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Chem 18.1 The Common Ion Effect of 4

Expt. #5: The Common Ion Effect EG1, Group # 5, Prof. Julius Andrew NuñezOrante, Kristian Leonard O. & Eleazar, Sara Nicole H. 26 April 2012

Abstract

Common ion effect is a special case of the application of the law of chemical equilibrium to ionizationreactions. It is the effect on the equilibrium of the system done by the addition of a common ion. In thisexperiment, common ion effect was observed in different ionizations of acids/bases, buffering actions and thesolubility of slightly soluble solid substances. The experiment started with the comparison of the pH values ofsolutions of strong and weak acids and bases with their similar solutions mixed with a strong electrolytecontaining the common ion. Next was the identification of buffer solutions from sets of different mixtures of acidsand bases by measuring the pH change when added with acid or base. Finally, the solubility of benzoic acid in asodium benzoate was calculated based from the volume of base titrated. Then, this was compared to thesolubility of benzoic acid in water.

Keywords: acids, bases, buffers, common ion effect, pH, solubility

Introduction

Common ion effect is the shift in equilibriumcaused by the addition of a compound having anion in common with the dissolved substance. Thisphenomenon is actually an application of LeChatelier's Principle. When a given ion is added toa solution in equilibrium which already contains thation, the system shifts away from producing more ofthe common ion. (Chang, 2010) Because of thesystem’s shift to undo the stress caused by theaddition of the common ion, the system reduces theextent of ionization and the solubility of the addedsubstance.

An application of the common ion effect isthe use of buffers, solutions which lessen abrupt pHchanges after the addition of an acid or base. Acid-base buffers most often consist of conjugate acid-base pairs of a weak acid or base, meaning thatthey contain an acidic and a basic componentwhich do not neutralize each other. This is the keyfeature of a buffer solution: small amounts ofadditional OH – and H+ ions in the solution areconsumed by one buffer component and convertedinto the other, thus resulting in a negligible changein pH. (Silberberg, 2010)

The common ion effect has a number ofchemical, biological and industrial uses. A practicalapplication is in the treatment of water to reduce its“hardness” or high mineral content. For example,the addition of sodium carbonate to hard waterreduces the solubility of dissolved calciumcarbonate and causes it to precipitate out from thewater.

Another application of the common ioneffect is the aforementioned use of buffers. In the

laboratory and/or industry, buffers are used inanalytical procedures, calibration of pH meters,preparation and maintenance of dosage formsapproaching isotonicity, and many more. (U.S.Pharmacopeia, n.d.) Many enzymes work onlyunder a specific pH range that biological buffersmaintain. Another important example is themaintenance of blood plasma’s pH level of 7.4 byseveral buffer systems, such as bicarbonate(HCO3

–) and carbonic acid (H2CO3). (Chang, 2010)

The objectives of the experiment are: todetermine the effect of the presence of a common

ion on the extent of ionization of an electrolyte; toidentify buffer solutions and what type of substancethey consist of; and to determine the effect of acommon ion on the solubility of a slightly solublesubstance.

Experimental

The experiment was divided into threeparts:

A. Effect of Common Ion on the Ionization of Acidsand Bases

In this part of experiment, six solutionswere prepared as specified in table 1. Their pH wasthen determined by the use of pH papers. This wasdone to compare each solution regarding the effectof a common ion on the ionization of acids andbases.

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Table 1. Solutions to be tested for Part A.

SOL’N REAGENTS

A 10 mL 0.1 M HCl + 2 mL H2O

B 10 mL 0.1 M HCl + 2 mL 0.1 M NaCl

C 10 mL 0.1 M HOAc + 2 mL H2O

D10 mL 0.1 M HOAc + 2 mL 0.1 M

NaOAc

E 10 mL 0.1 M NaOH + 2 mL H2O

F10 mL 0.1 M NaOH + 2 mL 0.1 M

NaCl

B. Buffering Effect

In the second part, the pH of distilled water(labeled as S in table 4) was determined via pHpaper. Afterwards, 10 mL of distilled water wasplaced in 2 separate test tubes. A drop of 6 M HClwas added in one of the test tubes, and the pH ofthe resulting mixture was taken. In another testtube, a drop of 6 M NaOH and 10 mL of distilledwater were mixed. That mixture’s pH was alsodetermined.

Furthermore, the following solutions in table2 were prepared, and each solution’s pH wasdetermined. Once their pH values were known,

each solution was then transferred equally to 2separate test tubes. To one test tube, a drop of 6 MHCl was added. To the other one, a drop of 6 MNaOH was placed. Finally, the pH value of eachnew solution was measured. Additionally, thesolutions which exhibited buffering action wereidentified.

Table 2. Solutions to be tested for Part B.

