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Kharkiv National Medical University Department of Medical and B ioorganic chemistry

Kharkiv National Medical University Department of Medical and B ioorganic chemistry «Medical Chemistry » Lecture № 7 ELECTRODE PROCESSES, THEIR BIOLOGICAL ROLE AND USE IN MEDICINE Lecturer : As. Professor, Department of Medical and Bioorganic Chemistry ,, Ph.D. Lukianova L.V.

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Kharkiv National Medical University Department of Medical and B ioorganic chemistry

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  1. KharkivNational Medical University Department of Medical and Bioorganic chemistry «Medical Chemistry» Lecture№ 7 ELECTRODE PROCESSES, THEIR BIOLOGICAL ROLE AND USE IN MEDICINE Lecturer: As. Professor, Department of Medical and Bioorganic Chemistry,, Ph.D. LukianovaL.V.

  2. PLAN OF LECTURE • The biological significance of redox reactions. • Oxidation-Reduction Reactions. • Basics of the theory oxidation-reduction reactions. • Redox processes and periodic system. • Influence of the medium on the stroke of the redox reaction. • Change of oxidants and reductants in the reaction. • Galvanic cells. • Origin of electrode potential. • Electrochemical series. • Nernst equation. • Reduction potentials. • Electrodes of the third type (oxidation-reduction electrodes). • Quinhydrone electrode. • Membrane electrodes. • 15. Measuring of pH. • 16. Cells for pH measuring. • 17. Oxidation-reduction in the organism.

  3. THE BIOLOGICAL SIGNIFICANCE • OF REDOX REACTIONS Oxidation-Reduction Reactions are of great importance in biological systems. Redox reactions occur in the body during metabolism. They are a source of energy in the process of cell respiration. Due to redox reactions in the human body is the synthesis of amino acids, carbohydrates, hormones and other biologically important substances. In humans, there exists a redox homeostasis.

  4. 2. OXIDATION-REDUCTION REACTIONS Redox reactions – are reactions that occur with a change in the oxidation state of the atoms making up the molecules of the reactants. Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion. Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.

  5. OXIDATION-REDUCTION REACTIONS Reduction Oxidant + e- → Product (Electrons gained; oxidation number decreases) Oxidation Reductant → Product + e- (Electrons lost; oxidation number increases) Examples of redox reactions: H2 + Cl2 → 2 HCl the oxidation reaction: H2 → 2 H+ + 2 e− the reduction reaction: Cl2 + 2 e− → 2 Cl−

  6. Oxidation-reduction or redox – a system in which the indifferent electrodes does not exchange ions with a solution, but only provides a supply or removal of electrons for the oxidation-reduction reaction in a solution containing oxidized and reduced forms of the same substance. Scheme of redox-system: Pt | FeCl2, FeCl3.

  7. Redox – the potential is calculated by the equation of Peters: where, ered° – normal redox potential and this potential occurring at the electrode immersed in a solution at ratio in it oxidized and reduced forms equal to 1; n – number electrons, which gives or receives a reducing oxidant. At 18 ° C: at 25 ° C:

  8. The degree of oxidation – the charge is conditioned atom which is calculated on the assumption that the molecule consists of ions and the total charge of the molecule is zero. When calculating the degree of oxidation based on the fact that the degree of oxidation of the hydrogen is always 1 (except hydrides), oxygen – 2 (except peroxide), one alkali metal, alkaline – earth metals 2.

  9. For example, oxidation of nitrogen in the nitrate HNO3 acid calculated from the fact that the degree of oxidation of hydrogen +1, oxygen – 2, three oxygen atoms give – 6, then: +1 + X + (-2) · 3 = 0, X = 5.

  10. 3. BASICS OF THE THEORY OXIDATION-REDUCTION REACTIONS: 1) Reductants– are molecules, atoms or ions that donate electrons. At the same time they are oxidized: Н20 –2e → 2Н+; 2Cl ˉ – 2e → Cl20 . The most common reducing agents: molecule: CO, H₂, formic aldehyde; atoms of metals, nonmetals (C, S, P); Negatively charged ions nonmetals: Clˉ, Brˉ, Iˉ, S⁻², N⁻³, P⁻²; metal ions in a lower degree of oxidation: Fe⁺², Cr⁺³, Mn⁺²; electrical current at the cathode.

