Redox Titrations Lecture Notes PDF 2022-2023

‫اﻟﻣﺣﺎﺿرة اﻷوﻟﻰ‬ Reduction-Oxidation Titrations Dr. Mohamed Oraby 2022-2023  2 Oxidation-Reduction (Redox) Definitions & Terms Reaction of ferrous ion with chlorine gas 2 Fe2+ + Cl2 2Fe3+ +...

‫اﻟﻣﺣﺎﺿرة اﻷوﻟﻰ‬ Reduction-Oxidation Titrations Dr. Mohamed Oraby 2022-2023  2 Oxidation-Reduction (Redox) Definitions & Terms Reaction of ferrous ion with chlorine gas 2 Fe2+ + Cl2 2Fe3+ + 2Cl − 2Fe2+ - 2e 2Fe3+ (Loss of electrons: Oxidation) Cl2 + 2e 2 Cl- (Gain of electrons: Reduction) In every redox reaction, both reduction and oxidation must occur. Substance that gives electrons is the reducing agent or reductant. Substance that accepts electrons is the oxidizing agent or oxidant. Overall, the number of electrons lost in the oxidation half reaction must equal the number gained in the reduction half equation.  3 Antioxidants act by one or more of the following mechanisms:. Reaction with the generated free radicals can counteract the effect of the reactive oxygen and nitrogen species.. Binding the metal ions needed for catalyzing the formation of reactive radicals.. Repairing the damage or induction of the enzymes that catalyze the repair mechanisms.  4 Electrical properties of Redox systems Electrode potential: When a metal rod is dipped into a solution of its salts, there are tendencies for the following: Solution Pressure: The tendency of the metal to dissolve in a solution of its salt. Ionic Pressure: The tendency of the metal cations to deposit on its metal dipped into its solution. Cu/Cu2+ system: ionic pressure > solution pressure. Cu2+ leaves the solution to be deposited on Cu rod Zn/Zn2+ system: solution pressure > ionic pressure. Zn metal tends to dissolve forming Zn2+ in solution. Electrode potential (E): It is the potential difference between the metal rod (electrode) and the solution. It is a measure of the tendency of the metal to be oxidized to ions or the tendency of the ions to be reduced to metal atoms. 5 Electrical properties of redox systems Electrochemical cell:  6 Nernst Equation for Electrode Potential (E) o RT n+ Et = E + n log [M ] It is a quantitative relation of theF potential of half cell to molar conc. of ions in solution. Et = Electrode potential at temperature t. E° = Standard electrode potential (constant depend on the system) R = Gas constant T = Absolute Temp. (t°C + 273) F = Faraday constant (96500 Coulombs) Loge = ln (natural logarithm = 2.303 log) n = Valency of the ion [Mn+]=Molar concentration of metal ions in solution 7 Measurement of the Electrode Potential By connecting the electrode to another electrode (galvanic cell), an electric current will flow from the electrode having -ve potential to that having +ve potential (from Zn electrode to Cu electrode). The emf of the current can then be measured. The normal hydrogen electrode is used as a reference electrode. +0.33 9  8 Normal Hydrogen Electrode (NHE) A piece of platinum foil coated with platinum black and immersed in a solution of 1N HCl (with respect to H+). H2 H2 gas (at 1 atm. Pressure) is passed. a H+ = 1 Platinum black layer absorbs a large activity of amount of H2 and can be considered as hydrogen ions =1 a bar of hydrogen, it also catalyses the half reaction: 2 H+ + 2e → H2 Under these conditions: H2 electrode potential = zero  9 Standard Electrode Potential (Eo) Eo is the electromotive force (emf) produced when a half cell (consisting of the elements immersed in a molar solution of its ions) is coupled with a standard hydrogen electrode (E°= zero). System E° (volts) System E° (volts) Li / Li+ –3.03 Cd/Cd2+ –0.40 K / K+ –2.92 Sn / Sn2+ –0.13 Mg/Mg2+ –2.37 H2 (pt) / H+ 0.00 Al / Al3+ –1.33 Cu / Cu2+ +0.34 Zn / Zn2+ –0.76 Hg / Hg2+ +0.79 Fe / Fe2+ –0.44 Ag / Ag+ +0.80  10 Standard Oxidation Potential (Eo) the e.m.f. produced when a half cell consisting of an inert electrode (as platinum) dipped in a solution of equal concentration of both the oxidized and reduced forms (such as 3+ 2+ Fe / Fe 1. The higher ) is connected E°, the with a NHEthe oxidizing power of its stronger oxidant and the weaker the reducing power of its reduced form. 2. The most powerful oxidizing agents are those at the top (higher +ve) and the most powerful reducing agents are at the bottom (higher – ve). 3. 11 If any two redox systems are combined, the o stronger Standard Oxidation Potential (E ) oxidizing agent gains electrons from the stronger reducing agent with the formation of weaker reducing and oxidizing agents. Cl2/Cl- (E° = + 1.36 V) & Fe3+/Fe2+ (E°= + 0.77 V) Cl2 + 2Fe2+ → 2Fe3+ + 2Cl- strong oxid. + strong red. → weak oxid. + weak red. agent agent agent agent 12 Standard Oxidation Potential (Eo) at 25ºC System E° (volt) MnO4− / Mn+2 1.52 Ce+4 / Ce +3 1.44 Cl2 / 2Cl− 1.36 Cr2O7−2 / 2Cr+3 1.33 Br2 / 2Br− 1.07 Fe+3 / Fe+2 0.77 I2 / 2I− 0.54 2H+ / H2 0.00 AsO4−3 / AsO 3−3 −0.67