SOL’N REAGENTS

A10 mL 0.5 M HOAc + 10 mL 0.5 M

NaOAc

B 10 mL 0.5 M HCl + 10 mL 0.5 M NaCl

C10 mL 0.5 M HNO3 + 10 mL 0.5 M

NaNO3

D10 mL 0.5 M NaH2PO4 + 10 mL 0.5 M

Na2HPO4

E 10 mL 0.5 M NH4OH + 10 mL 0.5 NH4Cl

C. Effect of Common Ion on the Solubility ofSlightly Soluble Salts

For the final part of the experiment, 50 mLof distilled water was placed in a 100-mL beaker.0.5 g of sodium benzoate was added to the distilled

water, and the resulting mixture was heated toapproximately 40°C. Benzoic acid crystals wereadded a small amount at a time, stirring after eachaddition, until the crystals no longer dissolved. Themixture was cooled with stirring to roomtemperature then filtered into a clean beaker. 10 mLof the collected filtrate in the beaker was pipetedand transferred into a 125-mL Erlenmeyer flask.Two drops of an indicator, phenolphthalein, wasadded to the Erlenmeyer flask, and the solution wastitrated with 0.01 M NaOH to a light pink endpoint.This part was executed to compare the solubility ofbenzoic acid in water and the solubility of benzoic

acid in sodium benzoate solution.

Results

A. Effect of Common Ion on the Ionization of Acidsand Bases

Table 3. Observed pH of prepared solutions in Part A.

SOL’N REAGENTS Obs. pH

A10 mL 0.1 M HCl + 2 mL

H2O1

B10 mL 0.1 M HCl + 2 mL 0.1

M NaCl

1

C10 mL 0.1 M HOAc + 2 mL

H2O3

D10 mL 0.1 M HOAc + 2 mL

0.1 M NaOAc5

E10 mL 0.1 M NaOH + 2 mL

H2O13

F10 mL 0.1 M NaOH + 2 mL

0.1 M NaCl13

B. Buffering Effect

Table 4. Observed results of Part B.

SOL’N pH

(orig)pH

(+ HCl)pH

(+ NaOH)EXP

CONCL.THEOR.CONCL.

S 5 2 11 -- --

A 5 5 5 B B

B 1 1 1 B NB

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C 1 1 1 B NB

D 8 7 8 B B

E 8 8 9 B B

Legend:S = distilled water

B = bufferNB = non-buffer

C. Effect of Common Ion on the Solubility ofSlightly Soluble Salts

After the titration of the filtrate with 0.01 MNaOH, a volume of 0.30 mL titrant was recorded.After computing, the solubility of benzoic acid insodium benzoate was found to be 3.0 x 10

-4M. In

the previous experiment, the observed solubility ofbenzoic acid solution was 7.6 x 10

-3M. The

calculations of these results are explained in detail

in the discussion part of the report.

Discussion

The common ion effect tends to suppressthe ionization of a weak acid or a weak base. According to Le Châtelier’s principle, the addition of the common ions from the strong electrolyte (salt)to a solution of weak acid will suppress theionization/dissociation of that weak acid (that is,shift the equilibrium from right to left), therebydecreasing the hydrogen ion concentration. Thus, asolution containing both the weak acid and thestrong electrolyte will be less acidic than a solutioncontaining only the weak acid at the sameconcentration. (Chang, 2010)

For solution A, the pH is 1 due to the H+

ions from HCl, which is known to be a strong acid.In solving for the pH, the equation used was:

M1V1 = M2V2 (0.1 M) (10 mL) = M2 (12 mL)

M2 = 0.0833 M

pH = -log[0.0833 M]= 1.08

For solution B, the pH is 1 because theaddition of NaCl, which contains a common ion withHCl (the Cl

–ion), did not much affect the

concentration of H+

ions since HCl is a strong acidthat has already completely dissociated.

For solution C and D, the acid used, whichwas HOAc, is a weak acid. The pH of D is higherthan the pH of C because the addition of OAc

–ion

from NaOAc, a strong electrolyte, caused thereaction to shift backwards. Therefore, theconcentration of H

+ions was decreased and the pH

was increased.

For solution E, the pH is 13 due to the OH –

ions from NaOH, which is known to be a strongbase. In solving for the pH, the equation used was:

M1V1 = M2V2 (0.1 M) (10 mL) = M2 (12 mL)

M2 = 0.0833 M

pOH = -log[0.0833 M] = 1.08pH = 14 – 1.08

= 12.92

For solution F, the pH is 13 because theaddition of NaCl, which contains a common ion withNaOH (the Na

+ion), did not much affect the

concentration of OH –

ions because NaOH is a

strong base.

