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Electrochemistry –An Oxidizing Agent Is a Species That Oxidizes Another Species; It Is Itself Reduced

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Electrochemistry –An Oxidizing Agent Is a Species That Oxidizes Another Species; It Is Itself Reduced Oxidation-Reduction Reactions Chapter 17 • Describing Oxidation-Reduction Reactions Electrochemistry –An oxidizing agent is a species that oxidizes another species; it is itself reduced. –A reducing agent is a species that reduces another species; it is itself oxidized. Loss of 2 e-1 oxidation reducing agent +2 +2 Fe( s) + Cu (aq) → Fe (aq) + Cu( s) oxidizing agent Gain of 2 e-1 reduction Skeleton Oxidation-Reduction Equations Electrochemistry ! Identify what species is being oxidized (this will be the “reducing agent”) ! Identify what species is being •The study of the interchange of reduced (this will be the “oxidizing agent”) chemical and electrical energy. ! What species result from the oxidation and reduction? ! Does the reaction occur in acidic or basic solution? 2+ - 3+ 2+ Fe (aq) + MnO4 (aq) 6 Fe (aq) + Mn (aq) Steps in Balancing Oxidation-Reduction Review of Terms Equations in Acidic solutions 1. Assign oxidation numbers to • oxidation-reduction (redox) each atom so that you know reaction: involves a transfer of what is oxidized and what is electrons from the reducing agent to reduced 2. Split the skeleton equation into the oxidizing agent. two half-reactions-one for the oxidation reaction (element • oxidation: loss of electrons increases in oxidation number) and one for the reduction (element decreases in oxidation • reduction: gain of electrons number) 2+ 3+ - 2+ Fe (aq) º Fe (aq) MnO4 (aq) º Mn (aq) 1 3. Complete and balance each half reaction Galvanic Cell a. Balance all atoms except O and H 2+ 3+ - 2+ (Voltaic Cell) Fe (aq) º Fe (aq) MnO4 (aq) º Mn (aq) b. Balance O atoms by adding H2O to one side of the equation 2+ 3+ - 2+ •A device in which chemical Fe (aq) º Fe (aq) MnO4 (aq) º Mn (aq) + 4 H2O c. Balance H atoms by adding H+ to energy is changed to electrical one side of the equation + - 2+ energy. 8 H (aq) + MnO4 (aq) º Mn (aq) + 4 H2O d. Balance the electric charge by adding electrons (e-) to the more positive side Fe2+ (aq) º Fe3+ (aq) + e- oxidation + - - 2+ 8 H (aq) + MnO4 (aq) + 5 e º Mn (aq) + 4 H2O reduction 4. Combine the two half-reactions to ob- Figure 17.1: Schematic of a method to tain the balanced oxidation-reduction separate the oxidizing and reducing agents of equation a redox reaction. (The solutions also contain a. Multiply each half reaction by a counterions to balance the charge.) factor which will allow the e-‘s to cancel when the equations are added 5 Fe2+ (aq) º 5 Fe3+ (aq) + 5e- oxidation + - - 2+ 8 H (aq) + MnO4 (aq) + 5 e º Mn (aq) + 4 H2O reduction b. Simplify the equation by canceling species which occur on both sides of the equation and reduce the coeffi- cients to the smallest whole number. Check 2+ + - 3+ 2+ 5 Fe (aq) + 8 H (aq) + MnO4 (aq) º 5 Fe (aq) + Mn (aq) + 4 H2O Additional Steps for Balancing Oxidation- Figure 17.2: Galvanic cells can Reduction Equations in Basic Solutions contain a salt bridge as in (a) or a 5. Note the number of H+ ions in the porous-disk connection as in (b). equation. Add this number of OH- ions to both sides of the equation 6. Simplify the equation by noting that H+ reacts with OH- to give H2O. Cancel any H2O’s that occur on both sides of the equation and reduce the equation to simplest terms 2 Figure 20.2: Atomic view of voltaic cell. Half-Reactions •The overall reaction is split into two half- reactions, one involving oxidation and one reduction. + − 2+ 2+ 3+ •8H + MnO4 + 5Fe → Mn + 5Fe + 4H2O + − − 2+ •Reduction: 8H + MnO4 + 5e → Mn + 4H2O •Oxidation: 5Fe2+ → 5Fe3+ + 5e− Anode and Cathode Figure 17.7: The schematic of a galvanic cell based on the half- reactions • OXIDATION occurs at the ANODE. • REDUCTION occurs at the CATHODE. Figure 17.3: An electrochemical process involves electron transfer at the interface between the electrode and the Cell Potential solution. (a) The species in the solution acting as the reducing agent supplies electrons to the anode. (b) The species in the solution acting as the oxidizing agent receives electrons from the cathode. •Cell Potential or Electromotive Force (emf): The “pull” or driving force on the electrons. 