Standardization of Non-standard solution, Chemistry tutorial

Introduction:

Two vital goals of chemical analysis in chemistry are:

1. The determination of what the constituents in a sample are and

2. The determination of how much of each constituent is in a sample.

The 1st is termed qualitative analysis and embraces essentially the detection of elemental ions and functional groups in a single compound; we will come across this in the next course. The 2nd is termed quantitative analysis, determines the amount of the elements or compounds present in a sample. There are definite techniques useful for this type of analysis, such as gravimetric analysis, volumetric analysis and instrumental analysis.

Volumetric processes of analysis are comprised of techniques in that the volume of a gas is determined or the volume of a solution of recognized concentration is measured. The latter technique is recognized as titrimetic analysis. Titrimetic analysis consists of titration. Titration is the procedure via which a solution of known concentration is added to another solution until the chemical reaction between the 2 solutes is complete.

The solution whose concentration is known as the standard solution or the titrant. In titration the standard solution (titrant) is gradually added from a burette to a solution that encloses a known mass of solute. The latter solution is usually referred to as the unknown. The point at which stoichiometrically equivalent quantities of materials have been brought together is recognized as the equivalence point of titration. This point signifies when the reaction must contain come to completion.

Therefore, in the process of titrating an unknown solution by a standard solution, there must be several ways to determine whenever the equivalence point of the titration has been readied. This is since, in most of the cases, the 2 solutions being titrated against each other are colorless and determining the equivalence point will be a guessing game and thereby introduce error. To avert this inappropriateness, the end of the equivalence point is usually done by using what is called an indicator.

The indicator in most cases is an organic substance (a dye) which changes color or gives some other obvious physiochemical change in the titrated solution when an amount of the standard solution (titrant) added is equivalent, or with negligible error, almost equivalent to the sample. The moment at which the indicator changes the color of the titrated solution is termed the end point of the titration. In other terms- the experimentally determined position of the stoichiometric point of the titration is termed the end point, whilst the actual theoretical stoichiometric point of the titration is the equivalence point. A successful titration, a titration through good quantitative consequences, is possible only if the end point as established via the indicator take places at the equivalence point of the titration. There are 4 kinds of titrimetic analysis each of which is depend on certain kind of reaction for example:-

a) Acid - Base titration

b) Oxidation - Reduction Titration

c) Precipitation Titration

d) Complexometric Titration

And there are 3 different techniques for determination of an unknown sample in titrimetic analysis.

a) Direct titration - Whenever standard solution directly reacts with the substance being determined.

b) Back titration - An accurately computed amount of 1 standard solution is added to the sample, in excess of the stoichiometric amount of the substance to be determined in the sample, and then unreacted excess is titrated through another standard solution. For this sort of titration 2 standard solutions are required.

c) Indirect titration - Whenever standard solution doesn't react by the substance being determined, but instead through another substance that is generated in the sample in equivalent amount of the determined substance.

In the subsequent chapter; we shall concentrate on acid - base titration and on oxidation -reduction titration. The others -precipitation and complexometric titrations shall be taken up in the higher course during the course of your programme.

In addition to the three techniques, we shall concern ourselves with direct and back titrations. I shall defer the treatment of indirect titration to next higher course in practical chemistry.

Acid - base Titration:

Introduction

Acid base titration entails a neutralization reaction. In this titration, acids are measured via titration through an appropriate base solution. On the other hand, bases are determined via titration through an appropriate acid solution. As we can remind from experiments 1 and 2 in the previous chapters, 2 kinds of solutions can be prepared from either the acidic or basic groups. Such as primary standards prepared from substance whose purity and properties have been exactly recognized and verified. The other group consists of substances whose purity can't be exactly known due to a number of causes: they are highly hygroscopic; they react with CO2   in air or moisture in air or since they are volatile. To find out the concentration of these solutions, the approximately computed concentrated solution is titrated against a primary standard solution.

Let us refresh our memory concerning standardization. Standardization generally involves:

Preparation of primary standard solution (titrant solution).

b) Preparation of a solution of approximate concentration to be standardized.

c) Titration - the procedure during which the exact concentration of solution is determined.

d) Calculation of the result from data obtained in titration .

