Types of titration. Analytical chemistry Titration in brief

24.11.202227.05.2017

Send your good work in the knowledge base is simple. Use the form below

Students, graduate students, young scientists who use the knowledge base in their studies and work will be very grateful to you.

Posted on http://www.allbest.ru/

Plan

1. Essence of precipitation titration

2. Argentometric titration

3. Thiocyanatometric titration

4. Application of precipitation titration

4.1 Preparation of standardized silver nitrate solution

4.2 Preparation of a standardized ammonium thiocyanate solution

4.3 Determination of the chlorine content of a sample according to Volhard

4.4 Determination of the content of sodium trichloroacetate in a technical product

1. The essence of precipitationtitration

The method combines titrimetric determinations based on the reactions of precipitate formation of poorly soluble compounds. For these purposes, only certain reactions that satisfy certain conditions are suitable. The reaction must proceed strictly according to the equation and without side processes. The resulting precipitate should be practically insoluble and precipitate fairly quickly, without the formation of supersaturated solutions. In addition, it is necessary to be able to determine the end point of the titration using an indicator. Finally, the phenomena of adsorption (co-precipitation) must be expressed so weakly during titration that the result of the determination is not distorted.

The names of the individual precipitation methods are derived from the names of the solutions used. The method using a solution of silver nitrate is called argentometry. This method determines the content of C1~ and Br~ ions in neutral or slightly alkaline media. Thiocyanatometry is based on the use of a solution of ammonium thiocyanate NH 4 SCN (or potassium KSCN) and serves to determine traces of C1- and Br ~, but already in strongly alkaline and acidic solutions. It is also used to determine the silver content in ores or alloys.

The expensive argentometric method for determining halogens is gradually being replaced by the mercurometric method. In the latter, a solution of mercury nitrate (I) Hg 2 (NO 3) 2 is used.

Let us consider in more detail argentometric and thiocyanatometric titration.

2. Argentometric titration

The method is based on the reaction of precipitation of C1~ and Br~ ions by silver cations with the formation of sparingly soluble halides:

Cl-+Ag+=AgClb Br^- + Ag+= AgBr

In this case, a solution of silver nitrate is used. If the substance is analyzed for silver content, then a solution of sodium (or potassium) chloride is used. titration solution drug

Titration curves are of great importance for understanding the method of argentometry. As an example, consider the case of titration of 10.00 ml of 0.1 N. sodium chloride solution 0.1 N. a solution of silver nitrite (without taking into account the change in the volume of the solution).

Before the start of titration, the concentration of chloride ions in the solution is equal to the total concentration of sodium chloride, i.e. 0.1 mol / l or \u003d -lg lO-i \u003d 1.

When 9.00 ml of silver nitrate solution is added to the sodium chloride solution being titrated and 90% of the chloride ions are precipitated, their concentration in the solution will decrease by a factor of 10 and become equal to N0 ~ 2 mol/l, and pC1 will be equal to 2. Since value nPAgci= IQ- 10 , the concentration of silver ions in this case will be:

10th / [C1-] \u003d 10-10 / 10-2 \u003d 10-8 M ol / l, OR pAg \u003d - lg \u003d - IglO-s \u003d 8.

Similarly, all other points are calculated to plot the titration curve. At the equivalence point pCl=pAg= = 5 (see table).

Table Change in pC\ and pAg during titration of 10.00 ml of 0.1 N. sodium chloride solution 0.1 N. silver nitrate solution

AgNO 3 solution added,


9.99 10.00 (equiv. point) 10.01	yu-4 yu-5 yu-6.	yu- 6 yu- 5 yu-*

The jump interval in argentometric titration depends on the concentration of the solutions and on the value of the solubility product of the precipitate. The smaller the PR value of the compound resulting from the titration, the wider the jump interval on the titration curve and the easier it is to fix the end point of the titration using an indicator.

The most common is the argentometric determination of chlorine by the Mohr method. Its essence consists in the direct titration of a liquid with a solution of silver nitrate with an indicator of potassium chromate until a white precipitate turns brown.

Mohr's method indicator - a solution of K2CrO 4 gives a red precipitate of silver chromate Ag 2 CrO 4 with silver nitrate, but the solubility of the precipitate (0.65-10 ~ 4 E / l) is much greater than the solubility of silver chloride (1.25X _X10 ~ 5 E / l ). Therefore, when titrating with a solution of silver nitrate in the presence of potassium chromate, a red precipitate of silver chromate appears only after adding an excess of Ag + ions, when all chloride ions have already precipitated. In this case, a solution of silver nitrate is always added to the analyzed liquid, and not vice versa.

The possibilities of using argentometry are quite limited. It is used only when titrating neutral or slightly alkaline solutions (pH 7 to 10). In an acidic environment, the silver chromate precipitate dissolves.

In strongly alkaline solutions, silver nitrate decomposes with the release of insoluble oxide Ag 2 O. The method is also unsuitable for the analysis of solutions containing the NH ^ ion, since in this case an ammonia complex is formed with the Ag + cation + - The analyzed solution should not contain Ba 2 +, Sr 2+ , Pb 2+ , Bi 2+ and other ions that precipitate with potassium chromate.Nevertheless, argentometry is convenient in the analysis of colorless solutions containing C1~ and Br_ ions.

3. Thiocyanatometric titration

Thiocyanatometric titration is based on the precipitation of Ag+ (or Hgl+) ions with thiocyanates:

Ag+ + SCN- = AgSCN|

The determination requires a solution of NH 4 SCN (or KSCN). Determine Ag+ or Hgi + by direct titration with a solution of thiocyanate.

Thiocyanatometric determination of halogens is carried out according to the so-called Volhard method. Its essence can be expressed in diagrams:

CI- + Ag+ (excess) -* AgCI + Ag+ (residue), Ag+ (residue) + SCN~-> AgSCN

In other words, an excess of a titrated solution of silver nitrate is added to a liquid containing C1~. The AgNO 3 residue is then back titrated with a thiocyanate solution and the result is calculated.

The indicator of the Volhard method is a saturated solution of NH 4 Fe (SO 4) 2 - 12H 2 O. As long as there are Ag + ions in the titrated liquid, the added SCN ~ anions bind to precipitate AgSCN, but do not interact with Fe 3 + ions. However, after the equivalence point, the slightest excess of NH 4 SCN (or KSCN) causes the formation of blood-red ions 2 + and +. Thanks to this, it is possible to determine the equivalent point.

Thiocyanatometric definitions are used more often than argentometric ones. The presence of acids does not interfere with the Volhard titration and even contributes to obtaining more accurate results, since the acidic medium inhibits the hydrolysis of the Fe** salt. The method makes it possible to determine the C1~ ion not only in alkalis, but also in acids. The determination does not interfere with the presence of Ba 2 +, Pb 2 +, Bi 3 + and some other ions. However, if the analyzed solution contains oxidizing agents or mercury salts, then the application of the Volhard method becomes impossible: oxidizing agents destroy the SCN- ion, and the mercury cation precipitates it.

