COLOR CHANGES AND pH RANGE OF SOME INDICATORS

>> Tuesday, February 8, 2011

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


If you can’t see the image clearly, please click on the image.

 


Read more...

Construction of a titration curve (Neutralization curve)

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

The mechanism of neutralization process can be understood by studying the changes in the hydrogen ion concentration during the course of the approximate titration. The change in pH in the vicinity of the equivalence point is of greatest importance as it enables an indicator to be selected with smallest titration error.

The curve (pH vs the moles of base or acid added) may be evaluated experimentally by
determination of the pH at various stages during the titration by a potentiometric method or by usage of universal indicator.

Universal indicator is a mixture of suitable indicators which gives a color change over a considerable portion of pH range. They are not suitable for quantitative titrations, but may be employed for the determination of approximate pH of a solution.


Procedure


Pipette out 25.00 mL of the solution A into a titration flask and add few drops of the universal indicator. Compare the color with the color chart and record the corresponding pH value of the solution. Fill the burette with the solution B. Add following amounts (X) of solution B from the burette to the flask and record the corresponding pH value after each addition by comparing the color charts provided.

X mL = 12.50, 18.50, 22.50, 24.50, 24.75, 24.80, 24.85, 24.90, 24.95, 25.00, 25.05 25.10,25.15, 25.20, 25.25, 25.30, 25.35, 25.40, 25.45, 25.50, 27.50, 31.50 37.50, 42.50, 50.00

(i) Construct the titration curve using your results
(ii) Thus obtain the ideal pH range for an indicator
(iii) Suggest possible indicators which could be used in this titration



NOTE:-
You can use a pH meter instead of the universal indicator. Follow the same procedure usingthe pH meter and construct the titration curve. You can find the equivalence point in this curve which is the exact pH (a point) the neutralization takes place. When you use the indicators it is a pH range what you get, not the exact point.

Read more...

Titration of a strong base with a strong acid(Determination of the concentration of sodium hydroxide using 0.10 M HCl)


All acid base titrations basically involve the following reaction

As the titration proceeds, initially the pH drops slowly, but nearer the end point, the pH of the solution drops rapidly. i.e. in the vicinity of the equivalence point the rate of change of pH of the solution is very rapid.

If you cant see the image clearly please click on it




Procedure
  1. Prepare a 250.00 mL solution of 0.10M HCl using the concentrated HCl acid provided.
  2. Pipette out 25.00 mL of the given NaOH solution in to the titration flask, add few drops of phenolphthalein indicator. Titrate this solution with the HCl solution you prepared. Repeat the titration using methyl orange as the indicator. 
  3. Obtain two readings in each.
  4. Calculate the concentration of NaOH solution using your results.


The concentration of NaOH calculated when methyl orange is used as the indicator
will be slightly higher than the concentration when phenolphthalein is used.
Phenolphthalein is a weak base which has a pH range 8.0-10.0. Methyl orange is a
weak acid which have the pH range 3.0-5.0. More HCl is consumed when methyl orange
is used as the indicator because it shows the acidic color when the pH of the
medium fall to pH 3.

Total free energy for above neutralization reaction (Room temperature 30o C)

Δ G = - RT ln K
= 8.314Jmol-1 K-1 × 303K × 2.303 × log 10 ^14
= -8.1222 × 1014 J
Where Δ G = Free energy change
R= Universal gas constant
T= Absolute temperature
K= Equilibrium constant


Read more...

Titration of a weak acid with a strong base key words: dissociation of acetic acid





For dilute solutions

The pH of the solution at the equivalence point is
pH = 1/2 pKw + 1/2 pKa - 1/2 pC

Where
= the ionic product of water = 1× 10^14
= the dissociation constant of acetic acid = 1.82 ×10^-5
= the concentration of the salt in moldm-3
Accordingly, for the neutralization of 0.1M acetic acid with 0.1M NaOH at the equivalence point,


pH = 7 + 2.37 - 1/2 (1.3) = 8.37

The initial pH of the 0.1M acetic acid can be calculated using its Ka value to be 2.87.
For other concentrations corresponding pH value can be calculated by,


If you cant see the image clearly, please click on the image
Procedure

Pipette out 25.00mL of the given acetic acid solution into a titration flask and add few drops of thymolphthalein indicator. Titrate this solution with 0.10M NaOH solution.


