the different steps.

Titration is a process by which the concentration of an unknown substance in solution is determined by adding measured amounts of a standard solution that reacts with the unknown. Then the concentration of the unknown can be calculated using the stoichiometry of the reaction and the number of moles of standard solution needed to reach the so called end point.
Precipitation titrations are based upon reactions that yield ionic compounds of limited solubility.
Classification of methods precipitation titration (on titrant):
Argentometry
Thiocyanatometry
Mercurometry
Sulphatometry
Hexacianoferratometry
Characteristics of Precipitation Titration:
They are fast and the stoichiometry is known and reproducible, (no secondary reactions of interference).
They are complete or can be quantified depending on the amount of solubility product (in general a precipitation titration is considered complete when Ksp < 10-8)
An indicator can be used to find the equivalence point or titration end point which, for this type of titration, corresponds to when precipitation of the analyte under analysis is complete.

The Solubility Product Constant
When an ionic compound is dissolved in water, it usually goes into solution as the ions. When an express of the ionic compound is mixed with water, equilibrium occurs between the solid compound and the ions in the saturated solution:
KxAy = xK+ + yA-
The equilibrium constant for this solubility process can be written:
Keq=[K+]x [A-]y[KxAy]
However, because the concentration of the solid remains constant (in heterogeneous systems), we normally combine its concentration with Keq to give the equilibrium constant Ks, which is called the solubility product constant:
Ks=Keq KxAy=[K+]x [A-]y
In general, the solubility product constant, Ks, is the equilibrium constant for the solubility equilibrium of slightly soluble (or nearly insoluble) ionic compounds. It equals the product of the equilibrium concentrations of the ions in the compound, each concentration raised to a power equal to the number of such ions in the formula of the compound.
At equilibrium in saturated solution of slightly soluble compound at given temperature and pressure the value of Ks is constant and not depend on ions concentration. The solubility product constant is thermodynamic constant and depends on temperature and ions activity (ionic strength).
The reaction quotient, Q, is an expression that has the same form as the equilibrium constant expression Ks, but whole concentration values are not necessarily those at equilibrium. Though the concentrations of the products are starting values: Q=[K+]x [A-]y
Here Q for a solubility reaction is often called the ion product, because it is product of ion concentrations in a solution, each concentration raised to a power equal to the number of ions in the formula of the ionic compound.
Precipitation is expressed to occur if the ion product Q for a solubility reaction is greater than Ks: Q > Ks.
If the ion product Q is less than Ks, precipitation will not occur (the solution is unsaturated with respect to the ionic compound): Q < Ks.
If the ion product Q equal Ks, the reaction is at equilibrium (the solution is saturated with the ionic compound): Q = Ks.

Calculation of solubility
Solubility, S, is the molar concentration of compound in saturated solution.
I. Saturated solution of slightly soluble ionic compound: S=x+yKsxx yy
II. Saturated solution of good soluble ionic compound.
This type of solutions not used in analytical practice. Such solutions are very concentrated and have large ionic strength. Components of these solutions (ion, molecules) can associate and form various polymers and colloids.
III. Saturated solution of slightly soluble compound with very small solubility:
the substance have limited solubility but create ion pairs and various molecular forms. The ionic strength of this solution is high and solubility depends on common concentration of all molecular and ionic forms;
slightly soluble compound takes part in protolytic reaction with water with the pH change.
The solubility is affected by pH. If the anion is the conjugate base of a weak acid, it reacts with H+ ion. Therefore, the solubility slightly soluble compound to be more in acid solution (low pH) than it is in pure water.
In sour environment solubility of slightly soluble compounds is more than more is its Ks and more is the hydrogen ion concentration:
SKxAy=K+=Ks[A-]=Ks (H+Ka+1)
when H+=Ka, SKxAy=2KS

