LESSON 2

TEME: Volumetric analysis. Titration. Acid – base titration.

Medicobiological value: the titrimetric analysis is one of prime and accessible

expedients of reception of the chemical information. It is applied in clinical

biochemistry to diagnostics of series of pathological states. In biochemical,

physiological, sanitary - hygienic and etc. laboratories for definition of chemical

composition and quantitative content of separate builders of bodies and tissues,

study of a metabolism, metabolism of medicines, the definitions of composition of

water, ground, air and etc. will widely be used methods of analytical chemistry.

Diagnostics of the majority of diseases includes study of the clinical analyses which

are carried out with use of methods quantitative and qualitative analysis.

In medicobiological investigations the methods of acid-base titration will

widely be used, they allow to solve many problems incipient at chemical analysis of

biological fluids as at statement of the diagnosis, and at treatment of the patients (for

definition of an acidity of gastric contents, alkaline reserve of a blood and plasma).

In sanitary - hygienic practice the methods of acid-base titration allow to estimate

quality of various foodstuff.

The laboratory part of lesson a sectional theme is extremely important, as

during performance of laboratory operation the practical skills of operation with

analytical ware, technique of performance of measure analysis on an example of a

method of acid-base titration are shaped.

the formation of practical skills acid - base titration promotes development of

major operations on definition of substances in chemical, pharmaceutical,

biochemical, sanitary - hygienic and clinical practice.

To lesson it is necessary:

1. TO STUDY the following program questions: devices of a quantitative

analysis. The titrimetric analysis. A chemical equivalent of a substance. The

equivalent’s law. An equivalence point and expedients of its fixation. Expedients of

titration: forward, reverse, oblique. An acidimetry and alkalimetry: titrants, them

standardization; indicators.

The following program questions: substance and methods of acid-base

titration. Acid base indicators, mechanisms of change of colouring of indicators. A

choice of the indicator. The curves of acid-base titration. Application of a method of

neutralization in medical and sanitary - hygienic practice.

Literature:

U.Kask, J.D.Rawn. “General chemistry” P.139 – 144, 580 – 591.

Theoretical material to lesson

Acids and Bases

There are several definitions for an acid and a base. An Arrhenius acid is a

substance that produces H+ ions in water solution, and an Arrhenius base is a

substance that produces OH- ions in water solution. A Brønsted acid is a proton

donor, and a Brønsted base is a proton acceptor. A Brønsted acid-base reaction is a

proton transfer reaction.

A strong acid dissociates in water solution almost completely into H+ (aq) ions

and anions characteristic of the acid. A strong base is completely dissociated into

OH-(aq) ions and cations characteristic of the base. Common strong acids include

hydrochloric acid, HCl; nitric acid, HNO3; and sulfuric acid, H2SO4. The hydroxides

of the alkali metals and of the alkaline earth metals are strong bases.

Weak acids and bases ionize only to a small extent in water solution. Common

weak acids include acetic acid, HC2H3O2, and nitrous acid, HNO2. Phosphoric acid,

H3PO4, is a moderately weak acid. The most common weak base found in most

laboratories is aqueous ammonia, NH3(aq).

Acid-Base Titration

Acid-base reactions, or neutralization reactions, are commonly used to

determine the concentrations of acids (or bases) in solutions. If the concentration

and volume of one of the reactants in a neutralization reaction is known, the

concentration of the second solution can be determined if its volume is known. This

procedure requires the measurement of volumes, and is therefore called volumetric

analysis.

Let us how volumetric analysis is used to determine the concentration of an

acid solution using a base of known concentration. The base solution is added slowly

from a buret to a measured volume of the acid solution until all of the acid is

neutralized. The point at which neutralization occurs is usually detected by the

change in color of an organic dye such as litmus or phenolphthalein, which has one

color in acidic solution and a different color in basic solution. The organic dye used

for this purpose is called an indicator. This type of volumetric analysis is known as

titration. Figure 1 illustrates a setup for an acid-base titration.

Fig 1. A setup for an acid-base titration

The procedure we just described, in which the concentration of one solution

is determined by titration with another solution of known concentration, is called

standardization. In a standardization procedure, the solution whose concentration is

accurately known is called a standard solution.