  1. Calculate the concentration of the acetic acid solution.
  2. Sketch the titration curve.
  3. Comment on choosing indicators for this titration.
  4. Derive an equation for the effective or conditional formation constant of water on the employment of a weak acid, HA, with an acid dissociation constant (Ka)
  5. Calculate the conditional formation constant of water for the above titration.
  6. Calculate the standard free energy change of this titration (Room temperature = 30 oC
  7. A sample of vinegar, weighing 12.0g, was titrated with a 0.500M NaOH solution. 17.50 mL being required to obtain a phenolphthalein end point and 5.5 mL being required to obtain a methyl orange end point. Explain the use of correct indicator system in this titration and calculate the percentage of acetic acid in vinegar.


Discussion

Acetic acid is a weak acid and NaOH is a strong base. When acetic acid is titrated with NaOH, with the final drop of NaOH at the end point, the medium turns alkaline. So an alkaline indicator is suitable for this titration. Thymolphthalein is an alkaline indicator with the pH range of 9.3-10.5, the pH range where this reaction has its end point.
If you cant see the image clearly please click on it.










Read more...

Titration of a weak base with strong acid (Key words: Dissociation constant of ammonia)

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

Determination of the concentration of ammonia using 0.10M HCl

Theory


The pH of the solution at the equivalence point is,

pH = 1/2 pKw -1/2 pKb + 1/2pC


Where,
Kb = the dissociation constant of ammonia = 1.85 x 10^-5

At the equivalence point of the neutralization of 0.1M aqueous ammonia with 0.1M HCl,

pH 7 – 2.87 + ½(1.3) = 5.28

For other concentrations, the pH may be calculated from,


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

Procedure

Pipette out 25.00 mL of the given aqueous ammonia solution into a titration flask, add few drops of methyl orange indicator. Titrate this solution with the 0.100M HCl acid you have prepared.
  1. Calculate the concentration of the aqueous ammonia solution.
  2. Draw the titration curve
  3. Comment on the selection of a suitable indicator for this reaction.
  4. Calculate the pH of the solution at the equivalence point if 100.00 mL of 0.100M methylamine (pKa = 10.64) solution is titrated with 0.100M HCl acid.
  5. Calculate the standard free energy change for the above titration at room temperature.

Read more...

Neutralization of a weak acid with a weak base

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

Theory

For the following reaction, pH of the solution at the equivalence point is,

For the titration of 0.1M acetic acid with 0.1m aqueous ammonia the pH at the equivalence point is 7.1

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



 There is no sudden pH change and hence no sharp end point can be formed with any simple indicator.
  • Calculate the standard free energy change for the neutralization of 0.1M acetic acid with 0.1M aqueous ammonia.
  • Compare this value with the values you obtained for the standard free energy change in previous experiments and comment on the feasibility of this titration.
  • Discuss the feasibility of titration of a solution of formic acid pKa = 3.75) with a solution of piperidine (pKb = 2.80)


Read more...

Neutralization of a polyprotic acid with a strong base (Key words: Phosphoric acid

>> Saturday, February 5, 2011

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

Determination of the concentration of phosphoric acid using NaOH

Theory

The shape of the titration curve will depend on the relative magnitudes of the various dissociation constants. If a dibasic acid H2A has two equivalent independent ionization processes with two dissociation constants K1 & K2, the ratio K1/K2 should be greater than 104 (K1/K2 > 104) in order to titrate the two hydrogen ions separately.
i.e. the acid behaves as a mixture of two acids.
The common polyfunctional acid used for titration is phosphoric (V) acid (orthophosphoric acid). The stepwise dissociation of the acid is represented by the following set of equilibria

Hence the acid will behave as a mixture of three monoprotic acids. The neutralization proceeds almost completely to the end of the primary stage before the secondary stage is affected, and the secondary stage proceeds almost to completion before the tertiary stage apparent.

