Week 6.2 voltaic cells

Prepared by:Mrs Faraziehan Senusi

PA-A11-7C

David P. White Prentice Hall ©

2003

Electrochemical Cells

Corrosion & Prevention

Chapter 3Oxidation and Reduction

Oxidation-Reduction Concepts

Voltaic Cell

Electrolytic Cell

• Electrochemistry

– The study of the relationship between chemical change (reactions) and the flow of electrons (electrical work)

• Electrochemical Systems

– Voltaic/Galvanic – Release of free energy from a spontaneous reaction to produce electricity (Batteries)

– Electrolytic – Work done by absorbing free energy from a source (passage of an electrical current through a solution) to drive a nonspontaneous reaction

Electrochemistry

Voltaic (Galvanic) Cells

• Use spontaneous reaction (G < 0) to generate electrical energy

• Difference in Chemical Potential energy between higher energy reactants and lower energy products is converted to electrical energy to power electrical devices

• Thermodynamically - The system does work on the surroundings


Electrolytic Cells

• Uses electrical energy to drive nonspontaneous reaction (G > 0)

• Electrical energy from an external power supply converts lower energy reactants to higher energy products

• Thermodynamically – The surroundings do work on the system

• Examples – Electroplating and recovering metals from ores




• Zinc metal (Zn) in solution of Cu2+ ions

Construction of a Voltaic Cell

• The oxidizing agent (Zn) and reducing agent (Cu2+) in the same beaker will not generate electrical energy

• Separate the half-reactions by a barrier and connect them via an external circuit (wire)

• Set up salt bridge between chambers to maintain neutral charge in electrolyte solutions

2+ -Cu (aq) + 2e Cu(s) [reduction]

2+ -Zn(s) Zn (aq) + 2e [oxidation]

2+ 2+Zn(s) + Cu (aq) Zn + Cu(s)


v

Theodore L. Brown , H. Eugene Lemay , Bruce E. Bursten , Catherine J. Murphy ,David P. White, Chemistry the central science.

Zinc-Copper Voltaic Cell


Oxidation Half-Cell• Anode Compartment

– Oxidation of Zinc

• Zinc metal in solution of

Zn2+ electrolyte (ZnSO4)

• Zn is reactant in oxidation

half-reaction

• Conducts released electrons

(e-) out of its half-cell

Reduction Half-Cell• Cathode Compartment

- Reduction of Copper

• Copper bar in solution of

Cu2+ electrolyte (CuSO4)

• Copper metal is product in

reduction half-cell reaction

• Conducts electrons into its

half-cell

• Relative Charges on the Anode/Cathode electrodes

Electrode charges are determined by the source of the electrons and the direction of electron flow

Zinc atoms are oxidized (lose 2 e-) to form Zn2+ at the anode

• Anode – negative charge (e- rich)

Released electrons flow to right toward cathode to be accepted by Cu2+ to form Cu(s)

• Cathode – positive charge (e- deficient)


Purpose of Salt Bridge

• Electrons from oxidation of Zn leave neutral ZnSO4 solution producing net positive charge

• Incoming electrons to CuSO4 solution would produce net negative charge in solution as copper ions are reduced to copper metal

• Resulting charge imbalance would stop reaction

• Salt bridge provides “liquid wire” allowing ions to flow through both compartments completing circuit

• Salt bridge constructed of an inverted “U-tube” containing a solution of non-reacting Na+ & SO4

2- ions in a gel


Cell notation is used to describe the structure of a voltaic (galvanic) cell• For the Zn/Cu cell, the cell notation is:

Zn(s)Zn2+(aq) Cu2+(aq)Cu(s)

= phase boundary (solid Zn vs. Aqueous Zn2+)

= salt bridge


• Anode reaction (oxidation) is left of the salt bridge (AnOx)

• Cathode reaction (reduction) is right of the salt bridge (RedCat)

• Half-cell components usually appear in the same order as in the half-reactions (Zn(s) + 2e- Zn2+).

• Zinc solid loses 2e- (oxidized) to produce zinc(II) at the negative ANODE

• Copper(II) gains 2e- (reduced) to form copper metal at positive CATHODE

• Active vs Inactive Electrodes

Active Electrodes

• Electrodes in Zn/Cu2+ cell are active

• Zinc & Copper bars are components of the cell reactions

• Mass of Zn bar decreases as Zn2+ ions in cell solution increase

• Mass of Copper bar increases as Cu2+ ions accept electron to form more copper metal

Inactive Electrodes

• In many Redox reactions, one or the other reactant/product is not capable of serving as an electrode,so inactive electrodes are used.

• Inactive electrodes - Graphite or Platinum

– Can conduct electrons into and out of half-cells– Cannot take part in the half-reactions


Voltaic Cellwith

Inactive Graphite Electrodes

Each half-cell consists of inactive electrodes immersed in an electrolyte solution that contains all the reactant species involved in that half reaction.

~ Anode : I– ions oxidized to solid I2.~ e– released flow into graphite anode, through the wire, and into graphite cathode. ~ e– are consumed by MnO4

–, which reduced to Mn2

+ ions.

graphite I– (aq)I2 (s) H+(aq), MnO4

–(aq),Mn2+(aq)graphite

Write the cell notation for a voltaic cell with the following cell reaction

Practice Problem

2+ 2+Ni(s) + Pb (aq) Ni (aq) + Pb(s)

Ans:

A mercury battery, used for hearing aids and electric watches, delivers a constant voltage (1.35 V) for long periods. The half reactions are given below. Which half reaction occurs at the Anode and which occurs at the Cathode? What is the overall cell reaction?

HgO(s) + H2O(l) + 2e- Hg(l) + 2 OH-(aq)

Zn(s) + 2 OH-(aq) Zn(OH)2(s) + 2e-

Practice Problem

Summary

• A voltaic cell consists of oxidation (anode) and reduction (cathode) half-cells, connected by a wi re to conduct electrons and a salt bridge to maintain charge neutrality as the cell operates.

• Electrons move from anode (left) to cathode (right), while cations move from the salt bridge into the cathode half-cell and anions from the salt bridge into the anode half-cell.

• The cell notation shows the species and their phases in each half-cell, as well as the direction of current flow.

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Week 6.2 voltaic cells

Technology