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COMMON ION EFFECT
- 1. Course Title: COMMON ION EFFECT Course Code: CHM-404 Semmester: 4th E
- 2. Common ion effect and industrial applications
- 3. The reduction of the degree of dissociation of a salt by the addition of a common-ion is called the common-ion effect. Example In a saturated solution of silver chloride, we have the equilibrium AgCl (s) Ag+ (aq) + Cl– (aq) When sodium chloride is added to the solution, the concentration of Cl– ions will increase. The equilibrium shown above will be shifted to the left to form more of solid Ag Cl. Thus the solubility of AgCl, a typical sparingly soluble salt, will decrease. Common ion effect
- 4. Factors effecting degree of dissociation Nature of solute Nature of solvent Concentration Temperature common ion effect
- 5. Example Dissociation of hydrogen sulphide in presence of hydrochloric acid H2S 2H+ + S2- By applying the law of mass action, we have Ka = [H+]2 [S2-]/ [H2S] To the above solution of H2S , if we add hydrochloric acid, then it ionizes completely as HCl H+ + Cl- common ion effect
- 6. Solubility equilibria and solubility product When an ionic solid substance dissolves in water, it dissociates to give separate cations and anions. For example, for a sparingly soluble salt, say Ag Cl, we can write the equilibrium equations as follows : AgCl(S) Ag+ + Cl- common ion effect
- 7. Solubility equilibria and solubility product According to law of mass action K= [Ag+] [Cl–] / [AgCl] The amount of Ag Cl in contact with saturated solution does not change with time and the factor [Ag Cl] remains the same. common ion effect
- 8. Solubility equilibria and solubility product As [AgCl] is constant so equilibrium expression is Ksp= [Ag+] [Cl–] where [Ag+] and [Cl– ] are expressed in mol/L The equilibrium constant in the new context is called the solubility product constant (or simply the solubility product) and is denoted by Ksp. common ion effect
- 9. Solubility of barium iodate in presence of barium nitrate Barium iodate, Ba(IO3)2, has a solubility product Ksp = [Ba2+][IO3 -]2 = 1.57 x 10-9 Its solubility in pure water is 7.32 x 10-4 M. common ion effect
- 10. Common ion effect with weak acids and bases common ion effect
- 11. Common ion effect with weak acids and bases common ion effect
- 12. Activity Calculate the solubility of silver chromate, Ag2CrO4, in a 0.100 M solution of AgNO3. (Ksp for Ag2CrO4 = 9.0 × 10– 12) Activity Calculate the Ksp for Bismuth sulphide (Bi2S3), which has a solubility of 1.0 × 10– 15 mol/L at 25°C. Activity The solubility of BaSO4 is 2.33 × 10– 4 g/ml at 20°C. Calculate the solubility product of BaSO4 assuming that the salt is completely ionised. common ion effect
- 13. Common ion effect on solubility Adding a common ion decreases solubility, as the reaction shifts toward the left to relieve the stress of the excess product. Adding a common ion to a dissociation reaction causes the equilibrium to shift left, toward the reactants, causing precipitation. Applications of common ion effect
- 14. Common ion effect on solubility Example AgCl(s) Ag+ + Cl– If S be the solubility of AgCl, we have Ksp = [S mol/l Ag+] [S mol/l Cl–] Suppose 0.25 mol/L excess of HCl is added to the solution. Then ion product (Q) will be Q = [S mol/l Ag+] [(S + 0.25) mol/l Cl– ] If Q > Ksp Precipitation If Q = Ksp Saturated solution If Q < Ksp No precipitation Applications of common ion effect
- 15. Salting out of soap RCOONa(aq.) RCOO- (aq.) + Na+ (aq.) NaCl(aq.) Na+ (aq.) + Cl- (aq.) Applications of common ion effect
- 16. Purification of salt Applications of common ion effect NaCl(aq.) Na+ (aq.) + Cl- (aq.) HCl(aq.) H+ (aq.) + Cl- (aq.)