The pH of the equivalence point is,
pH = 1/2pK1 + 1/2 pK2 = 4.6
At the second equivalence point,

pH = 1/2pK2 + 1/2pK3

and at the third equivalence point,

pH = 1/2pKw + 1/2 pK3 - 1/3 pC

for 0.1M phosphoric acid, pH at the third equivalence point = 12.35

In the third ionization stage, HPO42- behaves as an exceedingly weak acid and the titration curve is very flat (the magnitude of the free energy change is almost zero) and no indicator is available for direct titration. Therefore phosphoric acid behaves identical to a dibasic acid.

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

Procedure

Pipette out 25.00 mL of the phosphoric acid into a titration flask and add few drops of bromocresol green indicator and titrate with the 0.20M NaOH solution.
Repeat the titration using thymolphthalein indicator.

I. Calculate the concentration of phosphoric acid.
II. Draw the titration curve and indicate the use of correct indicators in the titration.
III. Suggest an experimental procedure that you would carry out to determine the dissociation constant of the acid. Carry out the suggested procedure in consultation with your demonstrator. Thus determine the concentration of the acid.
IV. Discuss the possibility of using citric acid as a triprotic acid. The dissociation constants of the acid are:

                  K1 = 9.2 x 10^-4, K2 = 2.7 x 10^-5, K3 = 1.3 x 10^-6

V. You are provided with a mixture of two acids H2A and HB the corresponding dissociation constants are given below.
For H2A K1 = 2 x 10^-2, K2 = 1 x 10^-13
For HB K1 = 3.2 x 10^-7
  • Is it possible to determine the concentration of the two acids in the mixture by titrating with a standard solution of NaOH? Explain your answer giving reasons and calculations.
  • Draw the appropriate titration curve indicating all the necessary features.
  • Thus determine the concentration of individual acids in the mixture if the concentration of NaOH is 0.01M.

Read more...

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

>> Monday, January 24, 2011

 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

>> Friday, January 21, 2011


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

>> Saturday, January 15, 2011


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


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


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...

Measuring the endpoint of a titration


Different methods to determine the endpoint include:

pH indicator: This is a substance that changes colour in response to a chemical change. An acid-base indicator (e.g., phenolphthalein) changes colour depending on the pH. Redox indicators are also frequently used. A drop of indicator solution is added to the titration at the start; when the colour changes the endpoint has been reached.

A potentiometer can also be used. This is an instrument that measures the electrode potential of the solution. These are used for titrations based on a redox reaction; the potential of the working electrode will suddenly change as the endpoint is reached.

pH meter: This is a potentiometer that uses an electrode whose potential depends on the amount of H+ ion present in the solution. (This is an example of an ion-selective electrode.) This allows the pH of the solution to be measured throughout the titration. At the endpoint, there will be a sudden change in the measured pH. It can be more accurate than the indicator method, and is very easily automated.

Conductance: The conductivity of a solution depends on the ions that are present in it. During many titrations, the conductivity changes significantly. (For instance, during an acid-base titration, the H+ and OH- ions react to form neutral H2O. This changes the conductivity of the solution.) The total conductance of the solution depends also on the other ions present in the solution (such as counter ions). Not all ions contribute equally to the conductivity; this also depends on the mobility of each ion and on the total concentration of ions (ionic strength). Thus, predicting the change in conductivity is harder than measuring it.

Colour change: In some reactions, the solution changes colour without any added indicator. This is often seen in redox titrations, for instance, when the different oxidation states of the product and reactant produce different colours.

Precipitation: If the reaction forms a solid, then a precipitate will form during the titration. A classic example is the reaction between Ag+ and Cl- to form the very insoluble salt AgCl. This usually makes it difficult to determine the endpoint precisely. As a result, precipitation titrations often have to be done as "back" titrations (see below).

Read more...

