Voltaic or Galvanic Cells D8 c34 Electrochemical Cell

Electrolytic vs. Voltaic Cells How do voltaic and electrolytic cells differ? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrolytic vs. Voltaic Cells The process in which electrical energy is used to bring about a chemical change is called electrolysis. You are already familiar with some results of electrolysis, such as gold-plated jewelry, chrome-plated automobile parts, and silver-plated dishes. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The apparatus in which electrolysis is carried out is an electrolytic cell. An electrolytic cell uses electrical energy (direct current) to make a non-spontaneous redox reaction proceed to completion. An electrolytic cell is an electrochemical cell used to cause a chemical change through the application of electrical energy. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrolysis Using electrical energy to produce chemical change. Sn2+(aq) + 2 Cl-(aq) ---> Sn(s) + Cl2(g) SnCl2(aq) Cl2 Sn

Electrolytic vs. Voltaic Cells Anode (oxidation) Cathode (reduction) Energy Battery Voltaic Cell Electrolytic Cell e– In both voltaic and electrolytic cells, electrons flow from the anode to the cathode in the external circuit. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrolytic vs. Voltaic Cells Anode (oxidation) Cathode (reduction) Energy Battery Voltaic Cell Electrolytic Cell e– The key difference between voltaic and electrolytic cells is that in a voltaic cell, the flow of electrons is the result of a spontaneous redox reaction, whereas in an electrolytic cell, electrons are caused to flow by an outside power source, such as a battery. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrolytic vs. Voltaic Cells Anode (oxidation) Cathode (reduction) Energy Battery Voltaic Cell Electrolytic Cell e– In a voltaic cell, the anode is the negative electrode and the cathode is the positive electrode. In an electrolytic cell, the cathode is considered the negative electrode. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In electrolysis, an electric current is used to do which of the following? A. Cause a chemical change B. Produce a battery C. Generate heat D. Run a motor Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In electrolysis, an electric current is used to do which of the following? A. Cause a chemical change B. Produce a battery C. Generate heat D. Run a motor Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Driving Nonspontaneous Processes What are some applications that use electrolytic cells? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Driving Nonspontaneous Processes Electrolysis of a solution or of a melted, or molten, ionic compound can result in the separation of elements from compounds. Electrolytic cells are also commonly used in the plating, purifying, and refining of metals. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrolysis of water Anode (+) 2 H2O ---> O2(g) + 4 H+ + 4e- Cathode (-) 4 H2O + 4e- ---> 2H2 + 4 OH-

Electrolysis of Water The overall cell reaction is obtained by adding the half-reactions (after doubling the reduction half-reaction equation to balance electrons). Oxidation: 2H2O(l) → O2(g) + 4H+(aq) + 4e– Reduction: 2 [2H2O(l) + 2e– → H2(g) + 2OH–(aq)] 6H2O(l) → 2H2(g) + O2(g) + 4H+(aq) + 4OH–(aq) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Driving Nonspontaneous Processes Electrolysis of Molten Sodium Chloride The electrolytic cell in which this commercial process is carried out is called the Downs cell. The cell operates at a temperature of 801°C so that the sodium chloride is maintained in the molten state. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrolysis of Molten Sodium Chloride Sodium and chlorine are produced through the electrolysis of pure molten sodium chloride, rather than an aqueous solution of NaCl. Chlorine gas is produced at the anode. Molten sodium collects at the cathode. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrolysis of Molten Sodium Chloride The overall equation is the sum of the two half-reactions: Reduction: 2Na+(l) + 2e– → 2Na(l) 2NaCl(l) → 2Na(l) + Cl2(g) Oxidation: 2Cl–(l) → Cl2(g) + 2e– Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electroplating Electroplating is the deposition of a thin layer of a metal on an object in an electrolytic cell. An object that is to be silver-plated is made the cathode in an electrolytic cell. The anode is the metallic silver that is to be deposited. The electrolyte is a solution of a silver salt, such as silver cyanide. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electroplating When direct current is applied, silver ions move from the anode to the object to be plated. Reduction: Ag+(aq) + e– → Ag(s) (at cathode) This figure shows statuettes that were electroplated with copper, nickel, and 24-carat gold. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

This technique is used to obtain ultrapure silver, lead, and copper. Electrorefining In the process of electrorefining, a piece of impure metal is made into the anode of the cell. It is oxidized to the cation and then reduced to the pure metal at the cathode. This technique is used to obtain ultrapure silver, lead, and copper. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electrowinning In a process called electrowinning, impure metals can be purified in electrolytic cells. The cations of molten salts or aqueous solutions are reduced at the cathode to give very pure metals. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Producing Aluminum 2 Al2O3 + 3 C ---> 4 Al + 3 CO2 Charles Hall (1863-1914) developed electrolysis process. Founded Alcoa.

Key Concepts The key difference between voltaic and electrolytic cells is that in a voltaic cell, the flow of electrons is the result of a spontaneous redox reaction, whereas in an electrolytic cell, electrons are caused to flow by an outside power source, such as a battery. Electrolysis of a solution or of a melted, or molten, ionic compound can result in the separation of elements from compounds. Electrolytic cells are also commonly used in the plating, purifying, and refining of metals. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Glossary Terms electrolysis: a process in which electrical energy is used to bring about a chemical change; the electrolysis of water produces hydrogen and oxygen electrolytic cell: an electrochemical cell used to cause a chemical change through the application of electrical energy Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

BIG IDEA Matter and Energy The two types of electrochemical cells are voltaic cells and electrolytic cells. In an electrolytic cell, a nonspontaneous redox reaction is driven by the application of electrical energy. Electrolytic cells are used to produce commercially important chemicals and to plate, purify, and refine metals. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Quantitative Aspects of Electrochemistry electrolysis of aqueous silver ion. Ag+ (aq) + e- ---> Ag(s) 1 mol e- ---> 1 mol Ag If measure the moles of e-, can know the quantity of Ag formed. But how to measure moles of e-?

Quantitative Aspects of Electrochemistry Measure the electrical current Units

Quantitative Aspects of Electrochemistry But how is charge related to moles of electrons? = 96,500 C/mol e- = 1 Faraday

Michael Faraday 1791-1867 Originated the terms anode, cathode, anion, cation, electrode. Discoverer of electrolysis magnetic props. of matter electromagnetic induction benzene and other organic chemicals

Quantitative Aspects of Electrochemistry 1.50 A flow thru a Ag+(aq) solution for 15.0 min. What mass of Ag metal is deposited? Solution: Calc. charge Charge (C) = current (A) x time (t) = (1.5 A)(15.0 min)(60 s/min) = 1350 C

Quantitative Aspects of Electrochemistry Charge = 1350 C NEXT: Calculate moles of e- used (c) Calc. quantity of Ag

Pb(s) + HSO4-(aq) ---> PbSO4(s) + H+(aq) + 2e- If a battery delivers 1.50 A, and you have 454 g of Pb, how long will the battery last? Solution a) 454 g Pb = 2.19 mol Pb b) Calculate moles of e- c) Calculate charge 4.38 mol e- • 96,500 C/mol e- = 423,000 C

NEXT: Use Charge = 423,000 C d) Calculate time About 78 hours