  11. 2) Oxidizing agents – are molecules, atoms or ions, which accept electrons. At the same time they • are restored: • Cl20 + 2e → 2Clˉ; • S0 + 2е → Sˉ2; • Fe+3 + 1е → Fe+2. • The most common oxidants: • molecule: O2, O3, KMnO4, MnO2, K2Cr2O7, PbO2, CrO3, HNO3, halogens; • nonmetals in the positive degree of oxidation: N+5, S+6, Cl+, Cl+3,Cl+5, Cl+7; • metal ions in a higher oxidation: Fe+3, Cr+6, Mn+7, Pb+4; electrical current at the anode.

  12. 3) Oxidation – a process recoil electrons of molecules, atoms or ion. I. e. oxidation by oxidation increases. 4) Recovery – is the process of joining the electron molecule, atom or ion. I.e. the reduction degree of oxidation is reduced. Oxidation is always accompanied by reduction and vice versa. The number of electrons, which gives a reducing agent equals the number of electrons which takes oxidant. 5) Compounds which contain a degree of oxidation of the intermediate may be either oxidizing or reducing agents: HN+3O2, H2S+4O3, H3As+3O3, K2Mn+6O4.

  13. 4. REDOX PROCESSES AND PERIODIC SYSTEM In periods with increasing atomic number of the element reducing properties decrease and increase oxidation, since it is easier to accept electrons to complete the energy level . For example, alkaline metals – strong reducing agents, halogens – strong oxidizing agents. In major subgroups recovery properties increase as it increases the radius of the atom and the electrons easier to split off. In side subgroups metals are only so they restorers.

  14. Redox properties associated with the electronegativity: the more electronegative element, the stronger its oxidizing properties (F – the most electronegative element). On the contrary, metals having a low electronegativity are reducing. Redox properties depend on the degree of oxidation: the more positive charge of the same element, the more expressed oxidative properties: KMn+7O4 Mn+4O2 Mn+2SO4 oxidant oxidant and reductantreductant

  15. INFLUENCE OF THE MEDIUM ON THE STROKE OF THE REDOX REACTION 1) creating an acidic environment using H₂SO₄. HCl Hydrochloric acid may be not only medium but also a reducing agent. Nitric acid HNO3, may be not only medium but also the oxidant. 2) to create an alkaline environment using alkali NaOH, KOH, and Na₂CO₃. Influence of the medium on the stroke of the redox reaction can be show on the reduction reaction of KMnO4.

  16. KMn+7O4 Н+ (+5е) → Mn+2 (MnSO4, MnCl2 – colorless solution); Н2О (+3е) → Mn+4 (MnО2↓ – brown precipitate); ОНˉ (+1е) → Mn+6 (К2MnO4 – green solution).

  17. CHANGE OF OXIDANTS AND REDUCTANTS IN THE REACTION 1) In an acidic medium the H⁺ ions and OH⁻ form water. 2) In an acidic medium with metal cations (+1, +2, +3) to form salts with acidic residues. 3) Metal ions, which give the water-insoluble base in alkaline and neutral environments, corresponding to provide base (Fe (OH)₃, Cu (OH)₂). 4) The metal ions which give amphoteric hydroxides in alkaline medium is allowed the corresponding salts (Na3[Cr(OH)6], Na2[Pb(OH)4).

  18. GALVANIC CELL Galvanic cell – the device in which electrical energy is produced from chemical reactions. Electrode- is a metal strip at which oxidation or reduction occurs. Oxidation takes place at the anode (negative pole): Zn0→Zn2+ + 2e- Reduction takes place at the cathode (positive pole): 2e- + Cu2+→ Cu0 Electrons flow from anode to cathode. Salt bridge is used for neutrality maintaining. The Daniell Cell

  19. + _ Cu Zn ZnSO4 CuSO4 Cu2+ + 2e  Cu Zn – 2e  Zn2+

  20. GALVANIC CELLS Chemical cells. The chemical reaction energy turns into the electric energy in such kind of cells. Daniell cell is an example of the chemical cell. The E.M.F. occurs in the chemical elements due to the different chemical nature of electrodes.

  21. GALVANIC CELLS The concentrating cells. The similar electrodes in such chains are plunged into the solutions of the same electrolytes having different concentrations: Ag | AgNO3 || AgNO3 | Ag C2 < C1 The E.M.F. occurs as a result of the concentration leveling between the solutions. The electrode in the solution with a higher ion concentration will serve as a cathode, the ions move from left to right. The E.M.F. depends on the concentrations ratio of the potential-determining ions and not on the electrodes nature.

