Displacement titrations (Key words: Bronsted bases, acetates, carbonates, borates)

 Click here to visit our official web site and for more info; www.generalchemistryguide.com

Neutralization of anions of weak acids (Bronsted Bases) with strong acids

Theory

Titrations are also possible with weak bases (Bronsted bases) such as acetates, carbonates, borate ions. Here the OH- produced due to hydrolysis of the salt will react with the strong acid.

The weak acetic acid was apparently displaced by the strong HCl, and the process was referred to as displacement titration.

The so called titration of solutions of hydrolysis is merely the titration of a weak base with a strong acid (highly ionized).

Titration of carbonate ions with a strong acid

Determination of the concentration of carbonate using 0.1M HCl

Theory

The dissociation constant for H2CO3 are,

K1 = 4.3 x 10^-7, pK1 = 6.37
K2 = 4.67 x 10^-11, pK2 = 10.33
A solution of carbonate ion can be titrated to the hydrogen carbonate stage with HCl acid.

Equation 1 -The equivalence point for the primary stage of ionization
Equation 2 -The solution may also be titrated until all the carbonic acid is displaced.
Equation 3 -The same end point is reached by titrating hydrogen carbonate solution with HCl.
The pH of this equivalence point can be calculated if the concentrations are known. It is approximately 3.7 (for equal volumes of 0.1M HCl and 0.1M Sodium hydrogen carbonate)

If you cant see the image clearly, please click on it.

Procedure I

Pipette out 25.00mL of the given carbonate solution into a titration flask, add few drops of phenolphthalein indicator and titrate the solution with 0.1M HCl. Repeat your titration using methyl orange indicator instead of phenolphthalein.

I. Calculate the concentration of carbonate solution using the data obtained for the two indicators separately.
II. Comment on the results obtained.


Procedure II

Pipette out 25.00 mL of the carbonate solution into a flask, add few drops of phenolphthalein indicator. Titrate the solution to end point with 0.1M HCl and then add few drops of methyl orange indicator. Continue the titration to its end point.

1. Calculate the concentration of carbonate solution.
2. Comment on the two titration readings.
3. Suggest an experiment to determine the dissociation constants of the carbonic acid. Draw the corresponding titration curve for your suggested procedure.
4. You are provided with a solution containing a mixture of sodium hydroxide and sodium carbonate, a solution of standardized HCl and some common indicators. Describe a method to determine the concentration of sodium hydroxide and sodium carbonate in the given mixture.

Read more...

Standardization of HCl and NaOH using a primary standard solution

Click here to visit our official web site and for more info; www.generalchemistryguide.com 

Theory

The net result of the displacement titration between the tetraborate ion with hydrochloric acid is,

The pH at the equivalence point in the titration of 0.2M Sodium tetraborate with 0.2M HCl is 5.6.

The boric acid so formed is a weak monoprotic acid (Ka = 6.4 x 10^ -10). Therefore it cannot be directly titrated with standard alkali. However by the addition of certain organic polyhydroxy compounds it is converted to a much stronger acid which can be titrated using phenolphthalein. This is due to the complex formation between hydrated borate ion and 1,2 or 1,3 diols.

If you can't see the image clearly please click on it.

Procedure

Prepare a 0.05M solution of borax in a 250mL volumetric flask by weighing required amount accurately. Pipette out 25.00 mL of the borax solution in to a titration flask and add few drops of methyl red indicator. Titrate this solution with the HCl solution. Repeat the titration with two other portions.

I. Calculate the mean and the standard deviation of your reading.
II. Using the mean value calculate the concentration of the HCl acid.
III. Draw the titration curve.
IV. What are the advantages of using borax as a primary standard to standardize strong acids?

Pipette out 25.00 mL aliquots of the borax solution into two conical flasks to perform a duplicate determination. Add the volume of standard HCl determined by the above titration. Cover the flask with a watch glass, heat to simmering temperature for 7-8 minutes to expel carbon dioxide and then cool the solution to room temperature. Introduce one drop of methyl red indicator and if necessary add just sufficient NaOH to restore the basic transition color of the indicator.

Introduce 2gof Glucose, swirling gently to dissolve. Add 2 drops of phenolphthalein and titrate with the NaOH provided. When the first permanent pink color is produced add further 0.5g glucose. If the pink color disappears, titrate with more NaOH until the pink color reappears.

Standardize the NaOH against the HCl using phenolphthalein indicator.

I. Calculate the concentration of NaOH,
a) From the titrations against HCl and
b) From the titrations in which glucose was added to the borax



II. Explain the important stages of the above procedure.
III. Write down the titration methods to determine concentrations of each component in a mixture of,
a) Boric acid and strong acid
b) Sodium tetraborate and boric acid

Read more...

Precipitation titrations Key words: Argentimetry, Fractional precipitation, Solubility product

Determination of the chloride ion concentration using silver nitrate.

Theory

A titrimetric method based on the formation of a slightly soluble precipitate is called a precipitation titration. The most important precipitation process in titrimetric analysis utilizes silver nitrate as the reagent (Argentimetric process).

Many methods are utilized in determining end points of these reactions, but the most important method, the formation of a colored precipitate will be considered here.

In the titration of a neutral solution of chloride ions with silver nitrate, a small quantity of potassium chromate solution is added to serve as the indicator. At the end point the chromate ions combine with silver ions to form the sparingly soluble brick-red silver chromate.This is a case of fractional precipitation, the two sparingly soluble salts being AgCl (Ksp = 1.2 x 10^-10) and Ag2CrO4 (Ksp = 1.7x10^-12).

