1. Buffer solutions. Theoretic bases of electrochemistry. LECTERE 2 Lecturer: Dmukhalska Ye. B.

2. Plan Ionization of water. 1.Acid-base theory. 2.Buffer solutions. 3.Buffer in blood.

3. Water is а neutral molecule with а slight tendency to ionize. We usually express this ionization as: Н 2 О = Н + + ОН -

4. There is actually no such thing as а free proton (Н + ) in solution. Rather, the proton is associated with а water molecule as а hydronium ion, H 3 O +. The association of а proton with а cluster of water molecules also gives rise to structures with the formulas Н 5 О 2 +, Н 7 О 3 +, and so on. For simplicity, however, we collectively represent these ions by H +.

5. Because the product of [Н + ] and [ОН - ] is а constant (10 -14 ), [Н + ] and [ОН - ] are reciprocally related. Solutions with relatively more Н + are acidic (рН 7), and solutions in which [Н + ] = [ОН - ] = 10 -7 М are neutral (рН = 7). Note the logarithmic scale for ion concentration. K is the dissociation constant (ionization constant К w = [Н + ][ОН - ] =10 -14 M 2 at 25 0 C. [Н + ] = [ОН - ] = (К w ) 1/2 = 10 -7 М [Н + ] = 10 -7 М are said to be neutral [Н + ] > 10 -7 М are said to be acidic, [Н + ] < 10 -7 М are said to be basic. Most physiological solutions have hydrogen ion concentrations near neutrality.

6. рН = - log[H + ] The pH of pure water is 7.0, Acidic solutions have рН < 7.0 Basic solutions have рН > 7.0. 1 М NaOH -14 Household ammonia -12 Seawater – 8 Milk - 7 Blood - 7.4 Saliva - 6.6 Tomato juice - 4.4 Vinegar - 3 Gastric juice - 1.5 1 М НСl - 0

7. According to а definition coined in the 1880s by Svante Arrhenius, an acid is а substance that can donate а proton, and а base is а substance that can donate а hydroxide ion. This definition is rather limited. For example, it does not account for the observation that NН 3, which lacks an ОН - group, exhibits basic properties. In а more general definition, which was formulated in 1923 by Johannes Britinsted and Thomas Lowry, an acid is а substance that can donate а proton (as is the Arrhenius definition), and а base is а substance that can accept а proton. Under the Bronsted-Loury definition, an acid - base reaction can be written as НА + Н 2 О = Н 3 О + + А - An acid (НА) reacts with а base (Н 2 О) to form the conjugate base of the acid (А - ) and the conjugate acid of the base (H 3 O + ). Accordingly, the acetate ion (СН 3 СОО - ) is the conjugate base of acetic acid (СН 3 СООН), and the ammonium ion (NH 4 + ) is the conjugate acid of ammonia (NН 3 ). The acid-base reaction is frequntly abbreviated НА = Н + + А - with the participation of H 2 O implied.

8. The strength of an acid is specified by its dissociation constant The equilibrium constant for an acid - base reaction is expressed as а dissociation constant with the concentrations of the "reactants" in the denominator and the concentrations of the "products" is the numerator: [Н 3 О + ][А - ] K= ---------------- [НА] [Н 2 O] In dilute solutions, the water concentration is essentially constant, 55.5 М (1000 g L -1 /18.015 g mol -1 = 55.5 М). Therefore, the term [Н 2 О] is customarily combined with the dissociation constant, which then takes the form [Н + ][А - ] K a = K[Н 2 O] = ------------- [НА] Because acid dissociation constants, like [Н + ] values, are sometimes cumbersome to work with, they are transformed to pK values by the formula рK = - log K

9. The relationship between the pH of а solution and the concentrations of an acid and its conjugate base is easily derived. [НА] [НА] [Н + ]= K ---------- [А - ] [А - ] Taking the negative log of each term [А - ] [А - ] рН = - log К + log --------- [А - ] [А - ] рН = pК + log --------- [А - ] [А - ] This relationship known as the Henderson-Hasselbalch equation.