  22. GALVANIC CELLS The oxidation-reduction cells are the galvanic cells consisting of two oxidation-reduction electrodes: The anode reaction is: Sn2+ – 2e– Sn4+ The cathode reaction is: Fe3+ + e– Fe2+  The combined reaction, which is a source of E.M.F. in this element, is: Sn2+ + 2Fe3+ ↔ Sn4+ + 2Fe2+

  23. 7. GALVANIC CELL REPRESENTATION • The anode is written on the left hand side and cathode on the right hand side. • 2. The anode of the cell is represented by writing metal first and then the electrolyte (or cation of the electrolyte). The cathode is represented by writing the cation first and then the metal. • 3. The salt bridge is indicated by two vertical lines. • Daniell cell: (-)Zn|Zn2+||Cu2+|Cu(+)

  24. Me - - - - - Me+z + + + + + 8. ORIGIN OF ELECTRODE POTENTIAL If metal has high tendency to get oxidized, its atoms lose electrons and form cations, which go into the solution. The electrons lost accumulate on the metal electrode and the electrode becomes negatively charge with respect to the solution. The equilibrium is established after some time as: Me(s) → Mez+ + ze- electrode solution (on electrode) Now charges are separated – (-) on electrode, (+) – in solution.

  25. Me + + + + + - - - - - Me+z If the metal ions have greater tendency to get reduced, they will take electrons from the electrode. In this case, a positive charge will be developed on the electrode with respect to the solution. This will also result into separation of charges (positive on the electrode with respect to the solution). Now charges are separated – (+) on electrode, (-) – in solution. Due to separation of charges between the electrode and the solution, an electrical potential is set up between metal electrode and its solution. Electrode potential–is the potential difference between the electrode and itssolution. Thus, the electrode potential is a measure of tendency of an electrode in a half cell to gain or to lose electrons.

  26. 9. Reduction potentials The half cells potentials are represented as reduction potentials! Reduction potential – is the tendency of an electrode to gain electronsor to get reduced. The electrode potential depends upon: 1. The nature of the metal and its ions 2. Concentration of the ions in the solution 3. Temperature

  27. Electro-Motive Force Galvanic cell consists of two half cells. The electrodes in these half cells have different reduction potentials. The electrode with a lower reduction potential will lose electrons and electrode having higher potential will gain these electrons. So, the result of a potential difference is a flow of electrons. Electromotive force – is the difference between the electrode potentials of the two electrodes constituting a galvanic cell. e.m.f. is a driving force for the cell reaction. The e.m.f. is expressed in volts. e.m.f. = e(cathode) - e(anode) e.m.f. is always positive!

  28. H2 gas at 1 atm Platinum wire Bubbles of H2 Platinum plate 1 M HCl solution Standard Electrode Potential It is not possible to determine the absolute value of electrode potential because electrode can work only in combination with the other electrode. Only relative potential can be found! Hydrogen electrode has been accepted as reference electrode. Potential of hydrogen electrode at standard conditions (T=298K, P=1 atm,conc. of H+=1M) is taken as zero! Standard electrode potential (e0) – is the electrode potential of a metal electrode as determined with respect to a standard hydrogen electrode (S.H.E.).

  29. THE STANDARD HYDROGEN ELECTRODE Definition: The standard hydrogen electrode is the standard measurement of electrode potential for the thermodynamic scale of redox potentials.The standard is determined by the potential of a platinum electrode in the redoxhalf reaction2 H+(aq) + 2 e- → H2(g) at 25 °C.The standard hydrogen electrode is often abbreviated SHE. Also Known As: normal hydrogen electrode or NHE

  30. Voltmeter e- e- H2 gas at 1 atm Me 1M H+ 1M Mez+ 1M solution 1M HCl To measure the e0 of a metal it should be connected to S.H.E. Concentrations of solutions are 1 mol/l. Electrodes are connected to a voltmeter. Voltmeter indicates e.m.f. e.m.f. = e(electrode of unknown potential) – e(standard electrode) or e.m.f. = e(standard electrode) – e(electrode of unknown potential) e(standard electrode) = o, so, e(electrode of unknown potential) can be calculated!