The alkaline test solution is neutralized before titration with nitric acid, otherwise the Fe 3 + ions that are part of the indicator will precipitate iron (III) hydroxide.

4. Application of precipitation titration

4.1 Preparation of a standardized solution of silver nitrate

The primary standards for standardizing silver nitrate solution are sodium or potassium chlorides. Prepare a standard solution of sodium chloride and approximately 0.02 N. silver nitrate solution, standardize the second solution according to the first.

Preparation of standard sodium chloride solution. A solution of sodium chloride (or potassium chloride) is prepared from chemically pure salt. The equivalent mass of sodium chloride is equal to its molar mass (58.45 g/mol). Theoretically, for the preparation of 0.1 l 0.02 N. solution requires 58.45-0.02-0.1 \u003d 0.1169 g of NaCl.

Take a sample of approximately 0.12 g of sodium chloride on an analytical balance, transfer it to a 100 ml volumetric flask, dissolve, bring the volume to the mark with water, mix well. Calculate the titer and normal concentration of the stock sodium chloride solution.

Preparation of 100 ml of approximately 0.02 N. silver nitrate solution. Silver nitrate is a scarce reagent, and usually its solutions have a concentration not higher than 0.05 N. For this work, 0.02 n is quite suitable. solution.

In argentometric titration, the equivalent mass of AgN0 3 is equal to the molar mass, i.e., 169.9 g / mol. Therefore, 0.1 l 0.02 n. the solution should contain 169.9-0.02-0.1 \u003d 0.3398 g AgNO 3. However, it does not make sense to take exactly such a sample, since commercial silver nitrate always contains impurities. Weigh on technochemical scales approximately 0.34 - 0.35 g of silver nitrate; weigh the solution in a volumetric flask with a capacity of 100 ml, a solution in a small amount of water and bring the volume with water; store the solution in the flask, wrapping it in black paper and pour it into a dark glass bottle. Standardization of the sulfur nitrate solution by sodium chloride. silver and prepare it for titration. Rinse the pipette with sodium chloride solution and transfer 10.00 ml of the solution into a conical flask. Add 2 drops of saturated potassium chromate solution and carefully titrate with silver nitrate solution drop by drop while stirring. Ensure that the mixture turns from yellow to reddish with one excess drop of silver nitrate. After repeating the titration 2-3 times, take the average of the convergent readings and calculate the normal concentration of the silver nitrate solution.

Let us assume that for titration 10.00 ml of 0.02097 N. sodium chloride solution went on average 10.26 ml of silver nitrate solution. Then

A^ AgNOj . 10.26 = 0.02097. 10.00, AT AgNOs = 0.02097-10.00/10.26 = 0.02043

If it is supposed to determine the content of C1~ in the sample, then, in addition, the titer of the solution of silver nitrate by chlorine is calculated: T, - \u003d 35.46-0. ml of silver nitrate solution corresponds to 0.0007244 g of titrated chlorine.

4.2 Preparation of standardized ammonium thiocyanate solutionI

A solution of NH 4 SCN or KSCN with a precisely known titer cannot be prepared by dissolving a sample, since these salts are very hygroscopic. Therefore, prepare a solution with an approximate normal concentration and set it to a standardized solution of silver nitrate. The indicator is a saturated solution of NH 4 Fe (SO 4) 2 - 12H 2 O. To prevent the hydrolysis of the Fe salt, 6 N is added to the indicator itself and to the analyzed solution before titration. nitric acid.

Preparation of 100 ml of approximately 0.05 N. ammonium thiocyanate solution. The equivalent mass of NH4SCN is equal to its molar mass, i.e. 76.12 g/mol. Therefore, 0.1 l 0.05 n. the solution should contain 76.12.0.05-0.1=0.3806 g of NH 4 SCN.

Take a sample of about 0.3-0.4 g on an analytical balance, transfer it to a 100 ml flask, dissolve, dilute the volume of the solution with water to the mark and mix.

Standardization of ammonium thiocyanate solution by silver nitrate. Prepare a burette for titration with the NH 4 SCN solution. Rinse the pipette with silver nitrate solution and measure 10.00 ml of it into a conical flask. Add 1 ml of NH 4 Fe(SO 4) 2 solution (indicator) and 3 ml. 6 n. nitric acid. Slowly, with continuous agitation, pour the NH 4 SCN solution from the burette. Stop the titration when a brown-pink 2+ color appears, which does not disappear with vigorous shaking.

Repeat the titration 2-3 times, take the average from the convergent readings and calculate the normal concentration of NH 4 SCN.

Let us assume that for titration 10.00 ml of 0.02043 N. silver nitrate solution went on average 4.10 ml of NH 4 SCN solution.

4.3 Definitioncontentchlorine in the sample according to Folgard

Volhard halogens are determined by back titration of the silver nitrate residue with a solution of NH 4 SCN. However, accurate titration is possible here only on the condition that measures are taken to prevent (or slow down) the reaction between silver chloride and an excess of iron thiocyanate:

3AgCI + Fe(SCN) 3 = SAgSCNJ + FeCl 3

in which the color that appears at first gradually disappears. It is best to filter the AgCl precipitate before titrating the excess silver nitrate with NH 4 SCN solution. But sometimes, instead, some organic liquid is added to the solution, it is not mixed with water and, as it were, isolating the ApCl precipitate from excess nitrate.

Definition method. Take a test tube with a solution of the analyte containing sodium chloride. A weighed portion of the substance is dissolved in a volumetric flask with a capacity of 100 ml and the volume of the solution is brought to the mark with water (the concentration of chloride in the solution should be no more than 0.05 N).

Pipette 10.00 ml of the analyzed solution into a conical flask, add 3 ml of 6N. nitric acid and add a known excess of AgNO 3 solution from the burette, for example 18.00 ml. Then filter the precipitate of silver chloride. Titrate the silver nitrate residue with NH 4 SCN as described in the previous paragraph. After repeating the definition 2-3 times, take the average. If the precipitate of silver chloride is filtered, then it should be washed and the washings added to the filtrate.

Let us assume that the sample weight was 0.2254 g. To 10.00 ml of the analyzed solution was added 18.00 ml of 0.02043 N. silver nitrate solution. For the titration of its excess, 5.78 ml * 0.04982 n. NH 4 SCN solution.

First of all, we calculate what volume 0.02043 n. silver nitrate solution corresponds to 5.78 ml of 0.04982 N spent on titration. NH 4 SCN solution:

consequently, 18.00 - 14.09 = 3.91 ml of 0.2043 n went to the precipitation of the C1 ~ ion. silver nitrate solution. From here it is easy to find the normal concentration of sodium chloride solution.

