Introduction to Oxidation-Reduction Reactions PDF

Introduction to Oxidation- Reduction Reactions Electron Transfer Reactions Redox reactions The term oxidation originates from oxygen. Initially gain in oxygen has been considered as oxidation, Loss of oxygen has been considered to be reduction. Redox currently says that electrons are transfe...

Introduction to Oxidation- Reduction Reactions Electron Transfer Reactions Redox reactions The term oxidation originates from oxygen. Initially gain in oxygen has been considered as oxidation, Loss of oxygen has been considered to be reduction. Redox currently says that electrons are transferred between reactants. 1 Introduction to Oxidation- Reduction Reactions Redox reactions are characterized by ELECTRON TRANSFER between an electron donor and electron acceptor. Transfer leads to 1. Increase in oxidation number of the donor 2. Decrease in oxidation number of acceptor 2 Introduction to Oxidation- Reduction Reactions Assigning Oxidation States oxidation state = oxidation number Definition: The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. 3 Introduction to Oxidation- Reduction Reactions Rules for determining oxidation numbers 1. The oxidation number of elements in their standard states is zero. Example: Na, Be, K, Pb, H2, O2, P4 = 0 2. Oxidation number for monatomic ions are the same as their charge. Example: Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 3. Oxygen is assigned an oxidation state of -2 in its covalent compounds except in peroxides such as H2O2 where it is -1. 4 Introduction to Oxidation- Reduction Reactions 4. The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1. 5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge of the molecule or ion. +1 +6 -2 +6 -2 H2SO4 (SO4) -2 2.(+1) + 1.(+6) + 4.(-2)=0 1.(+6) + 4.(-2) = -2 5 Introduction to Oxidation- Reduction Reactions Problem Find the oxidation numbers of all of the species present in the following compounds: KMnO4, K2SO4, Fe2O3, NO3–, CH4, CCl4, HCHO, CH3COOH. +1+7-2 Answer: KMnO4 +x -2 (NO3 )– x + 3(-2) = -1 x – 6 = -1 ; x = +5 6 Introduction to Oxidation- Reduction Reactions Mg + S→ Mg2+ + S2- The magnesium atom changes to a magnesium ion by losing 2 electrons, and is thus oxidized Lose Electrons = Oxidation The sulfur atom is changed to a sulfide ion by gaining 2 electrons, and is thus reduced Gain Electrons = Reduction 7 Introduction to Oxidation- Reduction Reactions - Losing electrons is oxidation - the substance that loses electrons is called reducing agent/reductant - oxidation number increases - Gaining electrons is reduction, - the substance that gains electrons is called the oxidizing agent/oxidant - oxidation number decreases reducing agent oxidizing agent Mg(s) + S(s) → MgS(s) substance oxidized substance reduced 8 Introduction to Oxidation- Reduction Reactions Term Electron Oxidation change number change Oxidation Loss of Increase electrons Reduction Gain of Decrease electrons Oxidizing Accepts Decrease agent electrons Reducing Donates Increase agent electrons 9 Introduction to Oxidation- Reduction Reactions Oxidized and reduced form of an ion or a substance is assigned as redox pair. Examples: Cu2+/Cu+; Fe3+/Fe; O2/ 2O2-; 2H+/H2; NAD+/NADH, FAD/FADH2 Mg(s) + S(s) → MgS(s) Pairs : Mg 2+ / Mg n=2 S / S2- n=2 n - redox capacity; number of electrons exchanged between