Let us consider the standardization of a sodium hydroxide solution with a

standard hydrochloric acid solution. An accurately measured volume of the standard

hydrochloric acid solution is delivered with a pipet into a conical flask called an

Erlenmeyer flask (Fig. 3.). A few drops of an indicator solution are also added to the

flask. The sodium hydroxide solution is then slowly added from a buret into the acid

solution, with swirling, until just enough of the base is added to neutralize the acid.

This point is called the neutralization point or the equivalence point. The indicator

changes color at or very near the neutralization point and signals the endpoint for

the titration.

In the titration just described, the volume of the NaOH solution used in the

titration is measured as the difference between the final and initial buret readings of

the liquid level in the buret. From the volume and the molarity of the standard HCl

solution, we can calculate the number of moles of HCl present by use of the equation

(volume of HCl in liters)(M HCl) = number of moles HCl

The number of moles of HCl is also the number of moles of NaOH used in the

titration because HCl and NaOH react in a 1:1 mole ratio. The molarity of NaOH is

the number of moles of NaOH divided by the volume of NaOH solution in liters:

M NaOH = litersin solution NaOH volume

NaOH moles ofnumber

Example 1.

A 25,00-mL sample of 0,1208 M standard HCl solution is titrated with

28,52mL of NaOH solution to the endpoint. What is the molarity of the NaOH

solution?

Solution: We convert the volume of HCl in milliliters to liters by moving the

decimal point three places to the left to obtain 0,02500 L. The number of moles of

HCl is the volume of the solution (soln) in liters times the molarity:

0,02500 L soln )soln 1L

HCl mol 0,1208( = 3,020∙10-3 mol HCl

At the equivalence point of the titration, the number of moles of HCl equals

the number of moles of NaOH because HCl reacts with NaOH in 1:1 mole ratio, as

we see from the balanced equation for the titration reaction:

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

Therefore, the number of moles HCl = the number of moles NaOH = 3,020∙10-

3 mol. The volume of NaOH used in the titration is 28,52 mL or 0,02852L. The

molarity of NaOH is the number of moles of NaOH divided by its volume in liters:

М soln = soln NaOH L 0,02852

NaOH mol103,020 3 = 0,1059 mlo L-1 or 0,1059 M

The steps outlined above can be carried out on a calculator in a single

multistep operation:

soln L 0,02852

)HCl mol 1

NaOH mol 1)(

soln L 1

HCl mol 0,1208soln( L 0,02500

= 0,1059 mol L-1 or 0,1059 M.

In a titration, 50.00 mL of 0,1204 M hydrochloric acid requires 48,54 mL of

sodium hydroxide solution for neutralization. What is the molarity of the sodium

hydroxide solution?

In Example 1, the acid and base react in a 1:1 mole ratio. But a complete

neutralization of 1 mol of a polyprotic acid requires more than 1 mol of sodium

hydroxide. For example, the complete neutralization of 1 mol of sulfuric acid, a

diprotic acid, requires 2 mol of sodium hydroxide:

H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)

Similarly, the complete neutralization of 1 mol of phosphoric acid, a triprotic

acid, requires 3 mol of sodium hydroxide:

H3PO4(aq) + 3NaOH(aq) → Na3PO4(aq) + 3H2O(l)

When an acid-base titration is carried out using a standard solution of NaOH

to neutralize a polyprotic acid, the calculations resemble those outlined in Example

1. The only difference is in conversion of the number of moles of the base to the

number of moles of the acid. The conversion factor is the acid/base mole ratio given

by the balanced equation for the reaction, as shown by Example 2.

Example 2.

Sulfuric acid is sold by chemical supply houses in an approximately 18M

solution. In an experiment, the 18M stock solution is diluted to approximately 0,3M.

A 25,00-mL aliquot (or portion) of this diluted solution of sulfuric acid requires

32,58 mL of 0,5000M sodium hydroxide for a complete neutralization. What is the

molarity of the sulfuric acid solution?

