Download Reduction and Oxidation

 Reduction and Oxidation Many chemical reactions include the transfer of electrons from one element to another. Electron gain is called reduction (decrease of oxidation number) and electron loss is oxidation (increase of oxidation number); the joint process is called a redox reaction. The species that supplies electrons is the reducing agent and the species that removes electrons is the oxidizing agent. Oxidation Number: Oxidation Number (ON) is a fictive charge obtained by exaggerating the ionic character of a bond. It can be considered as the charge that an atom would have if the more electronegative atom in a bond obtains the two electrons of the bond completely. When an element is ascribed a particular oxidation number, it is called the element is in a specific oxidation state. Oxidation numbers are given in positive or negative roman numbers. They are written in brackets or superscripted after the element symbol. Determination of oxidation numbers: ‐ draw Lewis structure and split the bonds heterolytically by using the concept of electronegativity ‐ determine ON by comparison of the valence electrons of a certain elements with the number of electrons after bond cleavage Rules of Oxidation Number: 1. Sum of ONs of all the atoms in a molecule = total charge (ON of Ca2+ =+II) 2. For atoms in their elemental form, the ON=0. (N2, Al, S8) 3. The ON of group 1 and 2 atoms have the number of their group in compounds (group 1‐I, group 2‐ II) 4. For atoms of group 13/ III (except B), the ON is + III (EX3) or + I (EX). And for atoms of group 14/ IV (except C, Si), the ON is + IV (EX4) or + II (EX2) 5. In compounds which are made up of non‐metals, the ON of the more electronegative element is equal to the charge of its anion. The other part of the compound has a positive oxidation number, and the sum of the ON in a neutral compound must be 0. 6. For hydrogen, + I in combination with nonmetals, and –I in combination with metals. 7. For fluorine, ‐ I in all its compound. 8. For oxygen, ‐ II except combined with F, ‐ I in peroxides(O22‐), ‐1/ 2 in superoxides(O2‐), ‐1/ 3 in ozonides(O3‐) 9. Halogens, ‐ I in most compounds, except the other elements include oxygen or more electronegative halogens. Determining Redox Equations: ‐ determine ON, find possible reaction products, define oxidation and reduction, count initial and final ON of each component and balance by adding electrons, equal the numberof elecrons for oxidation and reduction, combine both equations for final redox equation Example1: Determine the oxidation number of S in the following compounds. (a) H2SO4 ; (b) Na2S2O3; (c) K2S2O8; (d) SO32‐; (e) S4O62‐ Solution Set the ON of sulfur element x. (a) 2(+1)+1(x 1)+4(‐2)=0 x 1=+ VI (b) 2(+1)+2(x 2)+3(‐2)=0 x 2=+ II (c) 2(+1)+2(x 3)+8(‐2)=0 x 3=+ VII (d) 1(x 4)+3(‐2)=‐2 x 4=+ IV (e) 4(x 5)+6(‐2)=‐2 x 5=+5/ 2 Example2: xKMnO4 + yHCl ? Oxidation numbers: Mn(+VII), Cl(‐I), K(+I), O(‐II) , H(+I) Mn(+VII) is reduced to Mn2+(+II); Cl is oxidized (‐I ) to Cl2(0) xK MnO4 + yHCl aMnCl2 + b Cl2+ cH2O + dKCl Reduction: Mn(+VII) + 5e‐ Mn(+II) ×2 Oxidation: 2Cl(‐I) Cl2(each one 0) + 2 e‐ ×5 Redox: 2 Mn(+VII) +10 Cl(‐I) 2 Mn(+II) + 5 Cl2(0) 2K MnO4 + 16HCl 2MnCl2 + 5 Cl2+ 8H2O+ 2KCl Reference: D.F. Shriver, P. W. Atkins, Shriver & Atkins’ Inorganic Chemistry, 3th edition, 1999, 177,Oxford University Press, Oxford. D.F. Shriver, P. W. Atkins, Shriver & Atkins’ Inorganic Chemistry, 3th edition, 1999, 69,Oxford University Press, Oxford. Questions: 1. Complete the following reaction with the redox reaction method. (a) xCr2O72‐ + ySO32‐ + zH3O+ ? (b) x[Co(NH3)6]2+ + yO2 + zH2O ? 2. Determine the oxidation number of the element with a underline in the following compounds. HN3 NH2OH OF2 H3AsO4 BrO3‐ Ca3(PO4)2