3 Figure 17.4: Digital voltmeters draw only a negligible current Notation for Galvanic Cells and are convenient to use. • It is convenient to have a shorthand way of designating particular Galvanic cells. • The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s) anodesalt bridge cathode • The cell terminals are at the extreme ends in the cell notation. Notation for Galvanic Cells Notation for Galvanic Cells • It is convenient to have a shorthand way of • It is convenient to have a shorthand way of designating particular Galvanic cells. designating particular Galvanic cells. • The cell consisting of the zinc-zinc ion half-cell • The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written and the copper-copper ion half-cell, is written Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s) Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s) anode cathode anodesalt bridge cathode • The anode (oxidation half-cell) is written on the • A single vertical bar indicates a phase left. The cathode (reduction half-cell) is written boundary, such as between a solid terminal and on the right. the electrode solution. Notation for Galvanic Cells Notation for Galvanic Cells • It is convenient to have a shorthand way of • When the half-reaction involves a gas, an designating particular Galvanic cells. inert material such as platinum serves as a • The cell consisting of the zinc-zinc ion half-cell terminal and an electrode surface on which and the copper-copper ion half-cell, is written the reaction occurs. 2+ 2+ • Figure 20.5 shows a hydrogen electrode; Zn(s) | Zn (aq) || Cu (aq) | Cu(s) hydrogen bubbles over a platinum plate anodesalt bridge cathode immersed in an acidic solution. • The two electrodes are connected by a salt • The cathode half-reaction is bridge, denoted by two vertical bars. + − 2H (aq) + 2e → H2 (g) 4 Figure 20.5: A hydrogen electrode. Notation for Galvanic Cells • To fully specify a voltaic cell, it is necessary to give the concentrations of solutions and the pressure of gases. • In the cell notation, these are written in parentheses. For example, 2+ + Zn(s) | Zn (1.0 M) || H (aq) | H2 (1.0 atm) | Pt Return to slide 50 Notation for Galvanic Cells A Problem To Consider • When the half-reaction involves a gas, an • Give the overall cell reaction for the Galvanic cell 2+ + inert material such as platinum serves as a Cd(s) | Cd (1.0 M) || H (aq) | H2 (1.0 atm) | Pt terminal and an electrode surface on which the reaction occurs. • The half-cell reactions are 2+ − • The notation for the hydrogen electrode, written Cd(s) → Cd (aq) + 2e as a cathode, is + − 2H (aq) + 2e → H2 (g) + H (aq) | H2 (g) | Pt + 2+ Cd(s) + 2H (aq) → Cd (aq) + H2 (g) Standard Cell emf’s and Notation for Galvanic Cells Standard Electrode Potentials • When the half-reaction involves a gas, an •A cell emf is a measure of the driving force inert material such as platinum serves as a of the cell reaction. terminal and an electrode surface on which •A reduction potential is a measure of the the reaction occurs. tendency to gain electrons in the reduction half- • To write such an electrode as an anode, you reaction. The oxidation state of the reactant simply reverse the notation. becomes less positive. •The oxidation potential for an oxidation half- + Pt | H2 (g) | H (aq) reaction is the negative of the reduction potential for the reverse reaction. 5 Standard Cell emf’s and Standard Electrode Potentials • By convention, the Table of Standard Electrode Potentials (Table 17.1) are tabulated as reduction potentials. • Consider the zinc-copper cell described earlier. Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s) • The two half-reactions are Zn(s) → Zn2+ (aq) + 2e− Cu2+ (aq) + 2e− → Cu(s) Standard Cell emf’s and Figure 17.6: A Standard Electrode Potentials galvanic cell involving the half-reactions • By convention, the Table of Standard Zn º Zn2++ 2e- Electrode Potentials (Table 17.1) are (anode) tabulated as reduction potentials. and • The copper half-reaction is a reduction.. 2+ Cu + 2e- º Cu • Write E for the electrode potential. (cathode), Cu with ξº cell 2+ − = 1.10V Cu (aq) + 2e → Cu(s) (ECu) Standard Cell emf’s and Standard Cell emf’s and Standard Electrode Potentials Standard Electrode Potentials • By convention, the Table of Standard • By convention, the Table of Standard Electrode Potentials (Table 17.1) are Electrode Potentials (Table 17.1) are tabulated as reduction potentials. tabulated as reduction potentials. • The zinc half-reaction is an oxidation. • For this cell, the cell emf is the sum of the reduction potential for the copper half-cell and • If you write EZn for the reduction potential of the oxidation potential for the zinc half-cell. zinc, then –EZn is the oxidation potential of zinc. 2+ − Zn (aq) + 2e → Zn(s) (EZn) Ecell = ECu + (−EZn ) 2+ − Zn(s) → Zn (aq) + 2e -(EZn) Ecell = ECu − EZn = (+.34 V) - (-0.76 V ) = 1.10 V 6 An electrochemical cell with a A Problem To Consider measured potential of 1.10 V. • Calculate the standard emf for the following Galvanic cell at 25 oC using standard electrode potentials.
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