Solution: We follow the general method outlined in Example 1. The only

difference is in the conversion of the number of moles of base to the number of moles

of acid. From the balanced equation for the reaction we see that 1 mol of the acid

reacts with 2 mol of the base:

H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)

Thus, the molarity of H2SO4 solution can be obtained by the following

multistep operation:

soln L 0,02500

)NaOH mol 2

SOH mol 1)(

soln L 1

NaOH mol 0,5000(soln L 0,03258 42

=

= 0,3258 mol H2SO4/L soln or 0,3258 M

Practice Problem 1.: (a) How many milliliters of 0,2056 M sodium hydroxide

is required to completely neutralize 25,00 mL of 0,1000 M phosphoric acid solution?

(b) How many milliliters of 0,600 M H3PO4 is completely neutralized by 37,0 g of

solid KOH?

Practice Problem 2.: Hydrochloric acid is an ingredient of gastric juice is

neutralized with 15,72 mL of 0,04209 M barium hydroxide solution. What is the

molarity of hydrochloric acid in gastric juice? How many milliliters of 0,06800M

barium hydroxide solution would be required to neutralize all the hydrochloric acid

in the same sample of gastric juice?

Volumetric analysis is a method in which strength of an unknown solution

(solution containing the substance, the amount of which is to be determined) is

determined with the help of a standard solution (solution containing a known amount

of a substance). Thus, we know the volumes of the two solutions and strength of the

standard solution, so the strength of the unknown solution can be calculated with the

help of the equation N1V1 = N2V2. This process is called titration. The end-point in

this process is seen with the help of an indicator, an another substance.

Volumetric analysis are of different types e.g. (i) acid –alkali, (ii) redox, (iii)

iodometry and iodimetry, (iv) precipitation titrations etc.

The following terms are used in volumetric calculations:

(i) Normal solution (N): A normal solution of a substance may be defined as,

«the solution which contains one equivalent mass of a substance per litre».

It is shown by symbol «N». Similarly a solution containing ½ equivalent of a

substance per litre is indicated as N/2. When the equivalents of the substance per

litre are 1/5, 1/10, 1/20, 1/100 etc. they are indicated as N/5, N/10, N/20, N/100 etc.

respectively. These factors indicate normality of the solution.

(ii) Standard solution: A standard solution is a solution of definite

concentration, the strength of which is known to us. The strength of the solution is

generally expressed in terms of normality.

(iii) Equivalent mass: Gram equivalent mass of a substance is that mass in

grams which is chemically equivalent to 1,008 g of hydrogen, 8 g of oxygen or 35,46

g of chlorine.

(a) Equivalent mass of an acid:

Equivalent mass of an acid is that mass which is obtained by dividing its

molecular mass by the number of replaceable hydrogen atoms present in one

molecule of acid.

i.e. Equivalent mass of an acid

molecule onein present atomshydrogen ereplaceabl of No.

acid of mass Mol.

i.e. Eq. mass of acid = Basicity

Mol.mass

Thus Eq. mass of H2SO4 = 98/2 = 49.

(b) Equivalent mass of base:

Equivalent mass of a base is that mass which neutralizes completely one

equivalent mass of the acid.

i.e. HCl + NaOH → NaCl + H2O

36,5 40

Here, one equivalent mass (36,5 g) of HCl reacts completely with 40 grams

of NaOH hence equivalent mass of NaOH is 40.

The basicity of some acids and acidity of some bases are given in the

following table:

Acid Basicity Base Acidity

HCl 1 NaOH 1

HNO3 1 KOH 1

CH3COOH 1 NH4OH 1

H2SO4 2 Ba(OH)2 2

H2C2O4∙2H2O 2 Ca(OH)2 2

H3PO3 2

H3PO4 3

Acid-base titrations: When the strength of an acid is determined with the

help of a standard solution of base, it is known as acidimetry. Similarly, when the

strength of a base (alkali) is determined with the help of a standard solution of an

acid, it is known as alkalimetry. Both these titrations involve neutralization of an

acid with an alkali. In these titrations H+ ions of the acid combine with OH- ions of

the alkali to form ionized molecules of water.

HA + BOH → BA + H2O

Acid Alkali Salt Water

or H+ + A- + B+ + OH- → B+ + A- + H2O

or H+ + OH- → H2O

The end point in these titrations is determined by the use of organic dyes

which are either weak acids or weak bases. These change their colours within a

limited range of hydrogen ion concentrations, i.e., pH of the solution.