Factors which influence to solubility:
1. Temperature. Solubility for most of substances is endothermic process. Increase temperature occurs decrease solubility. But crystal compounds at various temperature form hydrates another structure (composition). Hydrates formation may be exothermic reaction.
2. Ionic strength of solution. Increasing of ionic strength causes decreasing of ions activity and, accordingly, Ks will increase. Because, solubility will increase. An example of it is salting effect.
Salting effect is increase the solubility of slightly soluble compounds in presence of strong electrolytes, which not have common ions with precipitate and not react with precipitate ions.
3. Common-ion electrolytes. Completeness of precipitation.
The importance of the solubility product constant becomes apparent when we consider the solubility of one salt in the solution of another having the same cation or anion. The effect of the common ion is to make slightly soluble salt less soluble than it would be in pure water. This decrease in solubility can be explained in terms of Le-Chatelier's principle. It is example of the common-ion effect.
Decrease of solubility of slightly soluble compounds in presence of electrolyte with common ions called common-ion effect.
But solubility of slightly soluble compounds decrease to moment when ionic strength of solution will begin to influence to solubility.
The ion is completely precipitated when its residual concentration (Cmin) is less than 1×10-6 M (Cmin < 1×10-6 M). Amount of precipitant must be more at 20-50 % it is necessary to stoichiometry equation.
If in solution are ions, which form slightly soluble compounds with precipitant, the sequence of its precipitation determines (depends on) Ks value.
Fractional precipitation is the technique of separating two or more ions from a solution by adding a reactant that precipitates first one ion, than another, and so forth.
4. The pH value (see above).
5. Complex compound formation. Solubility increases with increasing concentration of ligand, complex compound stability and Ks value.
6. Redox process. Redox reaction shift on equilibrium in heterogeneous system and change solubility of slightly soluble compounds.

ARGENTOMETRY

The most important precipitating reagent is silver nitrate (AgNO3). Titrimetric methods based upon silver nitrate as titrant are sometimes termed argentometric methods. Potassium chromate can serve as an end point indicator for the argentometric determination of chloride, bromide and cyanide ions by reacting with silver ions to form a brick-red silver chromate precipitate in the equivalence point region.
Fields of argentometry application:
The determination of the anions Cl-, I-, Br- and Ag+ is also common.
Environment: Determination of chloride in water.
Food and beverage: Determination of chloride in finished products (cooked meats, preserves). Determination of chloride in dairy products.
Precious metals: Determination of silver in various alloys (for jewellery).
Pharmaceuticals: Titration of halides.
The reaction rates for the silver salt precipitation is rapid. The reaction ratio is 1:1 and silver salts formed are generally quite insoluble. Table below gives the solubility product, Ksp, for the silver salts that are involved in precipitation titrations.
Table 1 Solubility products for silver salts
Anion
Ksp
Solubility (g/100mL)
Cl-
1.8 x 10-10

Br-
5.2 x 10-13

I-
8.3 x 10-17

SCN-
1.1 x 10-12
0.00002
CrO42-
2.6 x 10-12
0.0025

Endpoint Detection for Argentometric Titration. Another requirement of titrimetric analysis is that there must be some method of determining when the titration reaction has reached its equivalence point. In this experiment, you will compare two methods of endpoint detection, one using a chemical indicator and another using potentiometric detection. We will now describe the common methods of endpoint detection for argentometric titrations.
Chemical Indicators. There are three common chemical indicators that are associated with argentometric titrations:
1. The chromate ion, CrO42- (the Mohr method);
2. The ferric ion, Fe3+ (the Volhard method);
3. Adsorption indicators such as fluorescein (the Fajans method).
The curve for the titration of 100 cm3 of 0.0100 M KCl with 0.100 M AgNO3 is shown in Fig.1.