Since the equivalent mass of chlorine is 35.46 g/mol*, the total mass of chlorine in the sample is:

772 \u003d 0.007988-35.46-0.1 \u003d 0.02832 g.

0.2254 g C1 - 100%

x \u003d 0.02832-100 / 0.2254 \u003d 12.56%.:

0.02832 > C1 -- x%

According to the Folgard method, the content of Br~ and I- ions is also determined. At the same time, it is not required to filter out precipitates of silver bromide or iodide. But it must be taken into account that the Fe 3 + ion oxidizes iodides to free iodine. Therefore, the indicator is added after precipitation of all ions of I-silver nitrate.

4.4 Determination of trichl contentOsodium acetate | in a technical preparation (for chlorine)

Technical sodium trichloroacetate (TXA) is a herbicide for controlling grass weeds. It is a white or light brown crystalline substance, highly soluble in water. According to Folgard, the mass fraction of organochloride compounds is first determined, and then after the destruction of chlorine. By difference, find the mass fraction (%) of sodium chlorine trichloroacetate.

Determination of the mass fraction (%) of chlorine inorganic compounds. Accurately weighed 2–2.5 g of the drug is placed in a volumetric flask with a capacity of 250 ml, dissolve, dilute the solution with water to the mark, mix. Pipette 10 ml of the solution into a conical flask and add 5-10 ml of concentrated nitric acid.

Add from the burette 5 or 10 ml of 0.05 N. silver nitrate solution and its excess, titrate with 0.05 N. NH 4 SCN solution in the presence of NH 4 Fe(SO 4) 2 (indicator).

Calculate the mass fraction (%) of chlorine (x) of inorganic compounds using the formula

(V - l / i) 0.001773-250x100

where V is the volume exactly 0.05 n. AgNO 3 solution taken for analysis; Vi -- the volume is exactly 0.05 N. NH 4 SCN solution used for titration of excess AgNO 3 ; t is a sample of sodium trichloroacetate; 0.001773 is the mass of chlorine corresponding to 1 ml of 0.05 N. AgNO solution. Determination of the mass fraction (%) of total chlorine. Take 10 ml of the previously prepared solution into a conical flask, add 10 ml of a solution with a mass fraction of NaOH 30% and 50 ml of water. Connect the flask to a reflux bead condenser and boil the contents for 2 hours. Let the liquid cool, rinse the condenser with water, collecting the wash water in the same flask. Add 20 ml of dilute (1:1) nitric acid to the solution and pour 30 ml of 0.05 N. silver nitrate solution. Titrate excess silver nitrate with 0.05 N. NH 4 SCN solution in the presence of NH 4 Fe(SO 4) 2. Calculate the mass fraction (%) of total chlorine (xi) using the above formula. Find the mass fraction (%) of sodium trichloroacetate in the preparation (х^) using the formula

x2 \u003d (x1 - x) (185.5 / 106.5),

where 185.5 is the molar mass of sodium trichloroacetate; 106.5 is the mass of chlorine contained in the molar mass of sodium trichloroacetate.

Hosted on Allbest.ru

...

Classification of methods of titrimetric analysis.

acid-base titration (neutralization): this method determines the amount of acid or alkali in the analyzed solution;

precipitation and complexation (argentometry)

Ag + + Cl - "AgCl $

redox titration (redoximetry):

a) permanganatometry (KMnO 4);

b) iodometry (Y 2);

c) bromatometry (KBrO 3);

d) dichromatometry (K 2 Cr 2 O 7);

e) cerimetry (Ce(SO 4) 2);

f) vanadometry (NH 4 VO 3);

g) titanometry (TiCl 3), etc.

Titrimetric analysis

History and principle of the method

Titrimetric analysis (titrimetry) is the most important of the chemical methods of analysis. It originated in the 18th century, initially as an empirical method for testing the quality of various materials, such as vinegar, soda, bleach solutions. At the turn of the 18th and 19th centuries, burettes and pipettes were invented (F.Decroisille). Of particular importance were the works of J. Gay-Lussac, who introduced the basic terms of this method: titration, titrant and others derived from the word "title". A titer is the mass of a solute (in grams) contained in one milliliter of a solution. At the time of Gay-Lussac, the results of the analysis were calculated using titers. However, the titer as a way of expressing the concentration of a solution turned out to be less convenient than other characteristics (for example, molar concentrations), therefore, in modern chemistry analytics, calculations using titers are quite rare. On the contrary, various terms derived from the word "title" are used very widely.

In the middle of the 19th century, the German chemist K.Mohr summarized all the titrimetric methods created by that time and showed that the same principle underlies any method. Always add a solution with a precisely known concentration of the reagent R (titrant) to a sample solution containing the component X to be determined. This process is called titration. When performing a titration, the analyst monitors the progress of a chemical reaction between X and the added R. Upon reaching the equivalence point (t.eq.), when the number of mole equivalents of the introduced R is exactly equal to the number of mole equivalents of the substance X present in the sample, the titration is stopped and the volume of titrant consumed is measured. The end point of the titration is called the end point of the titration (k.t.t.), it, like t.eq., is expressed in units of volume, usually in milliliters. In the ideal case, Vk.t.t \u003d V t.eq. , but in practice an exact match is not achieved for various reasons, the titration is completed a little earlier or, conversely, a little later than the t.eq. Naturally, the titration should be carried out so that the difference between V t.eq. and V k.t.t. would be as small as possible.

Since the mass or concentration X is calculated from the volume of titrant used to titrate the sample (according to V k.t.t.), in the past titrimetry was called volumetric analysis. This name is often used today, but the term titrimetric analysis more accurate. The fact is that the operation of the gradual addition of a reagent (titration) is characteristic of any technique of this type, and the titrant consumption can be estimated not only by measuring the volume, but also by other methods. Sometimes the added titrant is weighed (measurement of mass on an analytical balance gives a smaller relative error than measurement of volume). Sometimes the time taken for the titrant to be injected (at a constant injection rate) is measured.

Since the end of the 19th century, titrimetric techniques have been used in research, factory, and other laboratories. Using the new method, it was possible to determine milligram and even microgram amounts of a wide variety of substances. The simplicity of the method, low cost, and versatility of the equipment contributed to the widespread use of titrimetry. Especially widely titrimetry began to be used in the 50s of the XX century, after the creation of a new version of this method (complexometry) by the Swiss analyst G. Schwarzenbach. Simultaneously, the widespread use of instrumental methods of control of c.t.t. By the end of the 20th century, the importance of titrimetry declined somewhat due to the competition of more sensitive instrumental methods, but even today titrimetry remains a very important method of analysis. It allows you to quickly, easily and fairly accurately determine the content of most chemical elements, individual organic and inorganic substances, the total content of similar substances, as well as generalized composition indicators (water hardness, milk fat content, acidity of petroleum products).