A buffer solution is a solution of (1) a weakacid or a weak base and (2) its salt; bothcomponents must be present. The solution has theability to resist changes in pH upon the addition ofsmall amounts of either acid or base. A buffersolution must contain a relatively largeconcentration of acid to react with any OH

–ions that

are added to it, and it must contain a similarconcentration of base to react with any added H

+

ions. Furthermore, the acid and the basecomponents of the buffer must not consume eachother in a neutralization reaction. (Chang, 2010)

A buffer has an effective range of pKa ± 1pH unit. (Silberberg, 2010) Therefore, Solutions A,B, C, D and E were considered as buffersexperimentally. However, theoretically, Solutions Band C should not be considered as buffers becausethey contain strong acids (HCl and HNO3); basedfrom the definition, buffers consist of a weakacid/base and a salt containing its conjugatebase/acid. Meanwhile, Solution A has HOAc as theweak acid and OAc

–as its conjugate base. Solution

D has H2PO4 –

as the weak acid and HPO42-

as itsconjugate base. Solution E has NH4OH as the weak

base and NH4+ as its conjugate acid. Therefore,solutions A, D and E are buffer solutions. Thisdiscrepancy in results might have been because ofcommon human errors like inaccuratemeasurements and contaminated reagents and pHpapers. The use of pH paper was also impractical infinding the accurate pH of a certain solution.

In the last part of experiment, the solubilityof benzoic acid in sodium benzoate solution was

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obtained. This solubility of benzoic acid depends onthe amount of H3O

+ions dissolved in the sodium

benzoate solution. That is, [H3O+] is equal to the

solubility of benzoic acid. Since titration was doneto neutralize the H3O

+ions with OH

–ions from 0.01

M NaOH, the concentration of H3O+

is equal to theconcentration of OH

–.

M1V1 = M2V2 (0.01 M) (0.30 mL) = M2 (10 mL)

M2 = 3.0 x10-4

M

Equations at Equilibrium:

NaC6H5CO2 (s) Na+

(aq) + C6H5COO –

C6H5COOH (s) H+

+ C6H5COO –

It is clearly shown above that the tworeactants have a common ion, which is thebenzoate ion. This common ion caused an effect inthe solubility of benzoic acid in water. The solubility

of benzoic acid in sodium benzoate solution is lowerthan the solubility of benzoic acid in water ascomputed from the results in experiment IV. Thiscan be explained by the relationship betweencommon ion effect and solubility. The addition ofbenzoate ions from the sodium benzoate causedthe reaction to shift from right to left according to LeChâtelier’s Principle. This increase in [C6H5COO

–]

made the ion product greater than the solubilityproduct. To reestablish equilibrium, some benzoicacid crystals precipitated out of the solution until theion product was again equal to the solubilityproduct. The effect of adding a common ion

(NaC6H5CO2), then, is a decrease in the solubility ofthe slightly soluble salt (C6H5COOH) in solution.

Conclusions and Recommendations

Common ion effect is the shift in equilibriumcaused by the addition of a compound having anion in common with the dissolved substance.(Chang, 2010) It affects the pH of a solutioncontaining a weak acid or weak base bysuppressing the ionization of that weak acid orweak base. It follows the Le Chatelier’s Principleinvolving the amount of reactants/products added tothe solution. The reversibility of the reaction was

made possible by the presence of weakelectrolytes. In this case, the addition of thecommon ions from the strong electrolyte to asolution of weak acid/base will suppress theionization/dissociation of that weak acid/base. Thiswould shift the equilibrium from right to left. So,there would be a decrease in the concentration ofhydrogen/hydroxide ions. Thus, the pH of thesolution containing the strong electrolyte would beless acidic/basic than the pH of the solution

containing the weak acid/base only at sameconcentration.

Moreover, common ion effect is present inbuffer solutions. A buffer is a solution consisting ofconjugate acid-base pair of a weak acid/base. Inorder for the buffer to resist pH changes, it mustcontain high concentrations of the said acidic and

basic components. This is because when smallamounts of OH

–or H3O

+ions are added, a small

amount of one buffer component becomesconverted to the other component. However, aslong as the added amount of OH

–or H3O

+is much

smaller than that of the original conjugate acid-basepair present in the buffer, there is only a smallchange in the relative concentrations of the buffercomponents, and thus, only a negligible change inthe pH of the solution. (Silberberg, 2010)

Furthermore, common ion effect has agreat effect in the solubility of slightly soluble solids.

The increase in the concentration of the commonion will cause the reaction to shift from left to right.Therefore, more of the slightly soluble solidsubstance would precipitate out of the solutionthereby consuming the hydrogen ions. This wouldthen lessen the solubility of that solid in the solution.

By the end of the experiment, all objectiveswere achieved; thus, the experiment can beconsidered a success. However, someimprovements can be made in order to refine thisparticular study and to ensure the success of futureexperiments similar to this one. The use of a pH

meter instead of pH papers is recommended formore accurate pH readings. Maintaininguncontaminated reagents and clean glassware isalso crucial to the success of this experiment.

References

Chang, R. (2010). Chemistry (10th

ed.). New York,NY: McGraw-Hill.

U.S. Pharmacopeia. (n.d.). Buffer Solutions .Retrieved 25 April 2012, fromhttp://www.pharmacopeia.cn/v29240/usp29nf24s0_ ris1s119.html

Silberberg, M. S. (2007). Principles of General Chemistry (1

sted.). New York, NY: McGraw-Hill

I hereby certify that I have given substantialcontribution to this report.

______________________ __________________ Kristian Leonard O. Orante Sara Nicole H. Eleazar

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