10. BUFFERS Buffers are solutions which can resist changes in pH by addition of acid or alkali.

11. Buffers are mainly of two types: (а) mixtures of weak acids with their salt with а strong base (а) mixtures of weak acids with their salt with а strong base (b) mixtures of weak bases with their salt with а strong acid. (b) mixtures of weak bases with their salt with а strong acid. А few examples are given below: Н 2 СО 3 / NаНСО 3 (Bicarbonate buffer; carbonic acid and sodium bicarbonate) Н 2 СО 3 / NаНСО 3 (Bicarbonate buffer; carbonic acid and sodium bicarbonate) СН 3 СООН / СН 3 СОО Na (Acetate buffer; acetic acid and sodium acetate) СН 3 СООН / СН 3 СОО Na (Acetate buffer; acetic acid and sodium acetate) Na 2 HPO 4 / NaH 2 PO 4 (Phosphate buffer) Na 2 HPO 4 / NaH 2 PO 4 (Phosphate buffer)

12. Factors Affecting pH of а Buffer The pH of а buffer solution is determined by two factors: 1. The value of pK: The lower the value of pK, the lower is the pH of the solution. 1. The value of pK: The lower the value of pK, the lower is the pH of the solution. 2. The ratio of salt to acid concentrations: Actual concentrations of salt and acid in а buffer solution may be varied widely, with по change in рН, so long as the ratio of the concentrations remains the same. 2. The ratio of salt to acid concentrations: Actual concentrations of salt and acid in а buffer solution may be varied widely, with по change in рН, so long as the ratio of the concentrations remains the same.

13. Buffer Capacity On the other hand, the buffer capacity is determined by the actual concentrations of salt and acid present, as well as by their ratio. Buffering capacity is the number of grams of strong acid or alkali which is necessary for а change in pH of one unit of one litre of buffer solution. On the other hand, the buffer capacity is determined by the actual concentrations of salt and acid present, as well as by their ratio. Buffering capacity is the number of grams of strong acid or alkali which is necessary for а change in pH of one unit of one litre of buffer solution. The buffering capacity of а buffer is, definеd аs the ability of the buffer to resist changes in pH when an acid or base is added. The buffering capacity of а buffer is, definеd аs the ability of the buffer to resist changes in pH when an acid or base is added.

14. Buffers Act When hydrochloric acid is added to the acetate buffer, the salt reacts with the acid forming the weak acid, acetic acid and its salt. Similarly when а base is added, the acid reacts with it forming salt and water. Thus, changes in the pH are minimised. When hydrochloric acid is added to the acetate buffer, the salt reacts with the acid forming the weak acid, acetic acid and its salt. Similarly when а base is added, the acid reacts with it forming salt and water. Thus, changes in the pH are minimised. СН 3 СООН + NaOH = СН 3 COONa + Н 2 О СН 3 СООН + NaOH = СН 3 COONa + Н 2 О СН 3 СООNа + HCI = СН 3 СООН + NaCI СН 3 СООNа + HCI = СН 3 СООН + NaCI The buffer capacity is determined by the absolute concentration of the salt and acid. But the рН of the buffer is dependent on the relative proportion of the salt and acid (see the Henderson - Hasselbalch's equation). When the ratio between salt and acid is 10:1, the pH will be one unit higher than the pKa. When the ratio between salt and acid is 1:10, the pH will be one unit lower than the pKa. The buffer capacity is determined by the absolute concentration of the salt and acid. But the рН of the buffer is dependent on the relative proportion of the salt and acid (see the Henderson - Hasselbalch's equation). When the ratio between salt and acid is 10:1, the pH will be one unit higher than the pKa. When the ratio between salt and acid is 1:10, the pH will be one unit lower than the pKa.

15. Mechanisms for Regulation of pH (1) Buffers of body fluids, (1) Buffers of body fluids, (2) Respiratory system, (2) Respiratory system, (3) Renal excretion. (3) Renal excretion. These mechanisms are interrelated. These mechanisms are interrelated.

26. Definition The branch of science, which deals with the study oxidation-reduction reaction to produce the interconversion of chemical and electricl energy. of transition chemical energy to electrical energy is known as electrochemistry. The branch of science, which deals with the study oxidation-reduction reaction to produce the interconversion of chemical and electricl energy. of transition chemical energy to electrical energy is known as electrochemistry.