You may note here that the titration curve is quite similar to the one for the titration between a strong base and a strong acid. According to this curve there is a sharp increase in the concentration of silver ions immediately after the equivalence point. This is indicated by a sharp decrease in the value of pAg+ ( -ln(Ag+)). Such a sharp increase in the concentration of silver ions can be detected in different ways and accordingly there are three different methods of detecting the end point of the titration.
According to end point detection method, three main procedures are widely used depending on the type of application. These are:
Mohr's Method
Volhard's Method
Fajan's Method
Table 2. Comparison of silver titration methods

1. The Mohr method uses chromate ions as an indicator in the titration of chloride ions with a silver nitrate secondary standard solution (normality: 0.0141). This corresponds to 1 mL of 0.0141N AgNO3 equals to 1 mg chloride in solution. The silver nitrate solution is standardized against standard chloride solution, prepared from sodium chloride (NaCl). During the titration, chloride ion is precipitated as white silver chloride:
Ag+ + Cl- = AgCl (Ksp=3×10-10)

After all the chloride has been precipitated as white silver chloride, the first excess of titrant results in the formation of a reddish-brown silver chromate precipitate, which signals the end point. This stage is taken as evidence that all chloride ions have been consumed and only excess silver ions have reacted with chromate ions. The reaction is:
2Ag++CrO42-=Ag2CrO4 (s) (Ksp=5×10-12)

The solution needs to be kept around a neutral pH neutral: silver hydroxide forms at high pH, and chromate will form H2CrO4 at low pH. This will reduce the concentration of chromate ions, and delay the formation of the precipitate. In addition, since carbonates and phosphates both precipitate with silver, they must be eliminated from the reaction to prevent inaccurate results. The Mohr method is a relatively simple and accurate method for chloride ion determination. As such, it has many applications where the concentration of chloride in water or in food must be determined.
By knowing the stoichiometry and moles consumed at the end point, the amount of chloride in an unknown sample can be determined.

2. The Volhard's method was first described by the Jacob Volhard, a German Chemist, in 1874.
This is an indirect titration procedure, where an excess amount of
standard Ag+ is added to the chloride solution containing Fe3+ as an indicator. The excess Ag+ is then titrated with standard SCN- solution until a red color is obtained which results from the reaction:
Ag+ + Cl- = AgCl + Ag+excess
white ppt
Ag+excess + KSCN -> AgSCN + K+
white ppt
At end point: Fe3+ + SCN- Fe(SCN)2+
red brown ppt

Here, initially thiocyanate react with silver ions and forms precipitate at end point, excess of thiocyanate (SCN-) react with Fe(III) and forms reddish brown complex which indicate the end point of reaction.
The indicator system is very sensitive and usually good results are obtained. The medium should be acidic to avoid the formation of Fe(OH)3.
However, the use of acidic medium together with added SCN- titrant increase the solubility of the precipitate leading to significant errors. This problem had been overcome by two main procedures.
Thiocyante is standardised against a standard silver solution, with the silver solution being in the titration flask and the thiocyanate in the burette.
This Volhard method is used to determine the concentration of Ag+ ions or concentration of halide ions (i.e. Cl-, Br-, I-) indirectly i.e. by back titration.

3. Fajan's Method (indicator adsorption method). The precipitation titration in which silver ions is titrated with halide or thiocyanate ions in presence of adsorption indicator is called fajan's method.
Adsorption indicators function in an entirely different manner than the chemical indicators and they can be used in many precipitation titrations.
Since the adsorption of indicator takes place at end point the method is also called indicator adsorption method.
Table 3. Common adsorption indicators

The indicator, which is a dye, exists in solution as the ionized form, usually an anion.
The method is generally used for the quantitative analysis of halide ions or thiocyanate ions.
Silver chloride forms colloidal particles. Before the equivalence point, the surface of the precipitant particles will be negatively charged due to the adsorption of excess Cl to the surface of the particles. When the equivalence point is reached, there is no longer an excess of analyte Cl , and the surface of the colloidal particles are largely neutral. After the equivalence point, there will be an excess of titrant Ag+, some of these will adsorb to the solid AgCl particles, which will now be surrounded by a diffuse negative counterion layer.
Adsorption indicators are dyes, such as dichlorofluorescein, that usually exist as anions in the titration solution. The doubly charged dichlorofluoroscein anion is attracted into the counterion layer immediately following the equivalence point, when the surface charge of the particles changes from negative to positive. For reasons that are not fully understood, the closer proximity of the dye to the particles changes the color of the molecule, providing a visual indication of the titration endpoint. In the case of dichlorofluorescein, the indicator changes to a pinkish color.
Fluorescein and its derivatives are adsorbed to the surface of colloidal AgCl. After all chloride is used, the first drop of Ag+ will react with fluorescein (FI-) forming a reddish color.
Ag+ + FI- AgF
Among these methods, the Volhard Method is widely used because we can detect the end point of precepitation titration very well.

