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Topic 23 Table of Contents Topic 23 Topic 23: Electrochemistry Basic Concepts Additional Concepts Electrochemistry: Basic Concepts Topic 23 Reviewing Redox • Suppose you could separate the oxidation and reduction parts of a redox reaction and cause the electrons to flow through a wire. • The flow of electrons in a particular direction is called an electrical current. Electrochemistry: Basic Concepts Topic 23 Reviewing Redox • In other words, you are using a redox reaction to produce an electrical current. • This is what occurs in a battery—one form of an electrochemical cell in which chemical energy is converted to electrical energy. Electrochemistry: Basic Concepts Topic 23 Reviewing Redox • You can reverse the process and use a current to cause a redox reaction to occur. Electrochemistry: Basic Concepts Topic 23 • • • • Electrolysis An electrochemical cell consists of two electrodes and a liquid electrolyte. One electrode, the cathode, brings electrons to the chemically reacting ions or atoms in the liquid; the other electrode, the anode, takes electrons away. The electrons act as chemical reagents at the electrode surface. The liquid electrolyte acts as the chemical reaction medium. Electrochemistry: Basic Concepts Topic 23 Electrolysis • You can remember that reduction always occurs at the cathode and oxidation always occurs at the anode by studying this diagram. Electrochemistry: Basic Concepts Topic 23 The Electrolysis Process • Electrolysis takes place in a type of electrochemical cell called an electrolytic cell, in which a source of electricity, such as a battery, is added to an external circuit connecting the electrodes. • The electrolysis process occurs when the electrons are transferred between the electronic conductors—the metal electrodes—and the ions or atoms at the electrode surfaces. Electrochemistry: Basic Concepts Topic 23 The Electrolysis Process Click box to view movie clip. Electrochemistry: Basic Concepts Topic 23 Electrolytic Cell • Electrolysis, the splitting of compounds by electricity, occurs when two electrodes, an anode and a cathode, are inserted into a liquid electrolyte such as molten sodium chloride and connected to a source of electrical energy such as a battery. Electrochemistry: Basic Concepts Topic 23 Electrolytic Cell Electrochemistry: Basic Concepts Topic 23 Electroplating • Reduction of silver ions onto cheaper metals forms silverplate. Click box to view movie clip. Electrochemistry: Basic Concepts Topic 23 Electroplating • The object to be plated is made the cathode. • At the pure silver anode, oxidation of silver metal to silver ions replaces the silver ions removed from the solution by plating at the cathode. Electrochemistry: Basic Concepts Topic 23 • • • • Electrolytic Cleaning Electrolysis can be used to clean objects by pulling ionic dirt away from them. The electrolysis cell for this cleaning process includes a cathode that is the object itself, a stainless steel anode, and an alkaline electrolyte. When an electric current is run through the cell, the chloride ions are drawn out. Hydrogen gas forms and bubbles out, helping to loosen corrosion products. Electrochemistry: Basic Concepts Topic 23 Electrophoresis • Electrophoresis is another electrochemical process that was used to restore some of the ceramic and organic artifacts from the Titanic. • Electrophoresis involves placing an artifact in an electrolyte solution between positive and negative electrodes and applying a current. Electrochemistry: Basic Concepts Topic 23 Electrophoresis • The current breaks up salts, dirt, and other particles as their charged components migrate to the electrodes. • Electrophoresis is also used in laboratories to separate and identify large molecules. Electrochemistry: Basic Concepts Topic 23 Potential Difference • Why do the electrons travel in one direction and not in the reverse? • The electron pressure at the cathode is kept low by the reduction reaction, and the electrons flow from a region of high pressure (negative potential at the anode) to a region of low pressure (positive potential at the cathode). This potential difference between the electrodes causes an electrical current to flow. Electrochemistry: Basic Concepts Topic 23 Potential Difference • In this model of a lemon battery, the level of the electron sea is raised or lowered by the chemical reactions at the electrode surfaces, creating a potential difference across the battery. Electrochemistry: Basic Concepts Topic 23 Potential Difference • A spontaneous oxidation reaction raises the electron pressure (potential) at the anode, and a spontaneous reduction reaction reduces the pressure at the cathode. Electrochemistry: Basic Concepts Topic 23 Potential Difference Electrochemistry: Basic Concepts Topic 23 Potential Difference • Because the redox reactions that take place during electrolysis are not spontaneous, a battery is needed to pump electrons from an area of low potential to one