Buffer solutions. Theoretic bases of electrochemistry.

LECTERE 2LECTERE 2

Lecturer: Dmukhalska Ye. B.

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

Water is а neutral molecule with а slight tendency to ionize. We usually express this ionization as:Н2О = Н+ + ОН-

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, H3O

+. The association of а

proton with а cluster of water molecules also gives rise to structures with the formulas Н5О2

+, Н7О3+, and so on. For

simplicity, however, we collectively represent these ions by H+.

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), solutions with relatively more ОН- are basic (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

Кw = [Н+][ОН-] =10 -14 M2 at 25 0C.

[Н+] = [ОН-] = (Кw)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.

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 -14Household ammonia -12Seawater – 8Milk - 7Blood  - 7.4Saliva - 6.6Tomato juice - 4.4Vinegar - 3Gastric juice - 1.51 М НСl - 0

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 NН3, 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 + Н2О = Н3О+ + А-

An acid (НА) reacts with а base (Н2О) to form the conjugate base of the acid (А-) and the conjugate acid of the base (H3O

+). Accordingly, the acetate ion (СН3СОО- ) is the conjugate base of acetic acid (СН3СООН), and the ammonium ion (NH4

+ ) is the conjugate acid of ammonia (NН3). The acid-base reaction is frequntly abbreviated НА = Н+ + А- with the participation of H2O implied.

The strength of an acid is specified by its dissociation constantThe 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: [Н3О

+][А-]

K= ---------------- [НА] [Н2O]

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 [Н2О] is

customarily combined with the dissociation constant, which then takes the form [Н+][А-]Ka = K[Н2O] = -------------

[НА] Because acid dissociation constants, like [Н+] values, are sometimes cumbersome to work with, they are transformed to pK values by the formulaрK = - log K

The relationship between the pH of а solution and the concentrations The relationship between the pH of а solution and the concentrations of an acid and its conjugate base is easily derived. of an acid and its conjugate base is easily derived.

[НА] [НА] [Н[Н++]= ]= KK ---------- ---------- [[АА--]]Taking the negative log of each term Taking the negative log of each term [А[А--]]рН = - рН = - log log К + К + log log ------------------ [А[А--]] [А[А--]]рН = рН = ppК + К + log log ------------------ [[АА--]]This relationship known as the Henderson-Hasselbalch equation. This relationship known as the Henderson-Hasselbalch equation.

BUFFERSBUFFERS

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

Buffers are mainly of two types: Buffers are mainly of two types: • ((аа) mixtures of weak acids with their ) mixtures of weak acids with their

salt with salt with аа strong base strong base • (b) mixtures of weak bases with their (b) mixtures of weak bases with their

salt with salt with аа strong acid. strong acid. АА few examples are given below: few examples are given below:• НН22СОСО33 / N / NаНСОаНСО33 (Bicarbonate buffer; (Bicarbonate buffer;

carbonic acid and sodium bicarbonate)carbonic acid and sodium bicarbonate)• СНСН33СООНСООН / / СНСН33СООСОО Na (Acetate Na (Acetate

buffer; acetic acid and sodium acetate)buffer; acetic acid and sodium acetate)• NaNa22HPOHPO44/ NaH/ NaH22POPO4 4 (Phosphate buffer)(Phosphate buffer)

Factors Affecting pH of Factors Affecting pH of аа BufferBuffer The pH of The pH of аа buffer solution is determined by buffer solution is determined by

two factors:two factors:

• 1. The value of pK: The lower the value of 1. The value of pK: The lower the value of pK, the lower is the pH of the solution.pK, the lower is the pH of the solution.

• 2. The ratio of salt to acid concentrations: 2. The ratio of salt to acid concentrations: Actual concentrations of salt and acid in Actual concentrations of salt and acid in аа buffer solution may be varied widely, with buffer solution may be varied widely, with попо change in change in рНрН, so long as the ratio of , so long as the ratio of the concentrations remains the same.the concentrations remains the same.

Buffer CapacityBuffer Capacity • On the other hand, the buffer capacity is On the other hand, the buffer capacity is

determined by the actual concentrations of determined by the actual concentrations of salt and acid present, as well as by their salt and acid present, as well as by their ratio. Buffering capacity is the number of ratio. Buffering capacity is the number of grams of strong acid or alkali which is grams of strong acid or alkali which is necessary for necessary for аа change in pH of one unit of change in pH of one unit of one litre of buffer solution.one litre of buffer solution.