In the titration method, one solution, generally the standard, is gradually added from the burette to another solution situated usually in a conical flask that encloses a recognized volume or known mass of solute, awaiting the chemical reaction between the reactants is complete.

Indicators:

The point at that stoichicometrically equivalent quantities of substances have been brought together is recognized as equivalence point of titration. To determine whenever the equivalence point of titration is attained, an indicator is utilized. In acid base titrations, the indicators these as methyl orange, phenolphthalein and bromothymol blue that are generally organic dyes that transform colour according to the hydrogen ion concentration of the solution or liquid to which they added or found themselves, are used.

For example, phenophtalein is an indicate or that is colorless in acidic solution but changed to pink in basic solution.

1334_methyle orange.jpg

Methyl orange (yellow form)

1839_Phenophtalein.jpg

Phenophtalein (Colourless form)

Such indicators can be regarded as weak acids of which either the undissociated molecule or the dissociated anion, or both are coloured. If we acquire methyl orange for instance,

HMe     →       H+           +             M

Red 4     ←   colourless                 yellow

Addition of acid moves the equilibrium to the left so that the solution becomes red. On the other hand, whenever alkali having hydroxyl ions are added, equilibrium moves to the right and the colour becomes yellow. It is as well said that it is hardly ever possible to discover an indicator that will indicate exact equivalence point, therefore the transform in colour we examine when we stop the titration indicates the end of the titration. Even though the end point of titration doesn't coincide by the equivalence point, the error is negligible.

The table below shows the list of indicators usually employed in acid -base titration, the colours they exhibit in diverse solutions and the pH range wherever colour changes occur.

Ph Scale:

The ph of a solution (or a liquid) is a compute of the hydrogen ion, H+, concentration in the solution. The ph scale ranges from 0 - 14. A solution by ph value less than 7 is acidic; exactly 7 is neutral, whilst the higher the value, the more alkaline the solution. Pure water has a ph of 7, it is neutral.

 When the colour of an indicator changes, it is due to the acidity of the change in the environment it has found itself. Near the end - point, both coloured forms will be present in appreciable quantities.

INDICATOR

                               COLOR in

Acid                          Neutral                                 ALKALI

pH

RANGE

1. Lithmus

Red

Purple

Blue

6.0 - 8.0 

2. Methylorange

Pink

Orange

Yellow

3.1 - 4.4

3.Phenolphtalein

Colorless

Colorless 

Pink

8.3 - 10.0

4.Bromothymol

Yellow

Blue

Blue

6.5 - 7.8

Choice of Indicator in Acid - base Titrations:

The choice of an indicator in an acid - base titration depends on the strengths (strong or weak) of the acid and base. The classification of the strengths of common acids and bases (alkali) is following in the table below:

Strength of Acid

  Strength of Base

STRONG 

WEAK

STRONG

WEAK

H2SO4 .HCl

HNO3

CH3COOH;H2C2

04;H2SO3   

NaOH;

KOH

Na2CO3;

aqNH3 

Kinds of Acid - Base titrations and the choice of indicator.

Acid -Base

Titration

Example

 pH of

solution end - point

Suitable

Indicator

Strong acid

vs strong base 

HCl vs NaOH

7

Any indicator:

Strong acid vs

weak base

H2SO4   vs

aqNH3

5 - 6

Methylorange

Weak vs

strong base

COON vs

NaOH

8 - 9

Phenolphthalein

Weak acid vs

weak base

CH3COOH vs

aq NH3

variable

 No suitable

Indicator

Wrong use of indicator:

The accuracy of a titration based on the utilize of the correct indicator. Wrong choice of indicator will lead to wrong consequence. For example, in the titration of a solution of a strong acid (HCl) with that of a weak base (Na2CO3), methyl orange should be used. If Phenolphthalein in used instead, the end point will appear when only half of the weak base, Na2CO3 has been used up, as shown in the following reactions.

