Utilizes relationship between chemical potential energy & electrical energy

Redox Reactions battery to start car prevent corrosion cleaning with bleach (oxidizing agent) Na, Al, Cl prepared or purified by redox reactions breathing O2  H2O and CO2

Redox Reactions Synthesis Decomposition Single Replacement Redox rxns DR rxns NOT redox rxn! Redox rxns

Predicting Redox Reactions Table J: predict if given redox reaction will occur metals donate electrons to ion of metals below itself nonmetal steals electrons from ion of nonmetal below itself

Predicting Single Replacement Redox Reactions Element + Compound  New Element + New Compound If element above swapable ion, reaction is spontaneous If element below swapable ion, reaction is not spontaneous

Predicting Redox Reactions A + BX  B + AX If metal A above metal B (Table J): reaction is spontaneous X + AY  Y + AX If nonmetal X above nonmetal Y (Table J): reaction is spontaneous

Spontaneous or not? Yes Cs + CuCl2  I2 + NaCl  Yes Cl2 + KBr  Li + AlCl3  Cs + CuCl2  I2 + NaCl  Cl2 + KBr  Fe + CaBr2  Mg + Sr(NO3)2  F2 + MgCl2  Yes Yes No Yes No No Yes

Which beaker had Zn ions & which had Ag ions? Started with: 1. Zn(NO3)2 & Cu 2. AgNO3 & Cu Which beaker had Zn ions & which had Ag ions?

Overview of Electrochemistry TWO kinds of cells: 1. Galvanic or Voltaic (NYS – Electrochemical) Use spontaneous rxn to produce flow of electrons (electricity) = Exothermic 2. Electrolytic Use flow of electrons (electricity) to force nonspontaneous rxn to occur = Endothermic

Vocabulary Redox Half-reaction Oxidation Reduction Cell Half-Cell Electrode Anode Cathode Galvanic Voltaic Electrochemical Electrolytic Salt bridge

Electrochemical Cells spontaneous SR redox rxn: produces flow of electrons Electrons flow from oxidized substance to reduced substance Names: Galvanic cells, voltaic cells, or electrochemical cells (NYS)

Electrochemical Cells Redox rxn arranged so electrons forced to flow through wire When electrons travel through a wire, can make them do work - light a bulb,ring a buzzer oxidation & reduction reactions must be separated physically

Half-Cell Place where each half-reaction takes place ½ cells: 2 needed for complete redox rxn connected by wire so electrons flow through connected by salt bridge to maintain electrical neutrality

Schematic of Galvanic/Voltaic Cell

Parts of a Galvanic/Voltaic Cell 2 half-cells: One for oxidation rxn & One for reduction rxn Each consists of: container with aqueous solution & electrode (surface where electron transfer takes place) Wire connects electrodes Salt bridge connects solutions

How much work can you get out of this reaction? can measure voltage by allowing electrons to travel through voltmeter galvanic cell is a battery not easy battery to transport or use in real-life applications

Electrode Surface at which oxidation or reduction half-reaction occurs: Anode & Cathode

An Ox Ate a Red Cat Anode – Oxidation the anode = location for the oxidation half-reaction Reduction – Cathode the cathode = location for the reduction half-reaction

Anode / Cathode How know which electrode is which? Table J: predict which electrode anode and which electrode is cathode

Anode Anode = Oxidation = Electron Donor anode is metal higher on Table J

Cathode Cathode = Reduction = Electron Acceptor cathode is metal lower on Table J

Zn is above Cu, Zn is anode

Direction of Electron Flow (through wire): Anode → Cathode Direction of Positive Ion Flow (salt bridge): Anode → Cathode

Positive & Negative Electrode Negative electrode (anode): where electrons originate here it’s Zn electrode Positive electrode (cathode): electrode that attracts electrons here it’s the Cu electrode

Aqueous Solution Solution containing ions of same element as electrode Cu electrode: Solution: Cu(NO3)3 or CuSO4 Zn electrode: Solution: Zn(NO3)2 or ZnSO4

Salt Bridge migration of ions between half-cells necessary to maintain electrical neutrality reaction can not proceed without salt bridge

A(s) + BX(aq)  B(s) + AX(aq) SR rxn occurs during operation of galvanic/voltaic cell One electrode gains mass (B) and one electrode dissolves (A) concen of metal ions ↑ in one soln (making AX) & ↓ in other soln (using up BX)

Half-Reactions Zn  Zn+2 + 2e- Cu+2 + 2e-  Cu _________________________ Zn + Cu+2  Zn+2 + Cu Which electrode is dissolving? Which species is increasing its mass? Zn Zn+2

Zn + Cu+2  Zn+2 + Cu Cu Cu+2 Which electrode is gaining mass? Which species is getting more dilute? Cu+2

When the reaction reaches equilibrium voltage is 0! electrons no longer flow

Construct Galvanic Cell with Al & Pb Use Table J to identify anode & cathode Draw Cell: put in electrodes & solutions Label: anode, cathode, direction of electron flow in wire, direction of positive ion flow in salt bridge, positive electrode, negative electrode Negative electrode: where electrons originate Positive electrode: attracts electrons

(-)  Electron flow  wire Positive ion flow  Pb = cathode Al = anode Salt bridge (-)  Pb+2 & NO3-1 Al+3 & NO3-1

What are half-reactions? Al  Al+3 + 3e- Pb+2 + 2e-  Pb Al metal is electrode that’s dissolving Al+3 ions go into solution Pb+2 ions are in the solution Ions pick up 2 electrons & plate together on surface of Pb electrode as Pb0

Overall Rxn 2(Al  Al+3 + 3e-) + 3(Pb+2 + 2e-  Pb) _____________________________ 2Al + 3Pb+2 + 6e- 2Al+3 + 3Pb + 6e- 2Al + 3Pb+2  2Al+3 + 3Pb

2Al + 3Pb+2  2Al+3 + 3Pb Al Pb Increasing Decreasing Which electrode is losing mass? Which electrode is gaining mass? What’s happening to the [Al+3]? What’s happening to the [Pb+2]? Pb Increasing Decreasing

Application: Batteries

Dry Cell

Mercury battery

Application: Corrosion

Corrosion Prevention

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