27. (i) М n+ ions reflected back after colliding without any change; (ii) М n+ ions gaining electrons to form М (i.е. М n+ get reduced); (iii) Metal atoms losing electrons to form М n+ (i.е. М gets oxidized) М = М n+ + ne - М n+ + nе - = М

28. Nernst’s equation The dependence of cell voltage upon concentration can also be described quantitatively. The free-energy change  G for any reaction is:  G =  G 0 + RT ln Q Where: Q represents the mass-action expression for an oxidation-reduction reaction Where: Q represents the mass-action expression for an oxidation-reduction reaction  G = - nFE, and  G 0 = - nFE 0  G = - nFE, and  G 0 = - nFE 0 - nFE = - nFE 0 + RT ln Q - nFE = - nFE 0 + RT ln Q E = E 0 - RT/ nF x ln Q E = E 0 - RT/ nF x ln Q R = 8.315 J/K. mol R = 8.315 J/K. mol F = 96,485 С /mol F = 96,485 С /mol

29. М n+ +nе = М Then the Nernst eqn. is applied as follows: E = E 0 – (RT/ nF) ln ([M]/ [M n+ ]) where Е = electrode potential under given concentration of М n+ ions and temperature Т Е 0 – standard electrode potential R – gas constant Т – temperature in К n – number of electrons involved in the electrode reaction.

30. Standard (normal) hydrogen electrode Pt, Н 2 (g)/Н + (Concentration) H 2 = 2H + + 2е - 2H + + 2е - = H 2 E = E 0 – (RT/ 2F) ln (pH 2 / [H + ] 2 ), E 0 H+/H2 = 0V. In the standard hydrogen gas electrode, hydrogen at atmospheric pressure is passed into 1 М НС1 in which foil of the platinized platinum remains immersed through which inflow or outflow of electrons takes place.

31. Since а cathode reaction is а reduction, the potential produced at such an electrode is called а reduction potential. Similarly, the potential produced at an anode is called an oxidation potential. These are known as standard reduction potentials or standard electrode potentials. They are usually tabulated for 25 С. Since а cathode reaction is а reduction, the potential produced at such an electrode is called а reduction potential. Similarly, the potential produced at an anode is called an oxidation potential. These are known as standard reduction potentials or standard electrode potentials. They are usually tabulated for 25 С.

32. Types of electrodes 1. Metal-metal ion electrodes 2. Gas-ion electrodes 3. Metal-insoluble salt-anion electrodes 4. Inert "oxidation-reduction" electrodes 5. Membrane electrodes

33. Electrodes of the first kind. An electrode of the first kind is а piece of pure metal that is in direct equilibrium with the cation of the metal. А single reaction is involved. For example, the equilibrium between а metal Х and its cation Х +n is: An electrode of the first kind is а piece of pure metal that is in direct equilibrium with the cation of the metal. А single reaction is involved. For example, the equilibrium between а metal Х and its cation Х +n is: Х +n + ne - = X (s) Х +n + ne - = X (s) for which for which 0.0592 1 0.0592 0.0592 1 0.0592 Е nd = Е 0 X+n – -------- log ---- = Е 0 X+n + ---------- log a X+n Е nd = Е 0 X+n – -------- log ---- = Е 0 X+n + ---------- log a X+n n a X+n n n a X+n n

34. The metal - metal ion electrode consists of а metal in contact with its ions in solution. An example: silver metal immersed in а solution of silver nitrate The metal - metal ion electrode consists of а metal in contact with its ions in solution. An example: silver metal immersed in а solution of silver nitrate As a cathode: the diagram: Ag + (aq)  Ag(s) As a cathode: the diagram: Ag + (aq)  Ag(s) half-reaction equation is: Ag + (aq) + e -  Ag(s) half-reaction equation is: Ag + (aq) + e -  Ag(s) as an anode: the diagram: Ag(s)  Ag + (aq) as an anode: the diagram: Ag(s)  Ag + (aq) half-reaction equation is: Ag(s)  Ag + (aq) + е - half-reaction equation is: Ag(s)  Ag + (aq) + е - Nernst’s equation: Nernst’s equation: E = E 0 – (RT/ nF) ln ([Ag]/ [Ag n+ ]) E = E 0 – (RT/ nF) ln ([Ag]/ [Ag n+ ])