AgCl is the less soluble salt and initially chloride concentration is high, hence AgCl will be precipitated. Once the chloride ions are over and with the addition of small excess of silver nitrate solution brick red color silver chromate becomes visible. The titration should be carried out in neutral solution or in very faintly alkaline solution. i.e. within the pH range 6.5-9.

In acid solutions following reaction occurs.

If you cant see the image clearly please click on it.

Consequently the chromate ions concentration is reduced and the solubility product of silver chromate may not be exceeded. In markedly alkaline solution, silver hydroxide (Ksp = 2.3 x 10^-8) might be precipitated.

Procedure

Pipette out 25.00 mL of the chloride solution into a titration flask and add 1mL of the potassium chromate solution. Titrate this solution with 0.1M silver nitrate solution.

1. Calculate the concentration of the chloride solution
2. Sketch the titration curve for the above titration.
3. Calculate the concentration of chloride, silver and chromate ions at the equivalence point.

Read more...

Chemical Titration as a Volumetric Analysis method in Physical Chemistry

 For more info; visit www.generalchemistryguide.com

Titration is a common laboratory method of quantitative chemical analysis that is used to determine the unknown concentration of a known reactant. Because volume measurements play a key role in titration, it is also known as volumetric analysis. A reagent, called the titrant or titrator, of a known concentration (a standard solution) and volume is used to react with a solution of the analyte or titrand, whose concentration is not known. Using a calibrated burette to add the titrant, it is possible to determine the exact amount that has been consumed when the endpoint is reached. The endpoint is the point at which the titration is complete, as determined by an indicator. This is ideally the same volume as the equivalence point—the volume of added titrant at which the number of moles of titrant is equal to the number of moles of analyte, or some multiple thereof (as in polyprotic acids). In the classic strong acid-strong base titration, the endpoint of a titration is the point at which the pH of the reactant is just about equal to 7, and often when the solution permanently changes color due to an indicator. There are however many different types of titrations

Read more...

Indicators in Titration

For more info; visit www.generalchemistryguide.com 

Many methods can be used to indicate the endpoint of a reaction; titrations often use visual indicators (the reactant mixture changes colour). In simple acid-base titrations a pH indicator may be used, such as phenolphthalein, which becomes pink when a certain pH (about 8.2) is reached or exceeded. Another example is methyl orange, which is red in acids and yellow in alkali solutions.

Not every titration requires an indicator. In some cases, either the reactants or the products are strongly coloured and can serve as the "indicator". For example, an oxidation-reduction titration using potassium permanganate (pink/purple) as the titrant does not require an indicator. When the titrant is reduced, it turns colourless. After the equivalence point, there is excess titrant present. The equivalence point is identified from the first faint pink color that persists in the solution being titrated.
Due to the logarithmic nature of the pH curve, the transitions are, in general, extremely sharp; and, thus, a single drop of titrant just before the endpoint can change the pH significantly—leading to an immediate colour change in the indicator. There is a slight difference between the change in indicator color and the actual equivalence point of the titration. This error is referred to as an indicator error, and it is indeterminate.

Read more...

History of Titration

For more info; visit www.generalchemistryguide.com 

The word "titration" comes from the Latin word titulus, meaning inscription or title. The French word titre, also from this origin, means rank. Titration, by definition, is the determination of rank or concentration of a solution with respect to water with a pH of 7 (which is the pH of pure H2O under standard conditions).

The origins of volumetric analysis are in late-18th-century French chemistry. Francois Antoine Henri Descroizilles developed the first burette (which looked more like a graduated cylinder) in 1791. Joseph Louis Gay-Lussac developed an improved version of the burette that included a side arm, and coined the terms "pipette" and "burette" in an 1824 paper on the standardization of indigo solutions. A major breakthrough in the methodology and popularization of volumetric analysis was due to Karl Friedrich Mohr, who redesigned the burette by placing a clamp and a tip at the bottom, and wrote the first textbook on the topic, Lehrbuch der chemisch-analytischen Titrirmethode (Textbook of analytical-chemical titration methods), published in 1855

Read more...

Types of titrations -Acid-base titrations, Redox titrations, Complexometric titrations

For more info; visit www.generalchemistryguide.com 

Titrations can be classified by the type of reaction. Different types of titration reaction include:
• Acid-base titrations are based on the neutralization reaction between the analyte and an acidic or basic titrant. These most commonly use a pH indicator, a pH meter, or a conductance meter to determine the endpoint.

• Redox titrations are based on an oxidation-reduction reaction between the analyte and titrant. These most commonly use a potentiometer or a redox indicator to determine the endpoint. Frequently either the reactants or the titrant have a colour intense enough that an additional indicator is not needed.

• Complexometric titrations are based on the formation of a complex between the analyte and the titrant. The chelating agent EDTA is very commonly used to titrate metal ions in solution. These titrations generally require specialized indicators that form weaker complexes with the analyte. A common example is Eriochrome Black T for the titration of calcium and magnesium ions.

• A form of titration can also be used to determine the concentration of a virus or bacterium. The original sample is diluted (in some fixed ratio, such as 1:1, 1:2, 1:4, 1:8, etc.) until the last dilution does not give a positive test for the presence of the virus. This value, the titre, may be based on TCID50, EID50, ELD50, LD50 or pfu. This procedure is more commonly known as an assay.

• A zeta potential titration characterizes heterogeneous systems, such as colloids. Zeta potential plays role of indicator. One of the purposes is determination of iso-electric point when surface charge becomes 0. This can be achieved by changing pH or adding surfactant. Another purpose is determination of the optimum dose of the chemical for flocculation or stabilization

Read more...