Phenolphthalein is a suitable indicator in the titrations of strong alkalies (free from

carbonate) against strong acids or weak acids. Methyl orange is used as an indicator

in the titrations of strong acids against strong and weak alkalies. As no indicator

gives correct results in the titrations of weak acids against weak bases, such titrations

are performed by some other methods (physical methods).

Acid-base indicators

An acid-base indicator is an organic dye that signals the end-point by

a visual change in colour. Phenolphthalein and methyl orange are two common examples of acid-

base indicators. Phenolphthalein is pink in base solution and colourless in acid

solution. Thus when added to the acid solution in the receiver flask, it shows no

colour. As the added base is in slight excess, it becomes pink. Thus phenolphthalein

signals the end-point by a colour change from colourless to pink.. Similarly methyl

orange indicates the end-point by a colour change from red (in acid) to yellow (in

base).

pH range of indicators

Most indicators do not change colour at a particular pH. They do so over a

range of pH from two to three units. This is called the pH range which is different

for various indicators.

pH ranges of some acid-base indicators

pH curves and Indicator range

During an acid-base titration the pH of the solution in the receiver flask changes

with the addition of the titrant from the burette. A plot of pH against the volume of

the solution being added is known as pH curve or titration curve. For illustration,

the pH curve produced by titration of HC1 solution with NaOH solution is shown in

Fig. As NaOH is added, the pH of the solution increases slowly at first, then rapidly

in the vicinity of the equivalence point and again slowly. The equivalence point lies

in the middle of the vertical portion of the curve (pH = 7). It must be clearly

understood that equivalence point is the theoretical end-point of a titration. The end-

point of a titration determined by a colour change of the indicator in titration solution

is the experimental estimate of the equivalence point.

Fig.2. pH curve for titration of a strong Fig. 3 Indicator ranges for phenolphthalein,

base with strong acid. litmus, and methyl orange.

The titration curve in Fig.2 shows that it remains vertical around the

equivalence point. From a study of this part of the curve, it is evident that the volume

of litre used at the experimental end - point will be very nearly the same as for the

equivalence point provided that: the indicator used has a small pH range and the

range wholly falls on the vertical portion of the curve.

Thus, a suitable indicator for a given titration may be defined as one which

has as narrow a pH range as possible that lies entirely on the upright part of the

Indicator Colour change

(acid-base)

pH range

Methyl orange Red-orange 3.1-4.4

Methyl red Red-yellow 4.4-6.0

Litmus Red-blue 5.0-8.0

Bromothymol blu Yellow-blue 6.0-7.6

Phenolphthalein Colourless-pink 8.3-10.0

titration curve. For example, as shown in Fig.3, phenolphthalein, litmus and methyl

orange may be used as indicators for acid-base titrations.

Сhoice of a suitable indicator

The choice of a suitable indicator for a particular acid-base titration depends

on the nature of the acid and the base involved in the titration. We may have the

titration of:

(a) a strong acid with a strong base

(b) a weak acid with a strong base

(c) a strong acid with a weak base

(d) a weak acid with weak base

Which indicator is suitable for a given titration, can be found by examining

the titration curve of that titration. We have already discussed that a suitable indicator

is one which has a small pH range that falls wholly on the upright portion of the

titration curve.

All the pH curves given in Fig.6 refer to addition of 0.1 M monoacid base to

25 ml of 0.1 M of monobasic acid. The equivalence point in all cases is at pH 7

when all the acid has been neutralized by the base to form a salt. If the titration is

performed so that acid is added to the base, the pH curve is the mirror image of that

shown.

Titrating a strong acid with a strong base Figure 4 (a) depicts the titration curve when NaOH is added gradually to HC1.

It shows that the pH of the titration solution rises extremely slowly in the beginning.

In the vicinity of the equivalence point, the pH rises dramatically and the curve

becomes vertical. Beyond this, the curve becomes almost flat that shows a slight rise

of pH when only excess base is present in the titration solution.

Fig.4. To find a suitable indicator from a study of the pH curves for: (a) a strong acid

and strong base; (b) weak acid and strong base; (c) strong acid and weak base;

(d) a weak acid and weak base.