Titrimetric Analysis Technique

The principle of the method will become clearer after the presentation of the technique of its implementation. So, let an alkali solution of unknown concentration be brought to you, and your task is to establish its exact concentration. For this you will need regent solution, or titrant- a substance that enters into a chemical reaction with alkali, and the concentration of the titrant must be precisely known. Obviously, to determine the concentration of alkali, we use an acid solution as a titrant.

1. We select with a pipette the exact volume of the analyzed solution - it is called aliquot. As a rule, the volume of an aliquot is 10-25 ml.

2. Transfer an aliquot to a titration flask, dilute with water and add an indicator.

3. Fill the burette with the titrant solution and perform titration is the slow, dropwise addition of a titrant to an aliquot of the test solution.

4. We finish the titration at the moment when the indicator changes its color. This moment is called the end point of the titration is k.t.t. K.t.t., as a rule, coincides with the moment when the reaction between the analyte and the titrant is completed, i.e. an exactly equivalent amount of titrant is added to the aliquot - this point is called equivalence point, i.e. Thus, i.e. and k.t.t. - these are two characteristics of the same moment, one is theoretical, the other is experimental, depending on the selected indicator. Therefore, it is necessary to choose the indicator correctly so that the k.t.t. coincided as closely as possible with t.e.

5. Measure the volume of the titrant used for titration and calculate the concentration of the test solution.

Types of titrimetric analysis

Titrimetric methods can be classified according to several independent features: namely: 1) according to the type of reaction between X and R, 2) according to the method of titration and calculation of results, 3) according to the method of monitoring t.eq.

Classification by type of chemical reaction- the most important. Recall that not all chemical reactions can be used for titrations.

First, as in other chemical methods, the component (analyte) to be determined must quantitatively react with the titrant.

Secondly, it is necessary that the equilibrium of the reaction is established as quickly as possible. Reactions in which, after the addition of the next portion of the titrant, the establishment of equilibrium requires at least several minutes, it is difficult or even impossible to use in titrimetry.

Thirdly, the reaction must correspond to a single and previously known stoichiometric equation. If the reaction leads to a mixture of products, the composition of this mixture will change during the titration and depend on the reaction conditions. Fixing the equivalence point will be very difficult, and the result of the analysis will be inaccurate. The combination of these requirements meets the protolysis (neutralization) reaction, many complexation and redox reactions, as well as some precipitation reactions. Accordingly, in titrimetric analysis, the following are distinguished:

Neutralization method,

complexometry,

Redox methods

Deposition methods.

Within each method, its individual variants are distinguished (Table 1). Their names come from the names of the reagents used in each of the options as a titrant (permanganatometry, iodometry, chromatometry, etc.).

Table 1.

Classification of titrimetric methods according to the type of chemical reaction used

Reaction	Method	Reagent (titrant)	method variant	Substances to be determined
Protolysis	Method of neutralization	H Cl, HClO 4 , HNO 3	Acidimetry	Oc innovations
Protolysis	Method of neutralization	KOH, NaOH, etc.	Alkalimetry	acids
Complex formation	Complex-metry	EDTA	Complexometry	Metals and their compounds
Complex formation	Complex-metry		Fluoridometry, cyanidometry	Some metals, organic substances
Redox	Redox meter	KMnO 4 K 2 C r 2 O 7	permanganatometry chromatometry	Restorers
		KJ and Na 2 S 2 O 3	Iodometry	Reducing agents, oxidizing agents, acids
		Ascorbic acid	Ascorbinometry	Oxidizers
precipitation	Sedimetry	AgNO3	Argentometry	Halides
		Hg 2 (NO 3) 2	Mercurymetry	Halides
		KSCN	Rodanometry	Some metals
		Ba(NO3)2	baryometry	sulfates

Classification by titration method. Usually there are three methods: direct, back and substitution titration. direct titration involves the direct addition of the titrant to the sample solution. Sometimes a different order of mixing reagents is used - a sample solution is gradually added to a known amount R, in which they want to determine the concentration X; but this is also a direct titration. In both cases, the analysis results are calculated using the same formulas based on the law of equivalents.

ν X = ν R

where ν X and ν R are the number of moles of X and R equivalents. Calculation formulas based on the ratio, as well as calculation examples, will be given below.

Direct titration is the most convenient and most common variant of titrimetry. It is more accurate than others. After all, random errors mainly arise when measuring the volume of solutions, and in this titration method, the volume is measured only once. However, direct titration is far from always possible. Many reactions between X and R do not go fast enough, and after adding the next portion of the titrant, the solution does not have time to establish equilibrium. Sometimes a direct titration is not possible due to side reactions or lack of a suitable indicator. In such cases, more complex back or substitution titration schemes are used. They include at least two chemical reactions.

Back titration carried out according to a two-stage scheme:

X + R 1 \u003d Y 1

R1 + R2 = Y2

Auxiliary reagent R 1 is introduced in a precisely known amount. The volume and concentration of the solution R 1 is chosen so that R 1 after completion of the reaction with Hostal in excess. The unreacted portion of R 1 is then titrated with R 2 titrant. An example would be the permanganatometric titration of organic substances. It is not possible to titrate many substances with permanganate "directly" because of the slowness of their oxidation and for other reasons. But you can first add a known (excessive) amount of KMnO 4 to the analyzed sample, acidify and heat the resulting solution. This will lead to the complete and rapid completion of the oxidation of organic substances. Then the remaining permanganate is titrated with some active reducing agent, for example, a solution of SnCl 2 or FeSO 4 .

The calculation of the back titration results is carried out based on the obvious relationship:

ν X \u003d ν R 1 - ν R 2

Since the volumes in this case are measured twice (first the volume of the reagent solution R 1 , then the volume of the titrant R 2 ), the random error of the analysis result is slightly higher than in direct titration. The relative error of the analysis increases especially strongly with a small excess of the auxiliary reagent, when ν R 1 ≈ ν R 2 .

Classification according to the method of control of t.eq. Several such methods are known. The simplest is indicatorless titration, the most common is titration with color indicators, and the most accurate and sensitive are instrumental titrimetry options.

Indicatorless titration is based on the use of reactions that are accompanied by a change in the visible properties of the titrated solution. As a rule, one of the reagents (X or R) has a visible color. The course of such a reaction is controlled without special instruments and without the addition of indicator reagents. So, colorless reducing agents are titrated in an acidic medium with a violet solution of an oxidizing agent - potassium permanganate (KMnO 4). Each portion of the added titrant will immediately discolor, turning into Mn 2+ ions under the action of the reducing agent. This will continue up to t.eq. However, the very first "extra" drop of titrant will color the titrated solution pink-violet, the color will not disappear even when the solution is stirred. When a non-vanishing color appears, the titration is stopped and the volume of the spent titrant is measured ( V k.t.t.). The end of the titration can be fixed not only by the appearance of the color of the titrated solution, as in the example considered, but also by the discoloration of the previously colored sample solution, as well as by the appearance of any precipitate, its disappearance or change in appearance. Indicatorless titration is rarely used, since only a few reactions are accompanied by a change in the visible properties of the solution.