of high potential. Click box to view movie clip. Electrochemistry: Basic Concepts Topic 23 Potential Difference Electrochemistry: Basic Concepts Topic 23 • • • • Galvanic Cells An electrochemical cell in which an oxidation-reduction reaction occurs spontaneously to produce a potential difference is called a galvanic cell. In a Galvanic cell, chemical energy is converted into electrical energy. Galvanic cells are sometimes called voltaic cells; both terms refer to the same device. A galvanic cell that has been packaged as a portable power source is often called a battery. Electrochemistry: Basic Concepts Topic 23 Batteries Perform Work • When a simple galvanic cell does useful work, it is called a battery. Click box to view movie clip. Electrochemistry: Basic Concepts Topic 23 Batteries Perform Work • If the external circuit is connected with a wire, electrons flow from the site of oxidation at the magnesium strip and through the LED to the surface of the copper strip, where reduction of Cu2+ ions takes place. Electrochemistry: Basic Concepts Topic 23 Batteries Perform Work • The voltage pushes electrons through the LED, causing it to light up. Electrochemistry: Basic Concepts Topic 23 Batteries Perform Work Electrochemistry: Basic Concepts Topic 23 Modern Batteries • Modern batteries come in a wide variety of sizes, shapes, and strengths. • Each type of battery serves a different purpose. Electrochemistry: Basic Concepts Topic 23 Modern Batteries Electrochemistry: Basic Concepts Topic 23 Modern Batteries • Although the term battery usually refers to a series of galvanic cells connected together, some batteries have only one such cell. • Other batteries may have a dozen or more cells. Electrochemistry: Basic Concepts Topic 23 Modern Batteries • When you put a battery into a flashlight, radio, or CD player, you complete the electrical circuit of a galvanic cell(s), providing a path for the electrons to flow through as they move from the reducing agent (the site of oxidation) to the oxidizing agent the site of (the reduction). Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • Whenever you put two or more common D batteries into a flashlight, you are connecting them in series. • They have to be placed in the correct order so that electrons flow through both cells. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • These relatively inexpensive batteries are carbon-zinc galvanic cells, and they come in several types, including standard, heavy-duty, and alkaline. • This type of battery is often called a dry cell because there is no aqueous electrolyte solution; a semisolid paste serves that role. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • A standard D battery is shown both whole and cut in half to reveal the structure of the carbon-zinc dry cell. • Beneath the outside paper cover of the battery is a cylinder casing made of zinc. • The zinc serves as the anode and will be oxidized in the redox reaction. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • The carbon rod in the center of the cylinder— surrounded by a moist, black paste of manganese (IV) oxide (MnO2) and carbon black—acts as a cathode. • Ammonium chloride (NH4Cl) and zinc chloride (ZnCl2) serve as electrolytes. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • Alkaline batteries contain potassium hydroxide (KOH) in place of the ammonium chloride electrolyte, and they maintain a high voltage for a longer period of time. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • The flow of electrons from the zinc cylinder through the electrical circuits of an appliance and back into the battery provides the electricity needed to power a flashlight, radio, CD player, toy, clock, or other item. • When electrons leave the casing, zinc metal is oxidized. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • The reactions in the carbon rod and the paste are much more complex, but one major reduction that takes place is that of manganese in manganese (IV) oxide. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • In this reaction, the oxidation number of manganese is reduced from . • Adding the two half-reactions together gives the major redox reaction taking place in a carbon-zinc dry cell. Electrochemistry: Basic Concepts Topic 23 • • • • Carbon-Zinc Dry Cell Each galvanic cell in a lead-acid battery has two electrodes—one made of a lead (IV) oxide (PbO2) plate and the other of spongy lead metal. In each cell, lead metal is oxidized as lead (IV) oxide is reduced. The lead metal is oxidized to Pb2+ ions as it releases two electrons at the anode. The Pb4+ ions in lead oxide gain two electrons, forming Pb2+ ions at the cathode. Electrochemistry: Basic Concepts Topic 23 Carbon-Zinc Dry Cell • The Pb2+ ions combine with SO42– ions from the dissociated sulfuric acid in the electrolyte solution to form lead (II) sulfate at each electrode. • Thus, the net reaction that takes place when a lead-acid battery is discharged results in the formation of lead sulfate at both