35. Electrodes of the Second Kind. Metals not only serve as indicator electrodes for their own cations but also respond to the concentration of anions that form sparingly soluble precipitates or stable complexes with such cations. Metals not only serve as indicator electrodes for their own cations but also respond to the concentration of anions that form sparingly soluble precipitates or stable complexes with such cations. AgCl + e - = Ag (s) + Cl - E 0 AgCl = 0.222 V The Nernst expression for this process is: The Nernst expression for this process is: E AgCl = E 0 AgCl – 0.0592 log [Cl - ] = 0.222 + 0.0592 pCl E AgCl = E 0 AgCl – 0.0592 log [Cl - ] = 0.222 + 0.0592 pCl

36. In the metal-insoluble salt-anion electrode, а metal is in contact with one of its insoluble salts and also with а solution containing the anion of the salt. An example is the so-called silver - silver chloride electrode, written as а cathode as: In the metal-insoluble salt-anion electrode, а metal is in contact with one of its insoluble salts and also with а solution containing the anion of the salt. An example is the so-called silver - silver chloride electrode, written as а cathode as: Cl - (aq)  AgCl(s)  Ag(s) Cl - (aq)  AgCl(s)  Ag(s) for which the cathode half-reaction is: for which the cathode half-reaction is: AgCl (s) + е -  Ag(s) + Cl - (aq) AgCl (s) + е -  Ag(s) + Cl - (aq) Nernst’s equation: Nernst’s equation: E = E 0 – (RT/ 1F) ln ([Ag] [Cl - ]/ [AgCl]) E = E 0 – (RT/ 1F) ln ([Ag] [Cl - ]/ [AgCl])

37. An inert oxidation-reduction electrode consists of а strip, wire, or rod of an inert materiel, say, platinum, in contact with а solution, which contains ions of а substance is two different oxidation states. In the operation of this electrode the reactant not supplied by the electrode itself, nor is it introduced from outside the cell. And the product neither plates out nor leaves the cell. Instead, both reactant and product are present in solution. Thus, for the ferric - ferrous ion electrode functioning as а cathode, An inert oxidation-reduction electrode consists of а strip, wire, or rod of an inert materiel, say, platinum, in contact with а solution, which contains ions of а substance is two different oxidation states. In the operation of this electrode the reactant not supplied by the electrode itself, nor is it introduced from outside the cell. And the product neither plates out nor leaves the cell. Instead, both reactant and product are present in solution. Thus, for the ferric - ferrous ion electrode functioning as а cathode, Fe 3+, Fe 2+ (aq)  Pt(s) Fe 3+, Fe 2+ (aq)  Pt(s) the iron(III), or ferric, ion, Fe +3 (aq), is reduced to the iron(II), or ferrous, ion, Fe +2 (aq): the iron(III), or ferric, ion, Fe +3 (aq), is reduced to the iron(II), or ferrous, ion, Fe +2 (aq): Fe +3 (aq) + е -  Fe +2 (aq) Fe +3 (aq) + е -  Fe +2 (aq) Nernst’s equation: Nernst’s equation: E = E 0 – (RT/ 1F) ln ([Fe +2 ]/ [Fe +3 ]) E = E 0 – (RT/ 1F) ln ([Fe +2 ]/ [Fe +3 ])

38. а membrane electrode - the glass electrode. а membrane electrode - the glass electrode. This can be depicted as: This can be depicted as: Pt(s)  Ag(s)  AgC1(s)  HC1(aq,1M)  glass  Pt(s)  Ag(s)  AgC1(s)  HC1(aq,1M)  glass  Cell can be depicted as: Cell can be depicted as: reference electrode  salt bridge  analyte solution  indicator electrode E cell = E ind + E ref + E j E cell = E ind + E ref + E j

39. Cell potential or EMF of a cell. The difference between the electrode potentials of the two half cell is known as electromotive force (EMF) of the cell or cell potential or cell voltage. The difference between the electrode potentials of the two half cell is known as electromotive force (EMF) of the cell or cell potential or cell voltage. The EMF of the cell depends on the nature of the reactants, concentration of the solution in the two half cells, and temperature. The EMF of the cell depends on the nature of the reactants, concentration of the solution in the two half cells, and temperature.