- 17. Lime Softening CO2 + Ca(OH)2 CaCO3 + H2O Ca(HCO3)2 + Ca(OH)2 2CaCO3 + 2H2O Mg(HCO3)2 + 2Ca(OH)2 2CaCO3 + 2H2O Applications of common ion effect
- 18. Harris, B. C. and W.H. Harris. 2010. Quantitative Chemical Analysis. 8th Edition. Freeman and Company New York. Christian, G.D. 2006. Analytical Chemistry. Sixth edition, John Wiley and Sons, New York. Skoog, D. A. and D.M. West 2005. Fundamentals of Analytical Chemistry. Hot Reinehart Inc., London. Kealey, D and P.J.Haines, 2002. Analytical Chemistry, Bios Scientific Publishers Limited, Oxford, UK. Sharma, B. K. 2004. Instrumental methods of chemical analysis, In; Introduction to Analytical chemistry: Goel Publishing House Meerut, 23th Edition. Reilley, C. 1993. Laboratory Manual of Analytical Chemistry. Allyn & Bacon, London. Reference books
Editor's Notes
- When a soluble salt (say A+C–) is added to a solution of another salt (A+B–) containing a common ion (A+), the dissociation of AB is suppressed. AB into A+ and B– . By the addition of the salt (AC), the concentration of A+ increases. Therefore, according to Le Chatelier’s principle, the equilibrium will shift to the left, thereby decreasing the concentration of A+ ions. Or that, the degree of dissociation of AB will be reduced. When solid NH4 Cl is added to NH4 OH solution, the equilibrium NH4 OH ----------NH4+ OH− shifts to the left. Thereby the equilibrium concentration of OH– decreases. This procedure of reducing the concentration of OH– ions is used in qualitative analysis
- Nature of Solute The nature of solute is the chief factor which determines its degree of dissociation in solution. Strong acids and strong bases, and the salts obtained by their interaction are almost completely dissociated in solution. On the other hand, weak acids and weak bases and their salts are feebly dissociated. (2) Nature of the solvent The nature of the solvent affects dissociation to a marked degree. It weakens the electrostatic forces of attraction between the two ions and separates them. This effect of the solvent is measured by its ‘dielectric constant’. The dielectric constant of a solvent may be defined as its capacity to weaken the force of attraction between the electrical charges immersed in that solvent. The dielectric constant of any solvent is evaluated considering that of vaccum as unity. It is 4.1 in case of ether, 25 in case of ethyl alcohol and 80 in case of water. The higher the value of the dielectric constant the greater is the dissociation of the electrolyte dissolved in it because the electrostatic forces vary inversely as the dielectric constant of the medium. Water, which has a high value of dielectric constant is, therefore, a strong dissociating solvent. The electrostatic forces of attraction between the ions are considerably weakened when electrolytes are dissolved in it and as a result, the ions begin to move freely and there is an increase in the conductance of the solution. (3) Concentration The extent of dissociation of an electrolyte is inversely proportional to the concentration of its solution. The less concentrated the solution, the greater will be the dissociation of the electrolyte. This is obviously due to the fact that in a dilute solution the ratio of solvent molecules to the solute molecules is large and the greater number of solvent molecules will separate more molecules of the solute into ions. (4) Temperature The dissociation of an electrolyte in solution also depends on temperature. The higher the temperature greater is the dissociation. At high temperature the increased molecular velocities overcome the forces of attraction between the ions and consequently the dissociation is great.
- Hydrogen sulphide (H2S) is a weak electrolyte. It is weakly ionized in its aqueous solution. There exists an equilibrium between unionized molecules and the ions in an aqueous medium as follows: H2S = 2H+ + S2- To the above solution of H2S , if we add hydrochloric acid, then it ionizes completely as HCl = H+ + Cl- This makes H+ a common ion and creates a common ion effect. Due to the increase in concentration of H+ ions, the equilibrium of dissociation of H2S shifts to the left and keeps the value of Ka constant. Thus the ionization of H2S is decreased. The concentration of unionized H2S is increased. As a result, the concentration of sulphide ions is decreased.
- At equilibrium the solute continues to dissolve at a rate that exactly matches the reverse process, the return of solute from the solution. Now the solution is said to be saturated. A Saturated solution is a solution in which the dissolved and undissolved solute are in equilibrium. A saturated solution represents the limit of a solute’s ability to dissolve in a given solvent. This is a measure of the “solubility” of the solute.
- The Solubility (S) of a substance in a solvent is the concentration in the saturated solution. Solubility of a solute may be represented in grams per 100 ml of solution. It can also be expressed in moles per litre. Molar Solubility is defined as the number of moles of the substance per one litre (l) of the solution. The value of solubility of a substance depends on the solvent and the temperature
- The value of K sp for a particular solubility equilibrium is constant at a given temperature. The product [Ag+] [Cl–] in the K sp expression above is also called the Ionic Product or Ion Product. The K sp expression may be stated as : the product of the concentration of ions (mol/l) in the saturated solution at a given temperature is constant. This is sometimes called the Solubility product principle.
- However in a solution that is 0.0200 M in barium nitrate, Ba(NO3)2, the increase in the common ion barium leads to a decrease in iodate ion concentration. The solubility is therefore reduced to 1.40 x 10-4 M, about five times smaller. .
- Adding a common ion prevents the weak acid or weak base from ionizing as much as it would without the added common ion. The common ion effect suppresses the ionization of a weak acid by adding more of an ion that is a product of this equilibrium.e.g.t he common ion effect of H3O+ on the ionization of acetic acid is shown above in figure. The common ion effect suppresses the ionization of a weak base by adding more of an ion that is a product of this equilibrium.