  31. 9. ELECTROCHEMICAL SERIES The standard electrode potentials of a large number of electrodes have been measured using standard S.H.E. The arrangement of elements in order of increasing reduction potential values is called electrochemical series (activity series). The elements at the bottom of the table are good oxidizing agents. The elements at the top of the table are reducing agents. A favorable or spontaneous reaction occurs between a reducing agent and any oxidizing agent in the table that is lower than the reducing agent! Al|Al3+||Ni2+|Ni

  32. 10. NERNST EQUATION e – electrode potential; e0 – the standard electrode potential; n - number of electrons gained or lost in reaction; C – molar concentration of solution

  33. 11. CLASSIFICATION OF ELECTRODESElectrodes of the first type Electrodes are classified depending on their composition and electrode reaction. The electrodes of the first type consist of metal strip placed in the solution of its salt. Scheme:Me/Men+ ; Electrode reaction:Me – ne- Men+ . Electrode potential depends on the concentration of metal ions, and, therefore such an electrode can be used for their determination. E. g. Cu|Cu2+; Zn|Zn2+

  34. 12. ELECTRODES OF THE SECOND TYPECALOMEL ELECTRODE These are electrodes consisting of metal covered with a hardly soluble compound of this metal (salt, oxide, hydroxide) and plunged into a solution of the readily soluble compound with the same anion. Calomel electrode consists of mercury covered with calomel: Hg, Hg2Cl2. This is deepen in the solution of KCl. Scheme: Hg, Hg2Cl2| KCl Electrode reaction: Hg2Cl2 + 2e-→ 2Hg + 2Cl- Electrode potential depends only on the concentration of chloride ions in the solution: ecal=e° - 0.059lgCCl- In the saturated solutionof KClecal is constantand equals ecal = 0.248V at T = 298K.

  35. SILVER-CHLORIDE ELECTRODE The silver-chloride electrode consists of a silver wire covered with a layer of a silver-chloride plunged into the KClsolution. Scheme: Ag, AgCl| KCl Electrode reaction: AgCl + e → Ag + Cl- Electrode potential depends on the concentration of chloride ions in the solution. es.c=e° - 0.059lgCCl- In the saturated solutionof KCles. c. is constantand equalses. c. = 0.222V at T = 298K.

  36. + 2H+ + 2 e- hydroquinone quinone QUINHYDRONE ELECTRODE A quinhydrone electrode is referred to red-ox electrodes. It consists of platinum wire soldered in the glass tube and immersed into then investigated solution to which quinhydrone crystals are added. Scheme: Pt|H+, quinhydrone Electrode potential: e quin.=e0 + 0.059lgCH+= e0 – 0.059pH; e0 = 0.7 V; e = 0.7 - 0.059pH

  37. ANTIMONY ELECTRODE Antimony electrode consists of antimony covered with antimony oxide (III) and plunged into the solution containing H+. Scheme: Sb,Sb2O3|H+ Electrode reaction: Sb2O3 + 6H+→ 2Sb3+ + 3H2O Electrode potential depends on the concentration of H+. e=e0+0.059lgCH+ e=e0–0.059pH

  38. H2 gas at 1 atm Platinum wire Bubbles of H2 Platinum plate 1 M HCl solution GAS ELECTRODES Hydrogen electrode is an example of gas electrodes. Scheme: Pt(H2)|H+ Electrode reaction: 2H++2 e- = H2 Electrode potential depends on the concentration of H+. e = e0 + 0.059lgC H+ e = e0 - 0.059pH; e0 = 0, e = - 0.059pH

  39. ELECTRODES OF THE THIRD TYPE(OXIDATION-REDUCTION ELECTRODES) Red-Ox electrodes consist of metal (platinum or gold) immersed in a solutioncontaining oxidized and reduced forms of the same substance. Scheme: e.g. Electrode reaction: OX + ne-→ RED Electrode potential:

  40. The E.M.F. can be calculated as: a, b, c, d – are stoichiometric coefficients in the reaction equation.

  41. silver-chloride electrode buffer solution ball made of special electrode glass MEMBRANE ELECTRODES Potential of membrane electrodes occurs on the boundary of a thin layer of theelectrode material and a solution. The most widespread membrane electrode isa glass electrode. Scheme: Glass el.|H+ Electrode potential: egl. = e0 – 0.059pH

  42. ION-SELECTIVE ELECTRODES Their mechanism is similar to the mechanism of the glass electrode. To make such electrodes the wide number of electro-chemically active substances such as liquid and solid ionites, mono and polycrystals, synthetic membrane-active chelates, is used. Depending on the type of the electrode material the ion-selective electrodes can be divided into three groups: solid, liquid and membrane. Today more than 20 ion-selective electrodes destined for the determination of and other ions concentration are created.