HCl + Na2CO3      →   NaHCO3 + NaCl

H2SO4 + Na2CO3   →  NaHCO3    +  NaHSO4

The cause is that Phenolphthalein is sensitive to weak acid these as NaHCO3. The table can easily be illustrated. For example if we add say methyl orange indicator the colour we will observe is yellow. As we run in the acid from the burette, the neutralisation reaction occurs. The colour in the conical flask will still be orange until complete neutralisation whenever there is stoichiometric equivalence of the acid and base. The next drop of acid into the conical flask will not contain a base to react through. Hence there will be an extra drop of acid. At the point, the colour in the flask will now have to transform since the indicator now finds itself in a new environment for example acidic environment. In the particular case of methyl orange, the color will now change to pink thus marking the end point.

Problem Solving in Acid - base titration:

How can the amount of substance be determined in a titrimetric analysis? Can we remember how we calculated the concentration of the hydrochloric acid solution in experiment 2b. Go back and study the calculation. To fully comprehend the principle of the computation, a thorough knowledge of the following is needed:

  • We must be able to write a balanced chemical equation of the reaction.
  • From the balanced chemical equation, we must derive the stoichiometric mole equivalent ratio of the reactants.
  • From the above, we must be able to utilize the full applications of the mole concept to arrive at the molar concentration terms.

Writing balanced chemical equation:

The ability to write a balanced chemical equation is obtained from having the full knowledge of the formula (e) of elements, molecules and compounds as a consequence of their respective valence. This ability, I believe we must have acquired before embarking on this programme. If we are though unsure of ourselves, I will refer you to any standard textbook in chemistry at the senior secondary school level. Put briefly, the coefficients of the molecules, compounds taking part in the neutralisation reaction must be accurately find out so that

i) Matter will neither have been produced or destroyed

ii) We will have no inequality in the mass balance between the reactants and products.

Calculation of molar concentration:

If we remember what I said whenever we computed the accurate concentration of the hydrochloric acid in experiment 2B. There, no explanation was given to the procedure utilized in arriving at the solution. The formula or process of calculation we employed in that particular problem should be well known to you from your SSCE experience. Even though it is still acceptable and correct to utilize this process, it doesn't provide a full understanding of how to solve quantitative acid- base problems in chemistry.

Let us now go through the proper method for calculation using the mole perception.

  • Can we remember the definition of a mole?
  • Can we remember the formula derivation of the mole for a solid substance?
  • What about the formula derivation of the mole for a solution?

The answers to such posers can be originated in chapter 2. Put in brief, a mole is the amount of the substance equivalent to the Avogadro number in amount of the substance. It is as well the mass of the substance separated via the formula mass of the substance. But let us reveal this mole concept via solving appropriately the problem encountered in experiment 2B. The facts for that experiment as presented are:

i. A standard solution of sodium carbonate was supplied by the following characteristics. The concentration of sodium carbonate was given as 0.5moledm-3(molar).The volume of sodium carbonate employed for the titration was 25.0cm3.

ii. The actual concentration of the HCl was not known but the volume of the acid consumed in the titration was found to be 24.40cm3.

Now let us discover the concentration of the HCl. The following steps might be followed.

i. First write a balanced equation for the reaction

2HCl + Na2CO3   →    2NaCl + CO2   +H2O

ii. Then state the stoichiometri F mole ratio 2 moles HCl react through 1 mole Na2CO3   

iii. Determine the amount in moles of the reactant whose volume and concentration are known. In this case, it is the Na2CO3 solution

Amount in moles   =  Mass/ Molar mass    or

Amount in mole = Concentration in moldm-3 x volume cm3 /1000cm3

= 0.5 x 25 /1000

= 0.0125 moles

This amount in moles we have determined is contained in the 25cm3 of Na2CO3 that entirely reacted through the acid.

iv. Determine the equivalent amount in moles of the acid with the Na2CO3.

Since from the stoichiometric balanced chemical equation, 1 mole Na2CO3 = 2moles HCl, hence 0.0125 moles Na2CO3 will, have reacted with 0.025 moles HCl.

(2 x 0.0125 moles).

This amount in moles of the acid we have determined we should remind ourselves is contained in the 24.40cm3 of acid that reacted with the base.

v. Finally we will determine the molar concentration of the acid Molar concentration = amount in mole x 1000cm 3 

 Volume = 0.025 x1000 

24.40

=1.03moldm-3  

This is exactly similar answer we got earlier. We will examine that this process is longer than the earlier one but we will as well agree it provides a better understanding of the calculation method in volumetric analysis. I recommend this process of calculation from now on.

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