The vertical portion of the curve extends from pH 3 to pH 7. The pH ranges

of methyl orange (3.1 - 4.4), and phenolphthalein (8.30 - 10.0) are fairly narrow and

fall on the vertical curve. Thus both methyl orange and phenolphthalein are

suitable indicators for strong acid/strong base titrations. Litmus with an

exceptionally wide pH range (4.5-8.3) is seldom used. Its colour does no; change

sharply from red to blue but goes through various shades of purple. Titrating a

Weak acid with a Strong base Figure 6 (b) represents the titration curve when NaOH (strong base) is added

to acetic acid (weak acid). The pH curve rises slowly in the beginning but near the

equilibrium point, the pH changes from 6 to 11 and the curve becomes vertical.

Beyond this the shape of the titration curve is similar to that for strong acid/strong

base.

Phenolphthalein has pH range 8.3-10.0 that falls on the vertical part of the

titration curve as marked in the figure. The pH range of methyl orange (3.1-4.4), on

the other hand, does not fall on the vertical curve. Thus if methyl orange is used as

indicator, the experimental end-point will be reached earlier than the equivalence

point. Therefore for weak acid-strong base titration phenolphthalein is a suitable

indicator, while methyl orange is not.

Titrating a strong acid with a weak base

The titration curve for HC1 (strong acid) with NH4OH (weak base) is shown in

Fig.4 (c). As NH4OH is added, the pH of the titration solution increases gradually.

Around the equivalence point, a sharp rise in pH occurs approximately from 3 to 8,

when the curve becomes vertical. The pH range of methyl orange (3.1-4.4) and that

of methyl red (4.4-6.0) falls on the vertical portion of the titration curve. Evidently,

methyl orange and methyl red are suitable indicators for strong acid/weak base

titrations.

Titrating a weak acid with a weak base The titration curve for acetic acid (weak acid) with NH4OH (weak base) is

shown in Fig.4 (d). The pH of the titration solution rises gradually and there is no

sharp change in pH around the equivalence point. The vertical portion is missing in

the titration curve. Under these conditions, all indicators change colour only

gradually and no indicator is suitable.

THEORIES OF ACID-BASE INDICATORS

An acid-base indicator is an organic substance used for the detection of

equivalence point or neutral point in an acid-base titration. An indicator has one

colour in acid solution and entirely different in basic solution. The end-point of the

titration is shown by a colour change of the indicator. Two theories have been put

forward to explain the indicator action in acid-base titrations:

(1) The Ostwald's theory

(2) The Quinonoid theory

We will discuss these with reference to two commonly used indicators,

namely, methyl orange and phenolphthalein. The Ostwald's theory According to

this theory:

(1) an acid-base indicator is a weak organic acid (HIn) or a weak organic base

(InOH), where the letter In stands for a complex organic group. Methyl orange

and phenolphthalein are both weak acids.

(2) the ionized indicator, HIn, has a colour different from the In- ions produced by

the ionization of the indicator in aqueous solution.

(3)the degree of ionization of the indicator determines the visible colour of the

indicator solution.

How an acid-base indicator works

Let us explain the indicator action by taking example of methyl orange. Methyl

orange is a weak acid and gives the following ionization equilibrium in solution.

HIn ↔ H+ + In-

red yellow

In accordance with the law of mass action,

[HIn]

]][In[HK in

where Kin is the dissociation constant of the indicator and is called the Indicator

constant. The anion In- is yellow and the unionized form HIn is red. If an acid is added

to the solution, the hydrogen ion concentration, [H+], in the equilibrium expression

(1) increases. To maintain Kiu constant, the equilibrium shifts to the left. Thereby the

concentration of [In-] is reduced and the concentration of [HIn] increases so that the

solution is red. On the other hand, upon addition of a base to the solution, H+ ions

are removed as H2O by reacting with OH- ions of the base. This shifts the equilibrium

to the right, resulting in the increase of In- ions that are yellow. Thus in acid solution

the unionized HIn molecules predominate and the solution is pink, while in basic

solution In- ions are in excess and the solution is yellow.

Relation of Indicator colour to pH

The indicator solution contains both the yellow In- and the red HIn molecules.

The actual colour shade of the indicator depends on the ratio of concentrations of In-

and HIn present in solution. From the equilibrium constant expression we can write

][

][

HIn

InKH in

If [H+] is large, the concentration of In- ions is also large and the colour is yellow.