1ST EXPERIMENT. Standardization of silver nitrate solution with sodium chloride

Apparatus: Burette, conical flask, pipette, measuring cylinder.
Reagents: 5% K2CrO4 (indicator) solution, 0.1M silver nitrate titrant and solid NaCl.

Aqueous silver nitrate is photosensitive and should not be exposed to light any more than is necessary during this procedure. It should be stored in darkened storage bottles, and be kept in your drawer except when being used.
Silver nitrate is an important precipitating reagent which can also be used for thedetermination of the halogens, halogenlike anions, mercaptans, fatty acids, and several divalentinorganic anions.
Silver nitrate solutions of known concentration can be prepared from known mass of dried AgNO3. However, if we don't have access to the high purity reagent, or if we have a solution of unknown concentration, we can easily standardize it against sodium chloride.
Reaction taking place during titration is
AgNO3 + NaCl AgCl + NaNO3

Preparation of 5% K2CrO4 (indicator): 1.0 g of K2CrO4 was dissolved in 20 mL of distilled water.
Preparation of AgNO3 solution: 9.0 g of AgNO3 was weighed out, transferred to a 500 mL volumetric flask and made up to volume with distilled water. The resulting solution was approximately 0.1 M. This solution was standardized against NaCl.
Preparation of standard NaCl solution: Reagent-grade NaCl was dried overnight and cooled to room temperature. Accurately weigh about 0.6g of sodium chloride in a clean dry weighing bottle, and transfer the same to a clean volumetric flask of 100 cm3 capacity through a glass funnel. Add about 20 mL of distilled water and swirl the contents of the flask until all the sodium chloride is dissolved. Make the volume upto the mark by adding more distilled water. In order to adjust the pH of the solutions, small quantities of NaHCO3 we are added until effervescence ceased.

PROCEDURE:
1. Take with pipette 25 ml NaCl standard solution in a conical flask. Measure sample pH.
2. Then add 2.0 ml 5% K2CrO4 indicator solution.
3. Fill up the burette with AgNO3 solution up to the zero mark (bottom menisque).
4. Titrate the sodium chloride solution silver nitrate titrant solution. Although the silver chloride that forms is a white precipitate, the chromate indicator initially gives the cloudy solution a faint lemon-yellow colour. The endpoint of the titration is identified as the first appearance of a red-brown colour of silver chromate Ag2CrO4 and note down volume of titrant used. Also measure sample pH.
5. All titration processes are done in three trials.
6. Calculate concentration and titer of titrant AgNO3.

Results of titration:
Volume of standard solution for titration: VNaCl= 10 mL
Normality of standard solution for titration: NNaCl = 0,1N
Equivalent weight of sodium chloride: EqNaCl=
Equivalent weight of potassium permanganate: EqAgNO3=
Volume of titrant consumed for the 1st titration: V1AgNO3 =
Volume of titrant consumed for the 2nd titration: V2AgNO3 =
Volume of titrant consumed for the 3rd titration: V3AgNO3 =
Average amount of titrant consumed for titration: VAgNO3*=V1+V2+V33=

Calculation:
The result is expressed as mol-eq/l concentration and based on the following formula:

VNaClVAgNO3*=NAgNO3NNaCl => NAgNO3=VNaCl NNaClVAgNO3*

TAgNO3=NAgNO3 EqAgNO31000

TAgNO3/NaCl=NAgNO3 EqNaCl1000

The calibration result can be accepted if 3 determinations give a result with a relative standard deviation of less than 0.5%

2nd EXPERIMENT. Use Mohr's method to determine the concentration of chloride ion in solid sample.

