Titration calculations precipitation

Run a blank on the reagents, but use 0.1M acid and alkali solutions for the titrations calculate the blank to 0.5M sodium hydroxide. Subtract the blank (which should not exceed 0.5 mL) from the volume neutralised by the original precipitate. 

Muller, B., 2004, ChemEQL V3.0, A program to calculate chemical speciation equilibria, titrations, dissolution, precipitation, adsorption, kinetics, pX-pY diagrams, solubility diagrams. Limnological Research Center EAWAG/ETH, Kastanienbaum, Switzerland. 

For each of the following precipitation titrations, calculate the cation and anion concentrations at equivalence as well as at reagent volumes corresponding to 20.00 mL, 10.00 mL, and 1.00 mL of equivalence. Construct a titration curve from the data, plotting the p-function of the cation versus reagent volume. 

Effect of complexation on solubihty (key equations 11.10, 11.11), p. 345 Calculating precipitation titration curves, p. 346 Indicators for precipitation titrations, p. 349... 

A piece of limestone weighing 0.1965 g was allowed to react with an excess of hydrochloric acid. The calcium in it was precipitated as calcium ethuedioace. The precipitate was dissolved in sulphuric acid, and the ethanedioate In the solution needed 3S.6cm of a 0,0200 mol dm solution of potssdum man ate(Vll) for titration. Calculate the percentage of CaCO, in the limestone. 

F test For two variances, s and s (with chosen to be the larger of the two), the statistic F is defined as F = s /sl. To decide whether Si is significantly greater than 52, we compare F with the critical values in a table based on a certain confidence level. If the calculated value of F is greater than the value in the table, the difference is significant. Fa.jans titration A precipitation titration in which the end point is signaled by adsorption of a colored indicator on the precipitate, false negative A conclusion that the concentration of analyte is below a certain limit when, in fact, the concentration is above the limit, false positive A conclusion that the concentration of analyte exceeds a certain limit when, in fact, the concentration is below the limit. 

Sketching the Titration Curve As we have done for acid-base, complexometric titrations, and redox titrations, we now show how to quickly sketch a precipitation titration curve using a minimum number of calculations. 

Assay of beryUium metal and beryUium compounds is usuaUy accompHshed by titration. The sample is dissolved in sulfuric acid. Solution pH is adjusted to 8.5 using sodium hydroxide. The beryUium hydroxide precipitate is redissolved by addition of excess sodium fluoride. Liberated hydroxide is titrated with sulfuric acid. The beryUium content of the sample is calculated from the titration volume. Standards containing known beryUium concentrations must be analyzed along with the samples, as complexation of beryUium by fluoride is not quantitative. Titration rate and hold times ate critical therefore use of an automatic titrator is recommended. Other fluotide-complexing elements such as aluminum, sUicon, zirconium, hafnium, uranium, thorium, and rate earth elements must be absent, or must be corrected for if present in smaU amounts. Copper-beryUium and nickel—beryUium aUoys can be analyzed by titration if the beryUium is first separated from copper, nickel, and cobalt by ammonium hydroxide precipitation (15,16). 

The holistic thermodynamic approach based on material (charge, concentration and electron) balances is a firm and valuable tool for a choice of the best a priori conditions of chemical analyses performed in electrolytic systems. Such an approach has been already presented in a series of papers issued in recent years, see [1-4] and references cited therein. In this communication, the approach will be exemplified with electrolytic systems, with special emphasis put on the complex systems where all particular types (acid-base, redox, complexation and precipitation) of chemical equilibria occur in parallel and/or sequentially. All attainable physicochemical knowledge can be involved in calculations and none simplifying assumptions are needed. All analytical prescriptions can be followed. The approach enables all possible (from thermodynamic viewpoint) reactions to be included and all effects resulting from activation barrier(s) and incomplete set of equilibrium data presumed can be tested. The problems involved are presented on some examples of analytical systems considered lately, concerning potentiometric titrations in complex titrand + titrant systems. All calculations were done with use of iterative computer programs MATLAB and DELPHI. 

The alkaline solution of thymol is made up to 100 or 200 c.c. as the case may require, using a 5 per cent, soda solution. To 10 c.c. of this solution in a graduated 500 c.c. flask is added a normal iodine solution in shgbt excess, whereupon the thymol is precipitated as a dark reddish-brown iodine compound. In order to ascertain whether a sufficient quantity of iodine has been added, a few drops are transferred into a test tube and a few drops of dilute hydrochloric acid are added. When enou iodine is present, the brown colour of the solution indicates the presence of io ne, otherwise the liquid appears milky by the separation of thymol. If an excess of iodine is present, the solution is slightly acidified with dilute hydrochloric acid and diluted to 500 c.c. From this 100 c.c. are filtered,off, and the excess of iodine determined by titration with normal solution of sodium thiosulphate. For calculation, the number of cubic centimetres required is deducted from the number of cubic centimetres of normal iodine solution added and the resultant figure multiplied by 5, which gives the number of cubia centimetres of iodine required by the thymol. 

A) Preparation of 3-Bromopropyltriphenylphosphonium Bromide Triphenylphosphine, 1.0 kg, and 770 grams of 1,3-dibromopropane are dissolved In 2.0 liters of xylene and the solution is stirred under a nitrogen atmosphere at 130°C. After 20 hours the mixture is cooled, and the crystalline product, which precipitates, is collected and washed with 20 liters of benzene. After drying in vacuo the product weighs 1,578 grams, MP 229°-230°C titration for bromide ion Found, T7.1% calculated, 17.2%. 