Instrumental titration. The course of the reaction between X and R can be monitored not just “by eye” (visually), but also with the help of instruments that measure some physical property of the solution. Variants of instrumental titrimetry are distinguished, depending on which property of the solution is controlled. You can use any property that depends on the qualitative and quantitative composition of the titrated solution. Namely, it is possible to measure the electrical conductivity of the solution (this option is called conductometric titration), the potential of the indicator electrode immersed in the titrated solution ( potentiometric titration), absorption of light by the titrated solution ( photometric titration), etc. You can stop titration when some pre-selected value of the measured property is reached. For example, an acid solution is titrated with alkali until a pH value of 7 is reached. However, more often they do it differently - the selected property of the solution is repeatedly (or even continuously) measured as the titrant is introduced, not only before, but also after the expected temperature. .eq. According to the data obtained, a graphical dependence of the measured property on the volume of the added titrant is plotted ( titration curve). Near the equivalence point, there is a sharp change in the composition and properties of the titrated solution, and a jump or break is recorded on the titration curve. For example, a jump in the potential of an electrode dipped into a solution. The position of the eq is estimated from the position of the inflection on the curve. This type of analysis is more laborious and time consuming than conventional titration, but gives more accurate results. In one titration, it is possible to determine separately the concentrations of a number of components.

More than a dozen variants of instrumental titrimetry are known. The American analyst I. Koltgof played an important role in their creation. Appropriate methods differ in the measured property of the solution, in the equipment used and in analytical capabilities, but all of them are more sensitive and selective than indicator or non-indicator visual titrimetry options. Instrumental control is especially important when indicators cannot be used, for example, in the analysis of cloudy or intensely colored solutions, as well as in the determination of trace impurities and in the analysis of mixtures. However, instrumental titrimetry requires the laboratory to be equipped with special instruments, preferably self-recording or fully automated, which is not always economically feasible. In many cases, sufficiently accurate and reliable results can be obtained in a simpler and cheaper way based on the use of indicators.

Using indicators. A small amount of a special reagent can be added to the titrated sample in advance - indicator. The titration will have to be stopped at the moment when the indicator under the action of the introduced titrant changes its visible color, this is the end point of the titration. It is important that the color change does not occur gradually, as a result of the addition of just one "extra" drop of titrant. In some cases, the indicator does not change its color, but the solubility or the nature of the glow. However, such indicators (adsorption, fluorescent, chemiluminescent, etc.) are used much less frequently than color indicators. The change in color of any indicator occurs due to the chemical interaction of the indicator with the titrant, leading to the transition of the indicator to a new form. The properties of the indicators must be considered in more detail.

Indicators

In analytical laboratories, several hundred color indicators of various types are used (acid-base, metal-chromic, adsorption, etc.). Once upon a time, tinctures obtained from plants - from violet flowers or from a special type of lichen (litmus) were used as indicators. R. Boyle was the first to use such indicators. Currently, natural indicators are not used, since they are always a mixture of different substances, so the transition of their color is not clearly expressed. Modern indicators are specially synthesized individual organic compounds. As a rule, indicators are compounds of the aromatic series, the molecules of which contain several functional groups (substituents). Many similar compounds are known, but only some of them can be used as color indicators. The proposed indicator must meet a number of requirements:

· the indicator should dissolve well, giving solutions that are stable during storage;

· in solution, the indicator must exist in several forms, different in molecular structure. A mobile chemical balance must be established between the forms. For example, the acid form of the indicator goes into the basic one (and vice versa), the oxidized one goes into the reduced one (and vice versa); the metallochromic indicator is reversibly complexed with metal ions, etc.;

· color indicator should intensively absorb light in the visible region of the spectrum. The color of its solution should be distinguishable even at very low concentrations (10 -6 - 10 -7 mol/l). In this case, it will be possible to introduce very small amounts of the indicator into the titrated solution, which contributes to obtaining more accurate analysis results;

· different forms of the indicator should be different in their color, that is, in the absorption spectrum in the visible region. In this case, during the titration, a contrasting color transition will be observed. For example, the transition of the indicator color from pink to emerald green is clearly visible to the eye. It is much more difficult to fix the end point of the titration (c.t.t.) by the transition of the pink color to orange or purple. It is very important how different the absorption spectra of the two forms of the indicator are. If one of the forms of the indicator maximally absorbs light with a wavelength of λ 1, and the other with a wavelength of λ 2, then the difference ∆λ = λ 1 - λ 2 characterizes the contrast of the color transition. The larger ∆λ, the better the color transition of the indicator is perceived by the eye. To increase the visual contrast of the color transition, mixtures of different indicators are sometimes used or an inert foreign dye is added to the indicator;

· the transition of the indicator from one form to another when changing the composition of the solution should take place very quickly, in a fraction of a second;

· the transition must be caused by a single factor, the same for all indicators of a given type. Thus, a change in the color of an acid-base indicator should not occur due to reactions of another type, for example, when interacting with oxidizing agents, or metal ions, or proteins! On the contrary, redox indicators should change their color only due to interaction with oxidizing and reducing agents, and this should occur at a certain potential specific to each redox indicator. The color of these indicators and the transition potential should not depend on the pH of the solution. Unfortunately, in practice, the transition potential of many redox indicators also depends on pH.

To reduce the influence of side processes, sometimes the indicator is not introduced into the titrated solution, but, on the contrary, during the titration, a drop of the titrated solution is periodically taken, mixed on a watch glass with a drop of the indicator solution and the color obtained is observed. This technique allows the use of irreversibly reacting indicators. It is more convenient to work with the “external indicator” if the impaper is soaked in advance.

The end point of the titration, as determined by the color change of the indicator, may not coincide with the equivalence point. Mismatch V k.t.t. And V t.eq leads to a systematic error in the analysis result. The error value is determined by the nature of this indicator, its concentration and the composition of the titrated solution.

The principle of selecting indicators is very simple and universal. : The transition characteristic of the indicator (pT-value of the titration, transition potential, etc.) must correspond to the expected composition of the titrated solution at the equivalence point. Thus, if an analyst titrates an aqueous solution of a strong acid with a strong base, the solution will have pH = 7 at the equivalence point. Therefore, an acid-base indicator should be used that changes color at approximately pH 7 (bromothymol blue, etc.). about pT - titration indicators for indicators of various types are in the reference literature.