reduced and oxidized form of a substance 10  Introduction to Oxidation- Reduction Reactions Consider the equation: K2Cr2O7 + CH3CH2OH + H2SO4 CH3CHO + Cr2(SO4)3 + K2SO + H2O Is this equation represents redox reaction? +1 +6 -2 -3+1-1+1-2+1 +1+6 -2 -3+1+1+1-2 +3 +6 -2 +1+6-2 +1-2 K2Cr2O7 + CH3CH2OH + H2SO4 CH3CHO + Cr2(SO4)3 + K2SO4 + H2O 11  Introduction to Oxidation- Reduction Reactions K2Cr2O7+CH3CH2OH+H2SO4 Cr2(SO4)3+CH3CHO+K2SO4+H2O  Identifythe oxidizing agent, reducing agent, substance oxidized, substance reduced  oxidizing agent - K2Cr2O7  substance reduced - K2Cr2O7  reducing agent - CH3CH2OH  substance oxidized - CH3CH2OH 12 Introduction to Oxidation-Reduction Reactions Which redox pairs take part in the reaction? K2Cr2O7 / Cr2(SO4)3 CH3CH2OH / CH3CHO Short equation: +6 -1 +3 +1 2Cr + 3C 2Cr + 3C 13 Introduction to Oxidation- Reduction Reactions c) All redox reactions can be thought of as happening in two halves: one produces electrons – an oxidation half; the other requires electrons – a reduction half. Write the separate half reactions for the process above; Oxidation C-1 – 2e- C+1 6 3 Reduction Cr+6 +3e- Cr+3 2 d) Balance the chemical equation K2Cr2O7 + 3CH3CH2OH + 4H2SO4 3CH3CHO + Cr2(SO4)3 + K2SO4 + 7H2O 14  : Introduction to Oxidation- Reduction Reactions TYPES OF OXIDATION-REDUCTION REACTIONS Combination reactions O2 + 2 H2 2 H2O Decomposition reactions 2 H2O2 2 H2O + O2 Displacement reactions; Zn +2 HCl ZnCl2 + H2 Disproportionation and comproportionation reactions. 15  : Introduction to Oxidation- Reduction Reactions TYPES OF OXIDATION-REDUCTION REACTIONS Disproportionation - two or more atoms of the same element originally having the same oxidation state react with other chemical(s) or themselves to give ions with different oxidation numbers on that element. Cannizzaro reaction (disproportionation) +1 +3 -1 2C6H5-CHO + KOH → C6H5COOK + C6H5CH2OH benzaldehyde potassium benzyl alcohol benzoate 16 Introduction to Oxidation- Reduction Reactions Other examples for disproportionation: The dismutation of a superoxide free radical to hydrogen peroxide and oxygen, catalyzed in living systems by superoxide dismutase: 2O2− + 2H+ → H2O2 + O2 The decomposition of hydrogen peroxide to water and oxygen by the enzyme catalase: 2H2O2 → 2H2O + O2 In the HiPco method for producing carbon nanotubes, high pressure carbon monoxide disproportionates when catalysed on the surface of an iron particle: 2CO → C + CO2 17 Introduction to Oxidation- Reduction Reactions  Comproportionation A reaction in which an element in a higher oxidation state reacts with the same element in a lower oxidation state to give the element in an intermediate oxidation state. Example: Ag2+(aq)+Ag(s) → 2Ag+(aq) -2 +4 0 2H2S + SO2 → 3S + 2H2O  It is the reverse of disproportionation. 18 Introduction to Oxidation- Reduction Reactions The energy of a chemical reaction can be converted in electrical energy by galvanic cell. If a Zn rod is dipped in a solution of copper sulphate, the following reaction occurs. Zn + CuSO4 → ZnSO4 + Cu Oxidation Zn – 2е- → Zn2+ Reduction Cu2+ + 2е- → Cu Zn (s)+ Cu2+ (aq) → Zn2+ (aq) + Cu (s) Zn/Zn2+ ; Cu/Cu2+ Energy is released as heat energy. Zinc is oxidized with releasing two electrons in a half reaction In the other half reaction Cu2+ gains two electrons and gets reduced to copper. 