When [H+] is small, [HIn] is large and the solution is red. At the equivalence point,

[In-] = [HIn] and the colour is orange (red + yellow). Obviously the indicator colour

is controlled by hydrogen ion concentration or pH of the solution.

Taking logarithms and using definition of pH and Kin, the expression can be

converted to the Henderson-Hasselbalch equation.

][

][log

HIn

InpKpH in

At the equivalence point, [In-] = [HIn] and methyl orange in solution is orange. Then,

inpKpH

The numerical value of the indicator constant Kin for methyl orange is 3.6 and the pH

of the orange solution is, therefore, about 4. As the values of Kin for the various

indicators are different, they will have intermediate intense colours (middle tint) at

different pH values. When a base is added to an acid solution in a titration, the

colour change of the indicator is gradual. It just becomes visible to the human eye

when [In-]/[HIn] = 1/10 and pH calculated from equation. The colour of the

indicator continues to change till [In-]/[HIn] = 10 when pH is 4.4. The pH range

between 3.1(red) and 4.4(yellow) is called the colour change interval of methyl

orange. The visible indicator colour change takes place between these pH values.

Indicator action of phenolphthalein

It can be explained as in case of methyl orange. It is a weak acid and exists

as the following equilibrium in solution,

HIn ↔ H+ + In- colourless pink

HIn molecules are colourless, while In- ions are pink. Thus in acid solution,

phenolphthalein is colourless and in basic solution it is pink. The value of Kin = 9.6

and the pH of the intermediate intense pink tint is also 9.6. The colour change

interval of phenolphthalein is 8.1 -10.0.

Quinonoid theory of Indicator colour change

The Ostwald's theory takes care of the quantitative aspect of indicator action

adequately. The Quinonoid theory, on the other hand, tells us the cause of colour

change of an indicator in acid-base solutions. It lays down that:

(1) the unionised HIn molecule and the anion In- are tautomeric forms of the

indicator which isan organic dye.

(2) one tautomeric form possesses the quinonoid structural unit and is called the

quinonoid form.

It has a deep colour. The other form has a lesser colouring group, say, -N=N- and

or simply benzene rings and is called the benzenoid form. This form has a light

colour or no colour.

(3) the colour change of the indicator occurs when one tautomeric form is

transformed into the other due to change of pH of the solution.

Let us illustrate the Quinonoid theory by taking example of methyl orange and

phenolphthalein.

Methyl Orange. The red quinonoid form of methyl orange exists in acid

solution. It is converted to yellow benzenoid form when pH alters to the basic side.

Phenolphthalein. Phenolphthalein exists in two tautomeric forms: (i) the

benzenoid form which is yellow and present in basic solution; and (ii) the quinonoid

form which is pink and present in acid solution.

Research work:

"Preparation of a solution of a sodium tetraborate and

standartization of its concentration on НСl"

Experiment 1. Preparation of a reference solution of a sodium tetraborate

(50,0 mls of a solution of a sodium tetraborate, c(1/2 Na2B4O7) ~ 0,1 mol/l to within

three significant digits)

Performance of experiment. Before a cut-in familiarize with the device

analytical balance with the instructions on their operation. Prepare a measuring flask

and dry funnel. Calculate mass of a shot of a sodium tetraborate necessary for

preparation of a given solution. The designed shot of a sodium tetraborate weigh on

counter balance. As container use test tube. For record of results of weighing on

analytical balance in a laboratory magazine beforehand prepare under the shape the

following records:

Mass test tube with Na2B4O7 • 10H2O … … ….

Mass test tube with oddments of substance … … ….

———————————————————

Mass of Na2B4O7 • 10H2O … … … … ….

Then test tube with a shot weigh on analytical balance. On a scale pan at once

establish a load relevant to result of prestress weighing. Write down the indications

of analytical balance and, having taken off test tube from a scale pan, is cautious add

a shot through a dry funnel in a measuring flask. Test tube cant above slowly by

funnel and do not suppose a diffusion and pulverization of substance. Test tube with

oddments of substance on walls immediately again weigh on analytical balance and

write down result. After that take off from weights all load and close doors.