  43. Indicating electrodes Reference electrodes Hydrogen Quinhydrone Antimony Glass Calomel Silver-chloride MEASURING OF pH In order to measure pH it needs a galvanic cell consisting ofindicating electrode and a reference electrode. Reference electrode has a constant potential. Potential of indicating electrode depends on pH.

  44. POTENTIOMETRY. DETERMINATION OF PH BY MEANS OF POTENTIOMETRY The internal media of an organism such as blood, lymph, gastric juice and urine are water solutions; pH of these solutions affects the vital activity of cells, tissues, organs and an organism as a whole. pH stability of organism systems is a pledge of its normal vital activity. Checking of this value enables to found out different kinds of pathology and to make the right diagnosis. The electrochemical method of pH measuring is widely used in chemistry, medicine, and biochemistry due to its high accuracy. Besides, it enables to measure pH without changing the composition and properties of the investigated solutions.

  45. MEASURING OF pH Hydrogen electrodeis a primary standard electrode used for pH measuring. Its potential linearly depends on the pH value: e= – 0,059pH It is rather inconvenient for serial measurements. A quinhydrone electrodeis convenient for the laboratory work though it can be used for the pH measuring only in acidic or weak-alkaline area (pH  8) An antimony electrodeis convenient for serial analyses, pH measuring of the concentrated solutions of salts and of the colloidal solutions. A glass electrodeis the most widely used one for pH measuring. To compose a galvanic cell for pH measuring it needs an indicating electrode and a reference electrode with a constant potential. The reference electrodes are calomelandsilver-chloride ones.

  46. CELLS FOR PH MEASURING • Cell consisting of hydrogen and reference electrode: • Pt(H2)|H+||KCl|AgCl, Ag • e.m.f. is indicating by voltmeter. • e.m.f. = e(cathode) – e(anode) = 0.222 - (-0.059pH) = = 0.222 + 0.059pH • 2. Cell consisting of quinhydrone and reference electrode: • Hg, Hg2Cl2|KCl||H+ quin.|Pt • A quinhydrone electrode's potential is more positive than calomel's, so it acts as cathode. • e.m.f. = e(cathode) – e(anode) = equin. – ecal. = • = 0.7 - 0.059pH - 0.248 • 3. A cell composed of an antimony electrode and a reference electrode. This kind of cell is used for pH measuring in a stomach cavity (Linar probe). • Sb, Sb2O3|H+||KCl|Hg2Cl2,Hg • Glass electrodes are very often used for pH measuring.

  47. CONCENTRATING CELL Concentrating cell consisting of two indicating electrodes plunged into the solutions with different concentrations of hydrogen ions. Pt(H2) | H+ ||H+ | (H2)Pt C2 < C1 For the pH measuring we can use a concentrating cell consisting of two glass electrodes, one of which is plunged into a standard solution with the determined pH value. This kind of chains can be used for the blood and gastric juice pH measuring.

  48. DIFFUSE POTENTIAL Diffuse potential – is the potential difference on the borderline of two electrolyte solutions with different concentration or different composition due to different ion mobility. Diffuse potentials can appear in biological objects in the result of impairment, e.g. cellular membranes.The electrolytes travel from the place of impairment to the intact areas. The damaged tissue is charged negatively when compared with the intact one, it means that diffuse potential of the impairment appears. It equals about 30–40 mV.

  49. МЕМBRANE POTENTIAL Membranepotentialisformedontheborderlineoftwosolutionsifthereis a semi-permeablemembrane, whichpassescationsandcapturesanions.Therefore, onesideofthemembraneischargedpositively, theother – negatively. The changes in membrane potential which accompany transmission of nerve impulses or muscular contraction are due to the flow of potassium cations outside the cell and sodium cations inside the cell. This results in potential diminishing, which can be registered using microelecrodes placed inside and outside the cell.

  50. REST POTENTIAL Rest potential – is potential difference measured in the state of physiological rest of the cell. Rest potential in different cells equals 50–100 mV. The cause of biopotential appearance is uneven distribution of K+ and Na+ ions. The number of K+ ions in the intracellular fluid is 20–40 times higher than in the extracellular fluid, while Na+ ions concentration is 10–20 times higher in the extracellular one. The ions of organic acids pass the membrane with difficulty. At rest K+ ions pass from the intracellular fluid to extracellular. Thus, inner surface of the cell is charged negatively, while external is charged positively.

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