19  Introduction to Oxidation-Reduction Reactions By carrying out the two half reactions in two different compartments, and allowing the electrons to flow through external circuit it is possible to convert the energy released to electrical energy. The combination of two metals inserted in solutions of their salts is a galvanic cell. The Zn compartment is the anode, and the copper compartment is the cathode. Electrode potential – emf of the half reaction Total emf = oxidation + reduction potential potential emf – electromotive force At 25° C and Oxidation Zn – 2е- → Zn2+ oxidation potential [ZnSO4]=[CuSO4]=1 mol/L Reduction Cu2+ + 2е- → Cu reduction potential emf = 1.107 V Introduction to Oxidation- Reduction Reactions Electrode potential – emf of the half reaction Total emf = oxidation + reduction potential potential Arbitrary standard – hydrogen electrode Electrode potential = 0.000 V at 25° C ½ H2 (1 atm) (g) ↔ H+ (aq) + e- - anode Cu2+ (aq) + 2е- → Cu (s) – cathode Total emf = oxidation + reduction = 0.000 V + reduction potential potential potential Total emf = reduction at standard conditions potential E0 – standard electrode potential (redox potential) Introduction to Oxidation- : Reduction Reactions REDOX POTENTIALS AND NERNST EQUATION Standard redox potential (Eo) is a constant characterizing the ability of a redox pair to act either as oxidant or reductant. The higher redox potential (positive value), the better oxidant; the lower redox potential (negative value), the better reductant. The real (measurable) redox potential (E) is concentration-dependent according to the Nernst equation: (1) E = Eo + R.T ln [oxidized form] n. F [reduced form] (2) E = Eo + 0.059 log [oxidized form] n [reduced form] n – number of electrons transferred between oxidized and reduced form 22 Introduction to Oxidation- : Reduction Reactions With regards to the standard redox-potentials find out in which direction the equilibrium is shifted in a solution. Use the following redox pairs accepting all concentrations of the substances to be equimolar. а) Fe3+ / Fe2+ and Sn4+ / Sn2+; Е0 Fe3+ / Fe2+ = +0.77V Е0 Sn4+ / Sn2+ = + 0.15 V Answer: 2Fe3+ + Sn2+ → 2Fe2+ + Sn4+ b) Zn2+ / Zn and Cu2+ / Cu Е0 Zn2+ / Zn = - 0.76 V Е0 Cu2+ / Cu = +0.34V Answer: Zn0 + Cu2+ → Zn2+ + Cu0 23 Biologic Oxidation Redox reactions going through electron transfer are not so common in organic compounds and therefore, in living systems. Most often the biological redox reactions represent a transfer of hydrogen atoms. The oxidation state of a molecule increases (oxidation) if its hydrogen content decreases or its oxygen content increases. Example: CH3CH2OH – 2H CH3CHO [O] CH3COOH Biological redox catalysts are either protein-bound (e.g. NAD+/NADH; FAD/FADH2), or low-molecular weight reductants that after oxidation are enzymatically recovered in their reduced forms (e.g. glutathione, ascorbic acid = vitamin C). 