- Consider the common ion effect of OH- on the ionization of ammonia as shown in above figure. dding the common ion of hydroxide shifts the reaction towards the left to decrease the stress (in accordance with Le Châtelier's Principle), forming more reactants. This decreases the reaction quotient, because the reaction is being pushed towards the left to reach equilibrium. The equilibrium constant, Kb=1.8*10-5, does not change. The reaction is put out of balance, or equilibrium. Qa=[NH+4][OH−][NH3](23)(23)Qa=[NH4+][OH−][NH3] At first, when more hydroxide is added, the quotient is greater than the equilibrium constant. The reaction then shifts right, causing the denominator to increase, decreasing the reaction quotient and pulling towards equilibrium and causing Q to decrease towards K. An example of such an effect can be observed when acetic acid and sodium acetate are both dissolved in a given solution, generating acetate ions. However, sodium acetate completely dissociates but the acetic acid only partly ionizes. This is because acetic acid is a weak acid whereas sodium acetate is a strong electrolyte. As per Le Chatelier’s principle, the new acetate ions put forth by sodium acetate facilitate the suppression of the ionization of acetic acid, thereby shifting the equilibrium to the left. Since the dissociation of acetic acid is reduced, the pH of the solution is increased. Therefore, the common ion solution containing acetic acid and sodium acetate will have an increased pH and will, therefore, be less acidic when compared to an acetic acid solution. Thus, the common ion effect, its effect on the solubility of a salt in a solution, and its effect on the pH of a solution are discussed in this article https://byjus.com/chemistry/common-ion-effect/
- A practical example used very widely in areas drawing drinking water from chalk or limestone aquifers is the addition of sodium carbonate to the raw water to reduce the hardness of the water. In the water treatment process, highly soluble sodium carbonate salt is added to precipitate out sparingly soluble calcium carbonate. The very pure and finely divided precipitate of calcium carbonate that is generated is a valuable by-product used in the manufacture of toothpaste. The salting-out process used in the manufacture of soaps benefits from the common-ion effect. Soaps are sodium salts of fatty acids. Addition of sodi
- The solubility product (K sp) of an insoluble substance is the product of the concentrations of its ions at equilibrium. However, the ion product is the product of actual concentrations of ions which may or may not be in equilibrium with the solid. The increase in concentration of Cl– will shift the equilibrium to the left to form a precipitate of AgCl. Essential of physical chemistry By BS Bhal
- Since soaps are the sodium salts of carboxylic acids containing a long aliphatic chain (fatty acids), the common ion effect can be observed in the salting-out process which is used in the manufacturing of soaps. The soaps are precipitated out by adding sodium chloride to the soap solution in order to reduce its solubility. The important applications of Solubility product is Salting out of soap. Soap is sodium salt of higher fatty acid . It is precipitated from the solution by adding concentrated solution of NaCl( sodium chloride) . NaCl is strong electrolyte , it ionises completely in solution then concentration of sodium ions (Na+) increases. Due to common ion effect dissociation of soap decreases and soap gets precipitated and easily removed from soap solution.This is processs of getting solid soap from soap solution by adding salt like NaCl and called salting out of soap.
- The common ion effect is the phenomenon in which the addition of an ion common to two solutes causes precipitation or reduces ionization. You know Le Chatelier's principle, which states that a reaction stays in equilibrium unless acted on by an outside force; then the reaction will shift to accommodate the force and re-establish equilibrium. This means that if you want a reaction to shift, you just apply an outside force, such as temperature, concentration or pressure. In the case of the common ion effect, a reaction can be shifted by adding an ion that is common to both solutes, thus changing the concentration of the ion in solution and shifting the equilibrium of the reaction. The net effect of the common ion is that it reduces the solubility of the solute in the solution. The common ion effect can make insoluble substances more insoluble. An example of the common ion effect is when sodium chloride (NaCl) is added to a solution of HCl and water. The hydrochloric acid and water are in equilibrium, with the products being H3O+ and Cl- . Then, some sodium chloride is added to the solution. The NaCl dissolves into the solution, forming Na+ and Cl-. As the NaCl dissolves, the concentration of Cl- ions increases. The system accommodates by combining the Na+ and Cl- back into NaCl, which is a solid and precipitates out of the solution. In effect, more of the Cl- ions, which are common to both of the reactions, were added to the solution in equilibrium, so the equilibrium shifted back to the left. http://www.csun.edu/~ml727939/coursework/695/common%20ion%20effect/recrystallization%20of%20NaCl%20solution%20with%20HCl.htm
- As lime in the form of limewater is added to raw water, the pH is raised and the equilibrium of carbonate species in the water is shifted. Dissolved carbon dioxide (CO2) is changed into bicarbonate (HCO− 3) and then carbonate (CO2- 3). This action causes calcium carbonate to precipitate due to exceeding the solubility product. Additionally, magnesium can be precipitated as magnesium hydroxide in a double displacement reaction.[4] In the process both the calcium (and to an extent magnesium) in the raw water as well as the calcium added with the lime are precipitated. This is in contrast to ion exchange softening where sodium is exchanged for calcium and magnesium ions. In lime softening, there is a substantial reduction in total dissolved solids (TDS) whereas in ion exchange softening (sometimes referred to as zeolite softening), there is no significant change in the level of TDS. Lime softening can also be used to remove iron, manganese, radium and arsenic from water. https://en.wikipedia.org/wiki/Lime_softening https://www.suezwatertechnologies.com/handbook/chapter-07-precipitation-softening