• The buffering capacity of The buffering capacity of аа buffer is, definеd buffer is, definеd ааs the ability of the buffer to resist changes s the ability of the buffer to resist changes in pH when an acid or base is added.in pH when an acid or base is added.

Buffers ActBuffers Act • When hydrochloric acid is added to the acetate buffer, the salt When hydrochloric acid is added to the acetate buffer, the salt

reacts with the acid forming the weak acid, acetic acid and its reacts with the acid forming the weak acid, acetic acid and its salt. Similarly when salt. Similarly when аа base is added, the acid reacts with it base is added, the acid reacts with it forming salt and water. Thus, changes in the pH are minimised.forming salt and water. Thus, changes in the pH are minimised.

• СНСН33СООН + NaOH = СНСООН + NaOH = СН33COONa + НCOONa + Н22ОО

• СНСН33СООСООNNа + HCI = СНа + HCI = СН33СООН + NaCIСООН + NaCI• The buffer capacity is determined by the absolute The buffer capacity is determined by the absolute

concentration of the salt and acid. But the concentration of the salt and acid. But the рНрН of the buffer is of the buffer is dependent on the relative proportion of the salt and acid (see dependent on the relative proportion of the salt and acid (see the Henderson - Hasselbalch's equation). When the ratio the Henderson - Hasselbalch's equation). When the ratio between salt and acid is 10:1, the pH will be one unit higher 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 than the pKa. When the ratio between salt and acid is 1:10, the pH will be one unit lower than the pKa.pH will be one unit lower than the pKa.

Mechanisms for Regulation of pHMechanisms for Regulation of pH

• (1)(1)      Buffers of body fluids, Buffers of body fluids,

• (2)(2)      Respiratory system, Respiratory system,

• (3) Renal excretion. (3) Renal excretion.

• These mechanisms are interrelated. These mechanisms are interrelated.

DefinitionDefinition

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

 

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

М = Мn+ + ne- Мn+ + nе- = М

Nernst’s equationNernst’s equation The dependence of cell voltage upon concentration The dependence of cell voltage upon concentration

can also be described quantitatively. The free-can also be described quantitatively. The free-energy change energy change G for any reaction is:G for any reaction is:

G =G =G G 00+ RT ln Q+ RT ln Q• Where: Q represents the mass-action expression Where: Q represents the mass-action expression

for an oxidation-reduction reactionfor an oxidation-reduction reactionG = - nFE, and G = - nFE, and GG00 = - nFE = - nFE00

• - nFE = - nFE- nFE = - nFE00+ RT ln Q+ RT ln Q• E = EE = E0 0 - RT/ nF - RT/ nF xx ln Q ln Q• R = 8.315 J/K R = 8.315 J/K ..molmol• FF = 96,485 С / = 96,485 С /molmol

• Мn++nе = М • Then the Nernst eqn. is applied as follows:• E = E0 – (RT/ nF) ln ([M]/ [Mn+])• where Е = electrode potential under given concentration

of Мn+ ions and temperature Т• Е0 – standard electrode potential• R – gas constant• Т – temperature in К• n – number of electrons involved in the electrode

reaction.

Standard (normal) Standard (normal) hydrogen hydrogen electrodeelectrode

Pt, НPt, Н2 2 (g)/Н(g)/Н++ (Concentration) (Concentration)

HH22 = 2H = 2H++ + 2е + 2е- -

2H2H++ + 2е + 2е-- = H = H22

E = EE = E00 – (RT/ 2F) ln (pH – (RT/ 2F) ln (pH22/ [H/ [H++]]22), E), E00H+/H2H+/H2 = 0V. = 0V.

In the standard hydrogen gas electrode, hydrogen at atmospheric In the standard hydrogen gas electrode, hydrogen at atmospheric pressure is passed into 1 М НС1 in which foil of the platinized pressure is passed into 1 М НС1 in which foil of the platinized platinum remains immersed through which inflow or outflow of platinum remains immersed through which inflow or outflow of electrons takes place. electrons takes place.