24  : Biologic Oxidation NAD+ - Nicotinamide Adenine Dinucleotide H H H CONH2 CONH2 +2H + H+ + -2H N N R R NAD+ NADH H + H3 C C O H + NAD+ H3 C C O + NADH + H H H NAD+ is a coenzyme in many enzymes - oxidizes different classes of compounds - removes hydrogen atoms from a C – H and O – H bonds, and two electrons from these bonds 25 Biologic Oxidation FAD FADH2 O O H H3 C N N H3 C NH + 2H , + 2e - NH H3 C - 2H , - 2e - H3 C C N N O N N O R H R FAD removes hydrogen atoms from two C – H bonds, and two electrons O O R CH CH C + FAD R CH CH C + FADH2 H H OH OH 26 Biologic Oxidation CoQ / CoQH2 - ubiquinone or coenzyme Q; exist in mitochondria in the oxidized form (aerobic conditions) and in the reduced form (anaerobic conditions) How does this redox pair work? O OH CH3O CH3 CH3O CH3 +2H -2H CH3O [CH2CH=CCH2]10H CH3O [CH2CH=CCH2]10H CH3 OH CH3 O oxidized form reduced form Biologic Oxidation Glutathione – tripeptide (glutamic acid, cysteine, glycine) HOOC-CH-CH2-CH2-CO-NH-CH-CO-NH-CH2-COOH NH2 CH2SH peptide bond Glutathion is abbreviated as G-SH -2H 2G-SH G-S-S-G +2H disulfide bridge (bond) 28 Biologic Oxidation Ascorbic acid (vitamin C) CH2OH CH2OH CHOH CHOH O O O -2H O H +2H H OH OH O O L-ascorbic acid L-dehydroascorbic acid 29 Biologic Oxidation L-cystein / L-cystine NH2 NH2 -2H S CH2 CH COOH 2 HS CH2 CH COOH +2H S CH2 CH COOH NH2 L - cysteine L - cystine Biologic Oxidation Lipoic acid / Dehydrolipoic acid COOH COOH (CH2)4 (CH2)4 -2H CH SH CH S CH2 +2H CH2 CH2 SH CH2 S Lipoic acid Dehydrolipoic acid SH S (L ) (L ) SH S  : Biologic Oxidation A catabolic reaction is a reaction that breaks macromolecules into constituent individual subunits. In a such reaction redox pairs are arranged in increase of their standard redox potential E0. An anabolic reaction is one that involves creating large macromolecules out of smaller molecules. Anabolic reactions are coupled with catabolic reactions. 32 Biologic Oxidation The mitochondria contain the series of catalysts known as respiratory chain that collect and transport H+ and e- and direct them to the final reaction with oxygen to form water. AH2 NAD+ FpH2 2 Fe 3+ H 2O substrate flavoprotein cytochromes A NADH Fp 2+ 1/2 O2 2Fe H+ H+ + 2H + 2H 33 Biologic Oxidation Problem Rearrange the following redox pairs in a catabolic reaction: NАD+ / NАDН + Н+ with E0 = –0,32 V; dehydrolipoic acid / lipoic acid with E0 = –0,29 V; Cytochrome c with E0 = 0,22 V; CоQ / CoQ.H2 with E0 = 0,10 V; and FАD / FАDН2 with E0 = –0,20 V. Answer: NАD+ / NАDН + Н+ dehydrolipoic acid / lipoic acid FАD / FАDН2 CоQ / CoQ.H2 Cytochrome c 34 Problems 1. What is the oxidation number? 2. Find the oxidation numbers of the species in the following compounds: H3PO4, CuSO4, NO2¯. 3. What is the oxidizing agent and what is the reducing agent in the reaction represented by the following equation? Sb2O3 + 2Fe → 2Sb + FeO3 4. Determine which substance is oxidized and which substance is reduced in the following reaction. Cu + Br2 → CuBr2 35  Problems 5. Whit regards of the standard redox potentials find out in which direction the equilibrium is shifted in a solution. Cu2+ + Cr2+ Cu+ + Cr 3+ Е0 Cu2+/Cu+ = + 0,153 V Е0 Cr3+/Cr 2+ = - 0,41 V? 6. Consider the following reaction for the metabolism of lactic acid. Is oxidation or reduction involved? Explain your answer. CH3-CH-COOH → CH3-CO-COOH OH 7. For the pair Cr6+/Cr3+ point out reduced form and oxidized form of chromium? Write the Nernst equation for this pair. 36  : Introduction to Oxidation- Reduction Reactions Terms Oxidation  Reduction Oxidation number Redox pair Rules for determining oxidation number Oxidazing agent Reducing agent  Disproportionation, comproportionation Biological oxidation and reduction - redox pairs Standard redox potential Nernst equation How to find direction of redox reaction Catabolic reaction Anabolic reaction 37