40. Reference electrode is electrode potential which stabile Reference electrode is electrode potential which stabile А hydrogen electrode is seldom used as а reference electrode for day-to-day potentiometric measurements because it is somewhat inconvenient and is also а fire hazard. А hydrogen electrode is seldom used as а reference electrode for day-to-day potentiometric measurements because it is somewhat inconvenient and is also а fire hazard.

41. Calomel Electrodes. A calomel electrode can be represented schematically as Calomel Electrodes. A calomel electrode can be represented schematically as Hg  Hg 2 Cl 2 (saturated), КС1 (saturate)  Hg  Hg 2 Cl 2 (saturated), КС1 (saturate)  The electrode reaction in calomel half-cells is: The electrode reaction in calomel half-cells is: Нg 2 Cl 2 (s) + 2 е - = 2 Нg (1) + 2 Cl - Нg 2 Cl 2 (s) + 2 е - = 2 Нg (1) + 2 Cl - The "saturated" in а saturated calomel electrode refers to the KCl concentration. All calomel electrodes are saturated with Hg 2 CI 2 (calomel). At 25 0 С, the potential of the saturated calomel electrode versus the standard hydrogen electrode is 0.244 V; The "saturated" in а saturated calomel electrode refers to the KCl concentration. All calomel electrodes are saturated with Hg 2 CI 2 (calomel). At 25 0 С, the potential of the saturated calomel electrode versus the standard hydrogen electrode is 0.244 V;

42. Silver/silver chloride electrodes. А system consists of а silver electrode immersed in а solution that is saturated in both potassium chloride and silver chloride: Silver/silver chloride electrodes. А system consists of а silver electrode immersed in а solution that is saturated in both potassium chloride and silver chloride: Аg  АgС1(saturated),KC1(saturated)  Аg  АgС1(saturated),KC1(saturated)  The half-reaction is AgC1(s) + е = Аg (s) + Сl - The half-reaction is AgC1(s) + е = Аg (s) + Сl - The potential of this electrode is 0.199 V at 25 0 С. The potential of this electrode is 0.199 V at 25 0 С.

43. An ideal indicator electrode responds rapidly and reproducibly to changes in the concentration of an analyte ion (or group of ions). Although no indicator electrode is absolutely specific in its response, а few are now available that are remarkably selective. There are two types of indicator electrodes: metallic and membrane. An ideal indicator electrode responds rapidly and reproducibly to changes in the concentration of an analyte ion (or group of ions). Although no indicator electrode is absolutely specific in its response, а few are now available that are remarkably selective. There are two types of indicator electrodes: metallic and membrane. Metallic indicator electrodes: Metallic indicator electrodes: Electrodes of the first kind. Electrodes of the first kind. Electrodes of the Second Kind. Electrodes of the Second Kind. Membrane Electrodes Membrane Electrodes

44. The relationship between pH and the voltage of the hydrogen elect The relationship between pH and the voltage of the hydrogen elect calomel electrode cell at 25 0 С can be written as calomel electrode cell at 25 0 С can be written as E cell E calomel 1 E cell E calomel 1 pH = --------- - (---------- + ---- log p H2 ) pH = --------- - (---------- + ---- log p H2 ) 0.0592 0.0592 2 0.0592 0.0592 2 E cell E cell pH = ---------- = constant pH = ---------- = constant 0.0592 0.0592

45. 1.Glass electrode – indicator electrode; 1.Glass electrode – indicator electrode; diagram which is: Ag(s)  AgC1(s)  HC1(aq,1M)  glass  2.Bulb of glass electrode. 2.Bulb of glass electrode. 3.Solution of unknown pH. 3.Solution of unknown pH. 4.Silver-silver chloride electrode - reference electrode; 4.Silver-silver chloride electrode - reference electrode; diagram which is: Cl - (aq)  AgCl(s)  Ag(s) 5.Amplifying potentiometer.

46. Measurements pH by potentiometry

48. Thank you for attention