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

Types ofTypes of electrodeselectrodes

1. Metal-metal ion electrodes1. Metal-metal ion electrodes

2. Gas-ion electrodes2. Gas-ion electrodes

3. Metal-insoluble salt-anion electrodes3. Metal-insoluble salt-anion electrodes

4. Inert "oxidation-reduction" electrodes4. Inert "oxidation-reduction" electrodes

5. Membrane electrodes5. Membrane electrodes

ElectrodesElectrodes of the first of the first kind.kind.• An electrode of the first kind is а piece of pure metal An electrode of the first kind is а piece of pure metal

that is in direct equilibrium with the cation of the that is in direct equilibrium with the cation of the metal. А single reaction is involved. For example, the metal. А single reaction is involved. For example, the equilibrium between а metal Х and its cation Хequilibrium between а metal Х and its cation Х+n+n is: is:

• ХХ+n+n + ne + ne-- = X (s) = X (s)• for which for which 0.0592 1 0.0592 0.0592 1 0.0592

• ЕЕndnd = Е = Е00X+nX+n – -------- log ---- = Е – -------- log ---- = Е00

X+nX+n + ---------- log a + ---------- log aX+nX+n

• n an aX+nX+n n n

• The metal - metal ion electrodeThe metal - metal ion electrode consists of а consists of а metal in contact with its ions in solution. An metal in contact with its ions in solution. An example: silver metal immersed in а solution of example: silver metal immersed in а solution of silver nitrate silver nitrate

• As a cathode: the diagram: AgAs a cathode: the diagram: Ag++(aq) (aq) Ag(s) Ag(s)

• half-reaction equation is: Aghalf-reaction equation is: Ag++ (aq) + e (aq) + e--Ag(s)Ag(s)

• as an anode: the diagram: Ag(s) as an anode: the diagram: Ag(s) Ag Ag++(aq)(aq)

• half-reaction equation is: Ag(s) half-reaction equation is: Ag(s) Ag Ag++(aq) + е(aq) + е--

• Nernst’s equation:Nernst’s equation:

• E = EE = E00 – (RT/ nF) ln ([Ag]/ [Ag – (RT/ nF) ln ([Ag]/ [Agn+n+])])

Electrodes of the Second Kind.Electrodes of the Second Kind.• Metals not only serve as indicator electrodes for their Metals not only serve as indicator electrodes for their

own cations but also respond to the concentration of own cations but also respond to the concentration of anions that form sparingly soluble precipitates or stable anions that form sparingly soluble precipitates or stable complexes with such cations. complexes with such cations.

AgCl + eAgCl + e-- = Ag (s) + Cl = Ag (s) + Cl-- E E00AgClAgCl = 0.222 V = 0.222 V

• The Nernst expression for this process is:The Nernst expression for this process is:

• EEAgClAgCl = E = E00AgClAgCl – 0.0592 log [Cl – 0.0592 log [Cl--] = 0.222 + 0.0592 pCl] = 0.222 + 0.0592 pCl

• In In the metal-insoluble salt-anion electrodethe metal-insoluble salt-anion electrode, а metal is in , а metal is in contact with one of its insoluble salts and also with а contact with one of its insoluble salts and also with а solution containing the anion of the salt. An example is solution containing the anion of the salt. An example is the so-called silver - silver chloride electrode, written as а the so-called silver - silver chloride electrode, written as а cathode as:cathode as:

• ClCl-- (aq) (aq) AgCl(s) AgCl(s) Ag(s) Ag(s)• for which the cathode half-reaction is:for which the cathode half-reaction is:• AgCl (s) + еAgCl (s) + е-- Ag(s) + Cl Ag(s) + Cl-- (aq) (aq)• Nernst’s equation:Nernst’s equation:• E = EE = E00 – (RT/ 1F) ln ([Ag] [Cl – (RT/ 1F) ln ([Ag] [Cl--]/ [AgCl])]/ [AgCl])

• An inert oxidation-reduction electrodeAn inert oxidation-reduction electrode consists of а strip, wire, or rod of consists of а strip, wire, or rod of an inert materiel, say, platinum, in contact with а solution, which contains 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 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 electrode the reactant not supplied by the electrode itself, nor is it introduced from outside the cell. And the product neither plates out nor introduced from outside the cell. And the product neither plates out nor leaves the cell. Instead, both reactant and product are present in solution. leaves the cell. Instead, both reactant and product are present in solution. Thus, for the ferric - ferrous ion electrode functioning as а cathode,Thus, for the ferric - ferrous ion electrode functioning as а cathode,