Mohr method of determination of chlorides by titration with silver nitrate is one of the oldest titration methods still in use - it was researched and published by Karl Friedrich Mohr in 1856.
The Mohr Method uses silver nitrate for titration (normality: 0.0141) (method applicability: 0.15 to 10 mg/L chloride ions). This corresponds to 1 mL of 0.0141 equals to 1 mg chloride in solution. The idea behind is very simple - chlorides are titrated with the silver nitrate solution in the presence of chromate anions. End point is signalled by the appearance of the red silver chromate.

Initial point white ppt of AgCl End point of Titration

Intense yellow color of chromate may make detection of first signs of formation of red silver chromate precipitation difficult. As some excess of silver must be added before precipitate starts to form, if concentration of titrant is below 0.1M, we may expect significant positive error. To correct for this error we can determine a blank, titrating a solution of the indicator potassium chromate with standard silver nitrate solution. To make result more realistic we can add small amount of chloride free calcium carbonate to the solution to imitate the white silver precipitate.
Solution during titration should be close to neutral. The Mohr method for determination of chloride in water is a pH sophisticated method. It must be perform between the pH levels 6.5 – 9.0. It is better to carry out between the pH ranges 7 – 8. At upper pH level, the silver ions react with hydroxide ions and precipitated as silver hydroxide. In contrast, at lower pH level, potassium chromate may be converted into potassium dichromate (K2Cr2O7) and mask the end point. Consequently, accurate result cannot be obtained. If the water sample is acidic, then gravimetric method or volhard's method is appropriate.
Ag+(aq) + OH–(aq) Ag(OH)(s)
CrO42-(aq) Cr2O72-(aq)
Both processes interfere with the determination accuracy.
Exactly the same approach can be used for determination of bromides. Other halides and pseudohalides, like I- and SCN-, behave very similarly in the solution, but their precipitate tends to adsorb chromate anions making end point detection difficult.

Most foods have sodium from dissolved salts, either naturally present or added in cooking or processing. Table salt known as sodium chloride (NaCl) is the most common source of sodium. It is made up of 40% sodium and 60% chloride and often used in processed and packaged foods as flavour enhancer or preservative. Other sources of sodium added in foods are monosodium glutamate (MSG), sodium nitrite, sodium saccharin, baking soda (sodium bicarbonate), and sodium benzoate.
The sodium content of food has implications on our health. Sodium is an essential mineral required in small amount by the body to control blood pressure and help the nerves and muscles to function properly. However, high sodium intake can cause health problems such as high blood pressure and cardiovascular diseases, which include heart, stroke, and blood vessel disease. Thus, knowing the sodium content in food and controlling the intake are of utmost important to keep diseases at bay.
The American Heart Association1 recommends consumption of less than 1,500 mg per day sodium for most American adults, which is the level with the greatest effect on blood pressure. This level does not apply to people who lose large amounts of sodium in sweat, such as competitive athletes, workers exposed to extreme heat stress, or to those directed otherwise by their healthcare provider.