Calculation of titrimetric analysis results

The results of titrimetric analysis are not recommended to be calculated directly from the reaction equation, for example, using proportions. Such a "school" way of solving computational problems is irrational and, as a rule, does not provide the required accuracy. The results of titrimetric analysis are calculated using one of several ready-made algebraic formulas derived from the law of equivalents. The initial data will be the volume of the spent titrant (in milliliters) and the concentration of the titrant (in mol / liter), they must be set with the required accuracy.

The method of calculation does not depend on the type of chemical reaction that occurs during the titration, and the method of monitoring the equivalence point (indicator, device, etc.). The choice of the calculation formula is determined by which titration method (direct, reverse, substitution) is used during the analysis. When choosing a formula, two cases should be distinguished: a) calculation of the concentration of X solution; b) determination of the mass fraction of the component (percentage of X in the sample).

The calculation formulas look the simplest if the concentrations of the component to be determined and the titrant are expressed as the number of moles of their equivalents per liter of the corresponding solutions, i.e. concentrations of the analyte are used ( N x ) and titrant (N T ), expressed as the number of moles of equivalent per liter of solution. Previously, these concentrations were called normal. Now this term is not recommended to be used, but in practice it is used very widely, especially in redoxmetry. But in complexometry and in some other methods, where 1 mol of analyte X always reacts with 1 mol of titrant, normal concentrations coincide with the usual molar concentrations ( C x and C T ), and therefore there is no need to use normal concentrations and equivalents when calculating the results.

Unlike conventional molar concentrations, the normal concentration is determined by taking into account the chemistry of the reaction that occurs during the titration. It is useful to remember that the normal concentration of X in a solution is either equal to its molar concentration, or exceeds it by several (2,3,4 ....) times, depending on how many protons (or electrons) participate in the reaction, per particle X. When writing the reaction equation, determining equivalents and calculating normal concentrations, one should take into account the conditions under which the titration takes place, and even the choice of indicator.

Weighttitrated Xatdirect titration equals (in mg):

m x =N T . V T . E x , (1),

where E x - the molar mass of the equivalent of X, corresponding to one proton (in acid-base reactions), one electron (in redox reactions), one ligand (in complex formation reactions), etc. V T is the volume of titrant (in ml). In complexometry, the mass of the analyte (in mg) is best calculated using the formula, which includes the value M x -molar mass X:

m x = C T . V T . M x (2).

From (4.11) it follows that the mass fraction of X in the sample sample, expressed in%, is equal to:

%X = N T . V T . E x . 100% / m S , (3),

where m S - weight of sample in mg. Usually, the result of titration does not depend on the volume of water in which the sample sample was dissolved before titration, and this volume is not taken into account in the calculations. If not all of the sample is titrated, but some part of it (an aliquot), then an additional coefficient must be taken into account TO , equal to the ratio V0 - the volume of the solution into which this sample was transferred and from which aliquots were taken, to Valiq - the volume of one aliquot:

m x = K. N T . V T . E x , (4).

When calculating concentrationaccording to the method of direct (or substitution) titration, a simple formula is used, directly following from the law of equivalents:

N x . V x \u003d N T. V T (5).

analysis, however, in factory laboratories, other methods of calculation are also used.

Preparation of working solutions in titrimetry

Working solutions of precisely known concentrations used in titrimetric analysis are prepared in several ways:

· according to the exact weighing of the chemical reagent taken on an analytical balance. This sample is dissolved in a small amount of solvent, and then the volume of the resulting solution is brought to the mark in a volumetric flask. The resulting solutions are called standard, and the corresponding reagents are called primary standards. Few substances can be primary standards - they must be pure chemicals of constant and precisely known composition, solid at room temperature, stable in air, non-hygroscopic and non-volatile. Examples are potassium dichromate, complexone III, oxalic acid. On the contrary, it is impossible to prepare a standard solution of hydrochloric acid (the “hydrochloric acid” reagent is a liquid with an inaccurately known composition), ferrous chloride (it oxidizes quickly in air), caustic soda (hygroscopic) and many other substances.

· from fixed channels. This term refers to a sealed glass ampoule that contains a certain amount of a reagent, usually 0.1000 mol equivalent. Fixanals are prepared in the factory. If in the laboratory the contents of fixanal are quantitatively transferred to a 1000 ml volumetric flask and brought to the mark with the solvent, a liter of exactly 0.1000 N solution will be obtained. The preparation of fixanal solutions not only saves the analyst's time, but also makes it possible to prepare solutions with precisely known concentrations from substances that do not have the complex of properties required for primary standards (for example, fixanal solutions of hydrochloric acid, ammonia, or iodine).

· according to an approximately known sample of a chemical reagent, taken on a technical scale. This sample is dissolved in an approximately known amount of solvent. Then an additional operation is carried out - standardization of the resulting solution. For example, an exact weight of another substance (primary standard) is titrated with the resulting solution. You can do it differently: take a known volume (aliquot) of the prepared solution and titrate it with a suitable standard solution. The exact concentration of the prepared solution is calculated from the volume used for titration. Such solutions are called standardized. For example, a KOH solution is standardized by a weighed portion of oxalic acid or by a fixanal solution of hydrochloric acid. If the substance in the laboratory is available in the form of a concentrated solution of approximately known concentration (for example, hydrochloric acid), then instead of weighing it, a certain pre-calculated volume of the concentrated solution is measured. This requires knowledge of the density of the initial solution. Then, as in the previous case, the resulting solution is standardized.

The concentration of solutions should not spontaneously change during storage. In this case, pre-prepared (standard or standardized) solutions can be used for titrations without any additional operations. It should be noted that the more dilute the solution, the less stable it is, as a rule, during storage (hydrolysis of the solute, its oxidation with oxygen air, adsorption on the inner surface of glassware, etc.). Therefore, working solutions with a low concentration, as a rule, are not prepared in advance. They are prepared only as needed, on the day of use. To do this, the initial (standard, fixed or standardized) solutions are diluted with a pure solvent in a precisely known number of times (usually, the solution is diluted 5 or 10 times in one operation). If even more dilute solutions are required, this operation is repeated. For example, 0.01 M is prepared from a 0.1 M solution, 0.001 M from that, etc.

The preparation of solutions with precisely known concentrations requires the use of a whole set of special volumetric utensils that allow you to measure volumes with the required accuracy. These are volumetric flasks, pipettes and burettes. The manuals for laboratory work contain descriptions of measuring utensils and the rules for working with them.

Titration Methods

Single sample method and aliquot method. To reduce the influence of random errors, the titration is usually repeated several times, and then the results are averaged. Repeated analyzes can be carried out in two different ways: by the method of separate portions by the method of aliquots. Both methods are used both in the standardization of working solutions and directly in the analysis of real objects.