• FeFe3+,3+, Fe Fe 2+2+(aq) (aq) Pt(s) Pt(s)• the iron(III), or ferric, ion, Fethe iron(III), or ferric, ion, Fe+3+3(aq), is reduced to the iron(II), or ferrous, (aq), is reduced to the iron(II), or ferrous,

ion, Feion, Fe+2+2(aq):(aq):• FeFe+3+3(aq) + е(aq) + е-- FeFe+2+2(aq)(aq)• Nernst’s equation:Nernst’s equation:• E = EE = E00 – (RT/ 1F) ln ([Fe – (RT/ 1F) ln ([Fe+2+2]/ [Fe]/ [Fe+3+3])])

•а membrane electrode а membrane electrode - the glass electrode. - the glass electrode. •This can be depicted as:This can be depicted as:•Pt(s)Pt(s)Ag(s)Ag(s) AgC1(s) AgC1(s)HC1(aq,1M)HC1(aq,1M) glass glass

•Cell can be depicted as:Cell can be depicted as:reference electrode reference electrode salt bridge salt bridge analyte solution analyte solution indicator electrode indicator electrode

•EEcell cell = E= Eindind + E + Erefref + E + Ej j

Cell potential or EMF of a cell.Cell potential or EMF of a cell. • The difference between the electrode The difference between the electrode

potentials of the two half cell is known as potentials of the two half cell is known as electromotive force (EMF) electromotive force (EMF) of the cell or of the cell or cell potential or cell voltage.cell potential or cell voltage.

• The EMF of the cell depends on the nature The EMF of the cell depends on the nature of the reactants, concentration of the of the reactants, concentration of the solution in the two half cells, and solution in the two half cells, and temperature.temperature.

• Reference electrode is electrode Reference electrode is electrode potential which stabilepotential which stabile

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

• Calomel Electrodes.Calomel Electrodes. A calomel electrode can A calomel electrode can be represented schematically asbe represented schematically as

• Hg Hg HgHg22ClCl2 2 (saturated), КС1 (saturate)(saturated), КС1 (saturate)• The electrode reaction in calomel half-cells is:The electrode reaction in calomel half-cells is:• НgНg22ClCl22(s) + 2 е(s) + 2 е-- = 2 Нg (1) + 2 Cl = 2 Нg (1) + 2 Cl--

• The "saturated" in а saturated calomel The "saturated" in а saturated calomel electrode refers to the KCl concentration. All electrode refers to the KCl concentration. All calomel electrodes are saturated with Hgcalomel electrodes are saturated with Hg22CICI22 (calomel). At 25(calomel). At 2500С, the potential of the С, the potential of the saturated calomel electrode versus the standard saturated calomel electrode versus the standard hydrogen electrode is 0.244 V; hydrogen electrode is 0.244 V;

• Silver/silver chloride electrodes.Silver/silver chloride electrodes. А system А system consists of а silver electrode immersed in а solution consists of а silver electrode immersed in а solution that is saturated in both potassium chloride and that is saturated in both potassium chloride and silver chloride:silver chloride:

• Аg Аg АgС1(saturated),KC1(saturated)АgС1(saturated),KC1(saturated)• The half-reaction is AgC1(s) + е = Аg (s) + СlThe half-reaction is AgC1(s) + е = Аg (s) + Сl--

• The potential of this electrode is 0.199 V at 25 The potential of this electrode is 0.199 V at 25 00С.С.

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

• The relationship between pH and the voltage of the The relationship between pH and the voltage of the hydrogen electhydrogen elect

• calomel electrode cell at 25calomel electrode cell at 2500С can be written asС can be written as

• EEcellcell E Ecalomelcalomel 1 1

• pH = --------- - (---------- + ---- log ppH = --------- - (---------- + ---- log pH2H2) )

• 0.0592 0.0592 2 0.0592 0.0592 2

• EEcellcell

• pH = ---------- = constantpH = ---------- = constant

• 0.0592 0.0592

• 1.Glass electrode – indicator electrode;1.Glass electrode – indicator electrode;diagram which is: Ag(s) diagram which is: Ag(s) AgC1(s) AgC1(s) HC1(aq,1M) HC1(aq,1M) glass 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: Cldiagram which is: Cl-- (aq) (aq) AgCl(s) AgCl(s) Ag(s) Ag(s)5.Amplifying potentiometer.5.Amplifying potentiometer.

Measurements pH by Measurements pH by potentiometry potentiometry

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