PROCEDURE
1. The chloride samples (NaCl) have been dried and stored in a desiccator. Take a sample and record its number on your data sheet.
2. Zero (tare) the analytical balance with a clean dry, capped weighing bottle. Add 1.1 to 1.3g of your chloride sample (NaCl) to the bottle; replace the cap and weigh as accurately as possible. Record on your data sheet.
3. Rinse a 250 mL volumetric flask with distilled water and quantitatively transfer the chloride sample to the flask. Dissolve the sample and dilute to the mark. Mix thoroughly.
4. Into each of three 200 - 250 mL erlenmeyer flasks carefully pipet 20.00 mL of your chloride sample. (Remember to rinse the pipet with the sample solution before the first transfer.) Use a graduated cylinder to add about 20 mL of distilled water to each flask.
5. Add 2.0 mL 5% K2CrO4 indicator solution to the first sample flask.
6. Rinse and fill a 50 mL burette with standartized AgNO3 titrant solution. Tap to remove bubbles and make certain that the burette tip is filled with solution. Adjust the solution level between 0 and 2 mL; read the burette to the nearest 0.01 mL. (Do not try to set 0.00 mL!) Record the initial reading.
7. Titrate first flask to the first appearance of a persistent brown color. Read the burette to the nearest 0.01 mL and record the value on your data sheet. Repeat the procedure with the other samples. You should have three endpoint volumes within a range of 0.1 - 0.2 mL. If this is not the case, prepare a few more sample solutions and repeat the titrations.
8. Empty the contents of the titration flasks as well as any excess AgNO3 solution into a waste container. (There are large bottles in the fume hoods for this purpose.)
9. Rinse the burette, pipet and volumetric flask with several portions of distilled water and put them away. Use a brush and soapy water to wash the Erlenmeyer flasks.

Weight of NaCl salt sample: msample =
Stock volume of NaCl salt solution: Vtotal NaCl = 250 mL
Equivalent weight of Cl- ion in Mohr salt: EqCl-=
Volume of NaCl salt solution for titration: VNaCl = 10 mL
Normality of standardized AgNO3 solution for titration: NAgNO3 =

Calculation:
The result is expressed as mol-eq/l concentration and based on the following formula:

VNaClVAgNO3*=NAgNO3NCl- => NCl-=VAgNO3 NAgNO3VNaCl

TCl-=NCl- EqCl-1000

mCl-=TCl- Vtotal NaCl

WCl-,%=m(Cl-)msample 100%

The calibration result can be accepted if 3 determinations give a result with a relative standard deviation of less than 0.5%

3rd EXPERIMENT. Use Mohr's method to determine the concentration of chloride ion in liquid sample.

Apparatus: Conical flask, burette with stand, pipette, measuring cylinder, volumetric flask, beakers, wash bottle.
Reagents: Deionized water, water sample, standard 0.1N silver nitrate solution, pH test paper, sodium hydroxide solution, nitric acid, Indicator potassium chromate solution.

Chloride in the form of chloride (Cl) ion is one of the major inorganic anions in water and wastewater. The chloride concentration is higher in wastewater than in raw water because sodium chloride is a common article of diet and passes unchanged through the digestive system (Average estimate of excretion: 6 g of chlorides/person/day; additional chloride burden due to human consumption on wastewater: 15 mg/L). Along the sea coast chloride may be present in high concentration because of leakage of salt water into the sewage system. It also may be increased by industrial process. In potable water, the salty taste produced by chloride concentration is variable and depends on the chemical composition of water. Some waters containing 250 mg/L Cl may have a detectable salty taste if sodium cation is present. On the other hand, the typical salty taste may be absent in waters containing as much as 1000 mg/L when the predominant cations are calcium and magnesium. In addition, a high chloride contents may harm metallic pipes and structures as well as growing plants.
The measured chloride ions can be used to know salinity of different water sources. For brackish water (or sea water or industrial brine solution), it is an important parameter and indicates the extent of desalting of apparatus required. It also interferes with COD determination and thus it requires a correction to be made on the basis of amount present or else a complexing agent, such as HgSO4 can be added. Further, chloride ions are used as tracer ions in column studies to model fate of different contaminants in soil and liquid media.

PROCEDURES:
Measure the pH of the water sample. Adjust the pH with nitric acid or sodium hydroxide, if needed.
Take a 25 ml collected water sample into a conical flask.
Add 2-3 drops potassium chromate (K2CrO4) indicator. The color of the water sample is turn into light yellow.
Add standard (normality: 0.0141N) silver nitrate solution from the burette and shake well. Titrate until the light yellow color changes to permanent brownish-red color (bricks-red color) precipitate with white color precipitate.
Note the volume of silver nitrate added.
Repeat the titration for concordant values.