Single weight method, as its name implies, assumes that several weighed portions of the analyzed material are taken for titration. Their masses should be approximately equal. The sample size is chosen taking into account the desired consumption of titrant per titration (no more than the volume of the burette) and taking into account the concentration of the titrant.

Let three weighed portions of oxalic acid be taken, the masses of which are indicated in Table 2. According to each titration, calculate (separately!) the concentration of KOH. Then the concentrations are averaged. The volumes spent on titration of different samples cannot be averaged!

Table 2. An example of calculating the results of analysis using the method of individual samples

Hinge number	Massanaveski, mg	Volume of titrant, ml	Found KOH concentration, mol/l
	95,7	14,9	0,102
	106,9	16,2	0,105
	80,8	12,7	0,101
Average result of analysisСKOH = 0.103 mol/l

Aliquot titration method (or pipetting method) is based on the titration of several separate aliquots - small volumes of the test solution, taken with pipettes.

The method of separate weights and the method of titration of aliquots is used not only for direct titration, as shown in the examples given, but also for reverse and substitution titration. When choosing a titration method, it should be taken into account that the method of separate weights gives more accurate results, but it is more laborious and requires more calculations. Therefore, it is better to use the method of separate weights for standardization of working solutions, and for serially performed analyzes, use the more express method of aliquots.

Shape of titration curves

Logarithmic titration curves represent a graphical dependence of the logarithm of the equilibrium concentration of one of the reagents on the volume of added titrant. Instead of the logarithm of the concentration, the pH value of the solution (pH) is usually plotted on the vertical axis. Other similar indicators are also used (for example, pAg \u003d - lg), as well as the value of those physico-chemical properties of the titrated solution, which linearly depend on the logarithms of equilibrium concentrations. An example would be the electrode potential (E).

If the solution contains only one substance that reacts with the titrant, and the reaction is described by a single chemical equation (that is, it does not proceed in steps), an almost vertical section is observed on the logarithmic curve, called jump titration . On the contrary, sections of the curve away from t.eq. close to horizontal. An example can be the dependence of the pH of solutions on the volume V of the added titrant, shown in Fig. 1

Fig.1. Type of titration curves

The higher the jump height on the tiding curve, the more accurately the equivalence point can be fixed.

Acid-base titration (neutralization method)

Method principle

The neutralization method is based on carrying out acid-base (protolytic) reactions. During this titration, the pH value of the solution changes. Acid-base reactions are most suitable for titrimetric analysis: they proceed according to strictly defined equations, without side processes and at a very high speed. The interaction of strong acids with strong bases leads to high equilibrium constants. To detect k.t.t. there is a convenient and well-studied way - the use of acid-base indicators. Instrumental methods can also be used and are especially important when titrating non-aqueous, cloudy or colored solutions.

The neutralization method includes two options − acidimetry(titrant is a strong acid solution) and alkalimetry(titrant is a solution of a strong base). These methods are respectively used to determine bases and acids, including ionic and multiprotonic ones. The ability to titrate strong protoliths is determined by their concentration; titration is possible if C x> 10 - 4 M .During such titration in an aqueous solution, the following reaction occurs:

H 3 O + +OH - ® 2 H 2 O

Titration of weak acids and weak bases in aqueous solutions follows the schemes:

ON+OH - ® H 2 O (alkalimetry)

B + H 3 O + ® HB + + H 2 O (acidimetry)

Examples of practical applications of acid-base titration:

· determination of the acidity of food products, soils and natural waters (alkalimetric titration of aqueous solutions with the indicator phenolphthalein);

· determination of the acidity of petroleum products (alkalimetric titration of non-aqueous solutions with instrumental control of rt);

· determination of carbonates and bicarbonates in minerals and building materials (acidimetric titration of aqueous solutions with two indicators);

· determination of nitrogen in ammonium salts and in organic substances (Kjeldahl method). In this case, organic nitrogen-containing substances are decomposed by boiling with concentrated sulfuric acid in the presence of mercury salts, ammonium nitrogen is driven off by the action of alkali when heated, ammonia is absorbed with a standard solution of HCl, taken in excess. The unreacted part of HCl is then titrated with alkali in the presence of methyl orange indicator. This method uses both the substitution principle and the back titration method.

working solutions.In acidimetric titration of aqueous solutions, titrants are used as solutions of strong acids (HCl, less often HNO 3 or H 2 SO 4). IN alkalimetry titrants - solutions of NaOH or KOH. However, the listed reagents do not have properties that would make it possible to prepare standard solutions from them simply by accurately weighing them. So, solid alkalis are hygroscopic and always contain impurities of carbonates. In the case of HCl and other strong acids, the starting reagent is not a pure substance, but a solution with an inaccurately known concentration. Therefore, in the neutralization method, a solution with an approximately known concentration is first prepared, and then it is standardized. Acid solutions are standardized for anhydrous sodium carbonate Na 2 CO 3 (soda) or sodium tetraborate Na 2 B 4 O 7 . 10H 2 O (storm). Borax, when dissolved, interacts with water:

B 4 O 7 2– + 3H 2 O \u003d 2H 3 VO 3 + 2VO 2 -

The resulting metaborate is a fairly strong base. It is titrated with acid:

IN 2 - + H 3 O + \u003d H 3 IN 3.

Obviously, the molar mass of the borax equivalent is M(½Na 2 B 4 O 7 . 10H 2 O) = 190.71 g/mol. The high molar mass equivalent is an advantage of borax as a primary standard. Alkali solutions are standardized for potassium hydrophthalate. The hydrophthalate molecule contains a mobile proton and has the properties of a weak acid:

Benzoic acid C 6 H 5 COOH, oxalic acid H 2 C 2 O 4 are often used as standards. . 2H 2 O and other weak organic acids (solid, pure stable substances). Standard 0.1000 M solutions of acids and bases in laboratories are usually prepared from fixanals. The prepared acid solution can be used to standardize the alkali solution, and vice versa. Standardized acid solutions are stable and can be stored without change for an arbitrarily long time. Alkali solutions are less stable; it is recommended to store them in waxed or fluoroplastic containers to prevent interaction with glass. It must be taken into account that alkali solutions absorb CO 2 from the air, and during storage they are protected with a tube filled with quicklime or soda lime.

Rice. 2. Neutralization curves for a strong acid.

1 - 0.1 M, 2 - 0.01 M, 3 - 0.001 M.

To detect k.t.t. with a color indicator, it is necessary that the height of the jump be greater than the width of the transition interval of the indicator. The latter is usually about two pH units.

The height of the jump on the neutralization curve for weak acids depends on the strength of the acid (the value of its acid constant, or pK a ). Namely, the weaker the acid (the greater the value of pK a), the smaller, all other things being equal, should be the height of the jump.

1 - hydrochloric acid, 2 - acetic acid (pK a = 4.8), 3 - hydrocyanic acid (pK a = 9.2).