Titration types - An isothermal titration, Thermometric titrimetry, Spectroscopy, Amperometry/amperometric titration


An isothermal titration calorimeter uses the heat produced or consumed by the reaction to determine the endpoint. This is important in biochemical titrations, such as the determination of how substrates bind to enzymes.
Thermometric titrimetry is an extraordinarily versatile technique. This is differentiated from calorimetric titrimetry by the fact that the heat of the reaction (as indicated by temperature rise or fall) is not used to determine the amount of analyte in the sample solution. Instead, the endpoint is determined by the rate of temperature change.
Spectroscopy can be used to measure the absorption of light by the solution during the titration, if the spectrum of the reactant, titrant or product is known. The relative amounts of the product and reactant can be used to determine the endpoint.
Amperometry can be used as a detection technique (amperometric titration). The current due to the oxidation or reduction of either the reactants or products at a working electrode will depend on the concentration of that species in solution. The endpoint can then be detected as a change in the current. This method is most useful when the excess titrant can be reduced, as in the titration of halides with Ag+. (This is handy also in that it ignores precipitates.)

Read more...

Titration curve & Henderson-Hasselbalch equation



Titrations are often recorded on titration curves, whose compositions are generally identical: the independent variable is the volume of the titrant, while the dependent variable is the pH of the solution (which changes depending on the composition of the two solutions). The equivalence point is a significant point on the graph (the point at which all of the starting solution, usually an acid, has been neutralized by the titrant, usually a base).

It can be calculated precisely by finding the second derivative of the titration curve and computing the points of inflection (where the graph changes concavity); however, in most cases, simple visual inspection of the curve will suffice (in the curve given to the right, both equivalence points are visible, after roughly 15 and 30 mL of NaOH solution has been titrated into the oxalic acid solution). To calculate the acid dissociation constant (pKa), one must find the volume at the half-equivalence point, that is where half the amount of titrant has been added to form the next compound (here, sodium hydrogen oxalate, then disodium oxalate). Halfway between each equivalence point, at 7.5 mL and 22.5 mL, the pH observed was about 1.5 and 4, giving the pKa.

In monoprotic acids, the point halfway between the beginning of the curve (before any titrant has been added) and the equivalence point is significant: at that point, the concentrations of the two species (the acid and conjugate base) are equal. Therefore, the Henderson-Hasselbalch equation can be solved in this manner:



Therefore, one can easily find the pKa of the monoprotic acid by finding the pH of the point halfway between the beginning of the curve and the equivalence point, and solving the simplified equation. In the case of the sample curve, the Ka would be approximately 1.78×10^-5 from visual inspection (the actual Ka2 is 1.7×10^-5)

For polyprotic acids, calculating the acid dissociation constants is only marginally more difficult: the first acid dissociation constant can be calculated the same way as it would be calculated in a monoprotic acid. The second acid dissociation constant, however, is the point halfway between the first equivalence point and the second equivalence point (and so on for acids that release more than two protons, such as phosphoric acid).

Read more...

QUANTITATIVE ANALYSIS - Error, Accuracy and Precision

For more info; visit www.generalchemistryguide.com

The error of a measured value is the numerical difference between it and the true value.
E = Absolute error
T = True value
O = observed or measured value
E = O – T
The accuracy of a determination may be defined as the agreement between measured value and the true or most probable value. It is the value of E compared to T that is of importance. Errors are usually expressed in percent (E/T × 100). Smaller the error, greater the accuracy. Accuracy expresses the correctness of a measurement.

The precision of a measurement means the reproducibility or the extent of agreement of the individual values among themselves. The magnitude of the difference between individual values and the arithmetic mean of all is a measure of the precision.
The smaller the difference the greater the precision.

Read more...

Mean and Standard deviation -Arithmatic mean, s, Relative Standard Deviation (RSD), Coefficient of Variation (CV)

In analytical chemistry one of the most statistical terms employed is the standard deviation of a population of observations.

Consider a series of observations X1, X2, X3, ……………….Xn-1, Xn


The spread of the values is measured most efficiently by the standard deviation; s

 







Read more...

About This Blog

Lorem Ipsum

  © Blogger templates Shiny by Ourblogtemplates.com 2008

Back to TOP