Equivalent weight of Cl- ion: EqCl-=
Volume of liquid sample solution for titration: Vsample = 10 mL
Volume of titrant consumed for titration: VAgNO3 =

Calculation:
Chloride ion concentration (mg/l): C(Cl-)=(VAgNO3 NAgNO3 35,45) 1000Vsample

Determination of Sodium chloride (salt content) in brine:
Direct titration of NaCl in brine with standardized silver nitrate solution based on the Mohr method is adequate for routine analysis.

Procedure:
Take 5 to 10 mg liquid portion from the drained weight determination. If it is acidic, neutralize it with standard Sodium hydroxide using phenolphthalein as indicator. Add 1 mL of 5% aqueous potassium chromate solution and titrate with 0.1 N AgNO3 solution to produce red-brown end point.
Equivalent weight of NaCl: EqNaCl=
Weight of liquid saple solution for titration: msample = ______mg
Volume of titrant consumed for titration: VAgNO3 =

Calculation: Equivalent weight of Cl- ion: EqCl-=
Volume of liquid saple solution for titration: Vsample = 10 mL

Calculation:
Sodium chloride concentration: WNaCl, %=VAgNO3 NAgNO3 58,5msample 100%

CONTROL OF MASTERING THE TOPIC

Typical calculation tasks
A mixture containing only KCl and NaBr is analyzed by the Mohr method. A 0.31720g sample is dissolved in 50 mL of water and titrated to the Ag2CrO4 end point, requiring 36.85 mL of 0.1120 M AgNO3. A blank titration requires 0.71 mL of titrant to reach the same end point. Report the %w/w KCl in the sample.
The %w/w I– ions in a 0.6712-g sample was determined by a Volhard titration. After adding 50.00 mL of 0.05619 M AgNO3 and allowing the precipitate to form, the remaining silver was back titrated with 0.05322 M KSCN, requiring 35.14 mL to reach the end point. Report the %w/w I– in the sample.
In the Volhard titration of 25 mL of 0.05 M of AgNO3 solution with 0.02 M KSCN solution, calculate the molar concentration of Ag+ in the conical flask solution after the following additions of titrant KSCN solution: (1) 30 mL (2) at equivalent point (3) 100 mL. (Ksp (AgSCN) = 1.010-12)
The As in a 9.13-g sample of pesticide was converted to AsO43- and
precipitated as Ag3AsO4 with 50.00 mL of 0.02015 M AgNO3. The excess
Ag+ was then titrated with 4.75 mL of 0.04321 M KSCN. Calculate the % of As2O3 in the sample.
400 mg of butter was heated and some water was added. After shaking and filtration, 10 ml 0.2 M AgNO3 solution, some HNO3, drops of Fe3+ solution and some nitrobenzene were added to the filtrate. The excess Ag+ in the aqueous layer was titrated with 0.1 M NH4SCN standard solution. If the volume of NH4SCN at the equivalent point was 15 mL. Calculate the percentage of NaCl in the butter sample.
Calculate the concentration of salt in a soy sauce 5.0g of sample, which is diluted accurately in 250 mL volumetric flask and its 10.25 mL aliquots of the diluted sample are mixed with 25 mL aliquots of the 0.1025N silver nitrate standard solution, and back titrated with an average of 14.6 mL of 0.1002N ammonium thiocyanate standard solution.
A 0.32 g sample containing KCl is dissolved in 50 mL of water and titrated to the Ag2CrO4 end point, requiring 16.9 mL of 0.1 M AgNO3. A blank titration requires 0.7 mL of titrant to reach the same end point. Report the %w/w KCl in the sample.
The sulphide contents of 100 mL of a water sample was titrated with a standard solution of 0.01 M AgNO3 according to the following reaction equation:
2Ag+ + S2- Ag2S
If the volume of AgNO3 solution at the equivalent point was 8.5 mL. Calculate the concentration of H2S in the water sample.
1.354 g sample of sodium nitrate contaminated with NaCl was dissolved in small amount of water and filled to the mark in the 100 mL volumetric flask. To titrate chlorides in 10.00 mL sample 35.70 mL of 0.01021 M AgNO3 was used. What is the percent purity of the sample?
A 2.2380g sample of a mixture of NaCl and NaI was dissolved in 200.0 mL and 25mL aliquots were titrated to endpoint, using fluorescein indicator with an average of 24.8 mL. Further 25 mL aliquots were titrated, using diiododimethylfluorescein indicator with an average of 10.3 mL. Calculate the %w/w of each salt in the mixture.
50.0 mL aliquots of a bromide solution were titrated to endpoint with an average of 12.3 mL of the silver solution which has been diluted accurately by a factor of 10. Calculate the concentration of bromide in mg/L.
Calculate the percentage of silver in an ingot, if a 0.9023g sample was dissolved in 100 mL and its 25 mL aliquots were titrated to endpoint with an average of 17.9 mL of thiocyanate titrant.
A waste water sample from a gold processing plant is analysed for its cyanide content by titration with silver nitrate. 50 mL aliquots of the sample are titrated to endpoint with an average of 13.2 mL of silver nitrate, which has been diluted five times. Does this sample contravene effluent regulations which limit cyanide waste to 100 mg/L?
The organic matter in a 3.776-g sample of a mercuric ointment is decomposed with HNO3. After dilution, the Hg2+ is titrated with 21.30 mL of a 0.1144 M soln of NH4SCN. Calculate the percent Hg (200.59 g/mol) in the ointment.
What is the solubility of barium sulfate in pure water at 25oC? (Ksp for barium sulfate is 1.1 x 10-10)
Calculate the solubility product constant for pure PbSO4 in water. The solubility of PbSO4 is 1.25 x 10-4 mol/L.
The solubility of CaF2 is 2.1 x 10 -4 mol/L. Find the Ksp of CaF2.