The height of the jump should be greater than the width of the transition zone of the indicator, which is usually 2 pH units. Therefore, to As in the case of strong electrolytes, titration criterion weak protolith with a 1% error can be derived from the condition ∆p Н ±1% ≥ 2. For an aqueous solution of a weak acid, we obtain the desired criterion in the following form:

R TOa+ p WITH≤ 8

At p C \u003d 2, the critical value p K a equals 6. In other words, if the acid is very weak and its pK A greater than 6, then it cannot be accurately titrated with color indicators.

Titration of mixtures of protoliths and multiproton protoliths. In mixed solutions, strong acids inhibit the protolysis of weaker ones. The same is observed in solutions containing a mixture of bases of different strengths. When a titrant is added to such a mixture, the stronger protolith is titrated first, and only then the weaker one reacts with the titrant. However, the number of jumps observed on the mixture titration curve depends not only on the number of protoliths present, but also on the absolute values of the corresponding acidity (basicity) constants, as well as on their ratio. The acidity (or basicity) constants of the components of the mixture should differ by more than 10 4 times, only in this case distinctly pronounced titration jumps will be observed separately on the titration curve, and the relative error in determining each component will not exceed 1%. The criterion for the possibility of separate titration of protoliths is the so-called "rule of four units":

(6)

Multiproton protoliths react with titrants stepwise, first in the first stage, then in the second, etc., if the corresponding acidity constants differ in accordance with condition (6). When calculating the neutralization curves, multiproton protoliths can be considered as mixtures of different electrolytes.

As an example, consider the possibility

Fig.5. Titration curve of a mixture of carbonate and bicarbonate ions with a solution HCl.

The pH values at which color transitions of the indicators are observed are indicated.

When titrating a mixture of two strong acids, a mixture of two equally weak acids, or a mixture of two bases with close p TOb there are no two separate jumps on the titration curve. However, it is still quite possible to determine the concentration of the components of such mixtures separately. These problems are successfully solved using differentiating non-aqueous solvents.

Acid-base indicators and their selection

To detect k.t.t. the neutralization method traditionally uses acid-base indicators - synthetic organic dyes that are weak acids or bases and change their visible color depending on the pH of the solution. Examples of some (most commonly used in laboratories) acid-base indicators are shown in table 3. Structure and properties indicators are given in reference books. The most important characteristics of each acid-base indicator are transition interval And titration index (pT). The transition interval is the zone between two pH values corresponding to the boundaries of the zone, within which a mixed color of the indicator is observed. So an observer will characterize an aqueous solution of methyl orange as pure yellow - at pH< 3,1 и как чисто красный при рН >4.4, and between these boundary values, a mixed, pink-orange color of different shades is observed. The width of the transition interval is typically 2 pH units. Experimentally determined transition intervals of indicators are in some cases less or more than two pH units. This, in particular, is explained by the different sensitivity of the eye to different parts of the visible region of the spectrum. For single-color indicators, the width of the interval also depends on the concentration of the indicator.

Table 3

The most important acid-base indicators

Indicator	Transition interval ΔрН Ind	R TOa(HInd)	Color change
methyl orange			Red - yellow
Bromocresol green			Yellow - blue
methyl red			Red - yellow
Bromocresol purple			Yellow - purple
Bromothymol blue			Yellow - blue
Phenol red			Yellow - red
thymol blue
Phenolphthalein			Colorless - red

Knowing the characteristics of different indicators, it is possible to theoretically reasonably select them in order to obtain the correct analysis results. The following rule is followed: the transition interval of the indicator must lie in the region of the jump on the titration curve.

When choosing indicators for titration of weak protoliths, it should be taken into account that t.eq. and the titration jump are shifted to a weakly alkaline medium when titrating an acid and to a slightly acidic medium when titrating a base. Hence, for titration of weak acids, indicators that change color in a slightly alkaline medium (for example, phenolphthalein) are suitable, and for titration of a weak base, indicators that change color in a slightly acidic medium (for example, methyl orange

There is another characteristic of each acid-base indicator - it is titration index ( RT ). This is the pH value at which the observer most clearly notices the change in the color of the indicator and at this moment considers the titration to be completed. Obviously pT = pH K.T.T. . When choosing a suitable indicator, one should strive to ensure that the pT value is as close as possible to the theoretically calculated value. pH T.EKV .. Typically, the pT value is close to the middle of the transition interval. But pT is a poorly reproducible quantity. Different people doing the same titration with the same indicator will get significantly different pT values. In addition, the pT value depends on the order of the titration, that is, on the direction of color change. When titrating acids and bases with the same indicator pT values will vary slightly. For monochromatic indicators (phenolphthalein, etc.), the pT value also depends on the concentration of the indicator.

Filled with titrant to the zero mark. Titration starting from other marks is not recommended, as the burette scale may be uneven. Burettes are filled with working solution through a funnel or with the help of special devices if the burette is semi-automatic. The end point of the titration (equivalence point) is determined by indicators or physico-chemical methods (by electrical conductivity, light transmission, indicator electrode potential, etc.). The results of the analysis are calculated by the amount of the working solution used for titration.

Types of titrimetric analysis

Titrimetric analysis can be based on various types of chemical reactions:

acid-base titration - neutralization reactions;
redox titration (permanganatometry, iodometry, chromatometry) - redox reactions;
precipitation titration (argentometry) - reactions occurring with the formation of a poorly soluble compound, while changing the concentration of precipitated ions in solution;
complexometric titration - reactions based on the formation of strong complex compounds of metal ions with a complexone (usually EDTA), while changing the concentration of metal ions in the titrated solution.

Titration types

A distinction is made between direct, back, and substituent titration.

At direct titration to a solution of the analyte (an aliquot or a sample, a titratable substance) add a titrant solution (working solution) in small portions.

At back titration first, a known excess of a special reagent is added to the solution of the analyte, and then its residue, which has not entered into the reaction, is titrated.

At substitution titration first, a certain excess of a special reagent is added to the solution of the analyte, and then one of the reaction products between the analyte and the added reagent is titrated.

Links

Wikimedia Foundation. 2010 .

Types of titration. Analytical chemistry Titration in brief

Send your good work in the knowledge base is simple. Use the form below

Similar Documents

Classification of methods of titrimetric analysis.

Types of titrimetric analysis

Titration types

see also

Links

Send your good work in the knowledge base is simple. Use the form below

Similar Documents

Classification of methods of titrimetric analysis.

Types of titrimetric analysis

Titration types

see also

Links

The concept of an electronic textbook, its distinctive characteristics Building diagrams using VBA

The problem of interethnic conflicts in the history of mankind

The reign of Nicholas I presentation for a history lesson (Grade 10) on the topic The main stages in the development of the Eastern question