Chlorinity When a sample of sea water is titrated with silver nitrate, bromides and iodides, as well as chlorides are precipitated. In calculating the chlorinity (Cl), the entire halogen content is taken as chloride, and chlorinity is defined as the weight in grams of silver required for precipitation of total halogen content per kilogram of sea water, multiplied by 0-328 533. (Chlorinity is always expressed as parts per thousand, using the symbol %o.)... 

The chemical compositions of the samples, obtained from chemical analyses are reported in Table 1. In order to check the chemical analyses, the mother and washing liquors were collected, analysed and their acidity was titrated. In all cases, the alkaline cations were detected only as traces. The acidimetric titration allowed us to determine the HPA amount remaining in the solution. On the other hand, the samples separated after precipitation and washings were weighted in order to calculate the precipitate yields. The results are reported in table 1 where the samples are designated as MxY (M being the alkaline or ammonium cation, Y the heteroatom, x the stoichiometry deduced from chemical analyses. 

The determination of neomycin by non-aqueous titration has been described by Penau et all2l. Neomycin base is allowed to react with standardised perchloric acid the excess acid is then back-titrated with potassium hydrogen phtha-late using crystal violet as indicator. To determine the neomycin content of the sulphate salt the same authors precipitated the sulphate with benzidine before reacting the neomycin with perchloric acid. The amount of benzidine required to precipitate the sulphate is calculated from the sulphate content which is itself determined by titration with sodium hydroxide. 

Common chemical titrations include acid-base, oxidation-reduction, precipitation, and complexometric analysis. The basic concepts underlying all titration are illustrated by classic acid-base titrations. A known amount of acid is placed in a flask and an indicator added. The indicator is a compound whose color depends on the pH of its environment. A solution of base of precisely known concentration (referred to as the titrant) is then added to the acid until all of the acid has just been reacted, causing the pH of the solution to increase and the color of the indicator to change. The volume of the base required to get to this point in the titration is known as the end point of the titration. The concentration of the acid present in the original solution can be calculated from the volume of base needed to reach the end point and the known concentration of the base. 

In the next chapter, you will extend your knowledge of equilibria involving aqueous ions. You will learn how to calculate the pH at an equivalence point, so you can select an appropriate indicator for any acid-hase titration. You will also learn why equilihrium is important to the solubility of compounds that are slightly soluble, and how to predict whether a precipitate will form as the result of a reaction between ions in solution. 

In some cases the molecular-weight distribution can be determined by turbi-dimetric titration, a technique which is based on the fractional precipitation. A precipitant is added to a very dilute solution of the polymer, and the resulting turbidity is measured as a function of the amount of added precipitant the preparative separation of the fractions is thereby avoided. If the polymer is chemically homogeneous, the mass distribution function can then be calculated. Tur-bidimetric titration is also suitable as a means for establishing the best fractionation conditions (e.g., choice of solvent/precipitant combination, size of fractions, etc.), before carrying out a full-scale fractionation by precipitation. 

By titration of the resulting acid the original quantity of thiosulphate can be calculated. This method has the especial advantage of being applicable in the presence of sulphides, which give a similar precipitate, but do not produce acidity.1 A similar process has been proposed using silver nitrate as precipitant 2... 

We now turn our attention to details of precipitation titrations as an illustration of principles that underlie all titrations. We first study how concentrations of analyte and titrant vary during a titration and then derive equations that can be used to predict titration curves. One reason to calculate titration curves is to understand the chemistry that occurs during titrations. A second reason is to learn how experimental control can be exerted to influence the quality of an analytical titration. For example, certain titrations conducted at the wrong pH could give no discernible end point. In precipitation titrations, the concentrations of analyte and titrant and the size of Ksp influence the sharpness of the end point. For acid-base titrations (Chapter 11) and oxidation-reduction titrations (Chapter 16). the theoretical titration curve enables us to choose an appropriate indicator. 

Example Calculating Concentrations During a Precipitation Titration... 

By now you should understand the chemistry that occurs at different stages of a precipitation titration, and you should know how to calculate the shape of a titration curve. We now introduce spreadsheet calculations that are more powerful than hand calculations and less prone to error. If a spreadsheet is not available, you can skip this section with no loss in continuity. Consider the addition of liters of cation M+ (whose initial concentration is C ) to liters of solution containing anion X- with a concentration C%. 

Concentrations of reactants and products during a precipitation titration are calculated in three regions. Before the equivalence point, there is excess analyte. Its concentration is the product (fraction remaining) X (original concentration) X (dilution factor). The concentration of titrant can be found from the solubility product of the precipitate and the known concentration of excess analyte. At the equivalence point, concentrations of both reactants are governed by the solubility product. After the equivalence point, the concentration of analyte can be determined from the solubility product of precipitate and the known concentration of excess titrant. 

The text claims that precipitation of I is not complete before Cl- begins to precipitate in the titration in Figure 7-8. Calculate the concentration of Ag+ at the equivalence point in the titration of 1 alone. Show that this concentration of Ag+ will precipitate Cl... 

Equation 11-1 is an example of the streamlined calculation introduced in Section 7-4 in connection with precipitation titrations. This equation tells us that the concentration of OH is equal to a certain fraction of the initial concentration, with a correction for dilution. The dilution factor equals the initial volume of analyte divided by the total volume of solution. 

---