Questions for test self-check
1. What is the role of chromate ions in chloride determination?
2. As potassium chromate is an oxidizing agent, what would happen to chloride determination if the sample were consists of organic matter (say 100 mg/L glucose) as well.
3. Why pH range is important in chloride determination?
4. Would the analytical results by the Mohr method for chlorides be higher, lower or the same as the true color value if any excess of indicator were accidentally added to the sample? Why?
5. What is the solubility constant expression for: a) Zn3(PO4)2; b) Ag3PO4, c) Mg(OH)2?
6. In most titrations it does not really matter which solution is in the burette and which is in the flask. However with Mohr's method, it is critical that the silver is in the burette and not in the flask. Explain why.
7. Why is a blank used in Mohr's method?
8. What indicator would be suitable to analyse the concentration of iodide in the presence of chloride?
9. How would you analyse the salt content of a sauce which contains a significant level of ethanoic acid?
References:
Harris, Daniel Charles (2003). Quantitative chemical analysis (6th ed.). San Francisco: W.H. Freeman. pp. 142–143.
James R. Fromm (1997). Precipitation Titrations
Yoder, Lester (1919). "Adaptation of the Mohr Volumetric Method to General Determinations of Chlorine". Industrial & Engineering Chemistry 11(8): pp.75.
Skoog D. A.; West D. M.; Holler F. J. Fundamentals of Analytical Chemistry, 7th Edition, Thomson Learning, Inc, USA, 1996.
Sheen R.T. and Kahler H. L. Effects of Ions on Mohr Method for Chloride Determination, Ind. Eng. Chem. Anal. Ed.; 1938; 10(11); 628-629.
Kraemer E. O. and Stamm A. J. Mohr's Method for the Determination of Silver and Halogens in other than Neutral Solutions, J. Am. Chem. Soc.; 1924; 46(12); pp. 2707- 2709.
Nielsen, Suzanne. Sodium Determination Using Ion Selective Electrodes, Mohr Titration, and Test Strips. Chapter 10. Food Analysis Laboratory Manual. 2nd Edition. USA: Springer. 2015

LABORATORY WORK PRECIPITATION TITRATION. MOHRS METHODS

Recommend Documents