Add heat distilled water in a test tube. A flask charge with distilled water

approximately on 2/3 volumes, dissolve all substance, cool a solution up to ambient

temperature and add by distilled water up to a score. Intermix a solution in a flask.

A flask close by a fuse and paste a label with the name of a solution and surname.

Experiment 2. An establishment of precise concentration of a hydrochloric

acid on a reference solution of a sodium tetraborate

Performance of experiment. Prepare ware necessary for carrying out of the

titrimetric analysis. A burette, pipet and flasks wash by distilled water. A laboratory

magazine prepare for record of results of the analysis:

Volume of a solution Na2B4O7 - 5,00 mls

c(1/2Na2B4O7) = … mol/l

V(HCl): 1 titration …

2 « …

3 « …

V(HCl): medial value …

In a burette fill by a working solution HCl. Measuring pipet by containment 5

mls wash by a solution of a sodium tetraborate. Then in flasks for titration select by

the pipet till 5,00 mls of a solution Na2B4O7 and 1 drop of a solution of the indicator.

The solution becomes yellow. Assay in a flask titrate by a solution HCl before

transferring colouring in light pink. (Sometimes for rising accuracy of titration will

use a flask - "witness". The witness plot, measuring off in a flask 5,0mls of water, 1

drop of an acid and 1 drop of the indicator. Assay titrate before occurrence of the

same colouring, as in a flask - "witness"). The pink colouring in an endpoint of

titration means presence in a flask of excess of an acid about 1 drop. (Therefore in

separate experience it is necessary to spot volume of one drop of a solution for a

sectional burette). Having concluded titration, wash a flask usual, and then distilled

water also lead titration some more time before reception not less than four results

concurrent in limits 0,04 mls.

Flasks for titration and pipet wash by distilled water. Flasks with solutions

clean(remove) in a table or case for performance of other laboratory operations.

Processing results of experiment. From plurality of separate results of titration

reject misses and find medial of concurrent results. From medial value take away

volume of one drop of a sectional burette. On obtained volume of an acid calculate

concentration of HCl. Spot also error of the analysis.

Formulate a deduction describing basic result of operation: specify

concentration of a working solution HCl to within three significant digits.

Problems for discussion

1. In an acid – base titration, what is the meaning of each of the following words?

“Neutralization”, “equivalence point”, and “endpoint”.

2. How many milliliters of 0.200 M sodium hydroxide is required to neutralize

40.0ml of 0.0500 M HCl solution?

3. How many milliliters of 0.200 M sodium hydroxide is required to completely

neutralize:

(a) 40.0 ml of 0.0500 M H2SO4;

(b) 40.0 ml of 0.0500 M H3PO4.

4. Write down the mathematical form of equivalence’s law.

5. What demands are showed to reactions used in the titrimetric analysis?

6. What titrants apply in an acidimetry and alkalimetries?

7. What molarity and titre of a solution HCl, if on titration 25 mls of it at the presence

of phenolphthalein 19.75 mls 0,1 M of a solution NaOH are spent.

8. Calculate molar mass of an equivalent for the following substances: NaOH,

H3PO4, CaCl2

9. What the medium and the equivalence point is observed at titration:

а) of weak acids by the strong bases;

b) of weak bases by strong acids?

10. In what case the spring of titration is more - at titration 0,1M or 0,01М by a

solution НСl?

11. What substances can be determined in a solution by a neutralization method:

Nа2SО4, НСl, С6Н12О6, Nа2СО3, КNО3, Н2SО4?

12. At clinical investigations in particular cases determine gastric acidity - content

of a hydrochloric acid and total acid. The curve of titration has 2 springs of

titration, at pH - 3-5 and pH - 8-10.

Offer: а) the plan of performance of the analysis;

b) what indicators will use.

13. On titration 60 mls of a solution of a potassium hydroxide went 30 mls 0,1M

solution of Н2SО4. How many grammes of a potassium hydroxide in 200 mls of

such solution?

14. Calculate molarity of a solution (ρ= 1,18 g/sm3), containing 24,5% of sulfuric

acid.

15. Calculate a percent by mass (%) of a hydrochloric acid in a gastric juice

(ρ=1,01g/sm3), if 5 mls of juice are spent for titration 3,1 mls 0,098М solution of

sodium hydroxide.