of the electrodes. Electrochemistry: Basic Concepts Topic 23 Lead Storage Batteries Electrochemistry: Basic Concepts Topic 23 Lead Storage Batteries • The reaction that occurs during discharge of a lead-acid battery is spontaneous and requires no energy input. • The reverse reaction, which recharges the battery, is not spontaneous and requires an input of electricity from the car’s alternator. Electrochemistry: Basic Concepts Topic 23 Lead Storage Batteries • Current enters the battery and provides energy for the reaction in which lead sulfate and water are converted into lead (IV) oxide, lead metal, and sulfuric acid. Electrochemistry: Basic Concepts Topic 23 Experimental Batteries • Two new experimental types of batteries for use in electric cars show early promise as candidates. • One is a rechargeable, nickel-metal hydride or NiMH battery. • This type of battery is less toxic and has a higher storage capacity than the batteries now used in electric cars. Electrochemistry: Basic Concepts Topic 23 Experimental Batteries • Another experimental battery is a lithium battery with a water-based electrolyte. • Lithium is more easily oxidized than any other metal but has a drawback that has limited its use in batteries: it explodes violently when it comes into contact with water. • Lithium is used in some batteries to power camcorders, but they require an expensive, nonaqueous electrolyte. Electrochemistry: Basic Concepts Topic 23 Aqueous Lithium Battery • How can a lithium battery have an aqueous electrolyte? Two facets of the construction of this new battery keep the lithium metal from reacting with water. • First, the lithium is in the form of individual atoms embedded in a material such as manganese (IV) oxide, rather than as a solid metal. Electrochemistry: Basic Concepts Topic 23 Aqueous Lithium Battery • Second, the electrolyte is full of dissolved lithium salts, so the lithium ions that are produced travel to the site of reduction without reacting with water. Electrochemistry: Basic Concepts Topic 23 Aqueous Lithium Battery Basic Assessment Questions Topic 23 Question 1 What term describes a battery that is not rechargeable? Basic Assessment Questions Topic 23 Answer primary battery Basic Assessment Questions Topic 23 Question 2 What element is oxidized in most dry cells? Basic Assessment Questions Topic 23 zinc Answer Basic Assessment Questions Topic 23 Question 3 Distinguish between a voltaic cell and an electrolytic cell. Basic Assessment Questions Topic 23 Answer spontaneous redox reaction; nonspontaneous redox reaction by electrolysis Electrochemistry: Additional Concepts Topic 23 Additional Concepts Electrochemistry: Additional Concepts Topic 23 Calculating Cell Potential • The two reduction half-reactions in this example represent the half-cells of a voltaic cell. • The standard reduction potentials for each half-reaction are given. Electrochemistry: Additional Concepts Topic 23 • • • • • Calculating Cell Potential Determine the overall cell reaction and the standard cell potential. Write the cell chemistry using cell notation with vertical lines separating components. Note that reduction of iodine has the higher reduction potential. This half-reaction will proceed in the forward direction as a reduction. The iron half-reaction will proceed in the reverse direction as an oxidation. Electrochemistry: Additional Concepts Topic 23 Calculating Cell Potential • Rewrite the half-reactions in the correct direction. Electrochemistry: Additional Concepts Topic 23 Calculating Cell Potential • Balance the reaction if necessary. Note that this reaction is balanced as written. • Calculate cell standard potential. Electrochemistry: Additional Concepts Topic 23 Calculating Cell Potential • Write the reaction using cell notation. • When representing a reaction in cell notation, the species in the oxidation halfreaction are written first in the following order: or in this case, . Electrochemistry: Additional Concepts Topic 23 Calculating Cell Potential • The species in the reduction half-reaction are written next in the order or in this case, . • Therefore, the complete cell is represented as . Electrochemistry: Additional Concepts Topic 23 Predicting the Spontaneity of a Reaction • Predict whether the following redox reaction will occur spontaneously. Electrochemistry: Additional Concepts Topic 23 Predicting the Spontaneity of a Reaction • Write the half-reactions. Note that the coefficients are simplified. • Find the standard cell potential, using E0 values from Table 21-1 in your textbook. Electrochemistry: Additional Concepts Topic 23 Predicting the Spontaneity of a Reaction • The voltage is negative, so the reaction is not spontaneous. • The reverse reaction will occur spontaneously. Additional Assessment Questions Topic 23 Question 1 For each pair of half-reactions, write the balanced equation for the overall cell reaction, calculate the standard cell potential, and express the reaction using cell notation. Use E0 values from Table 21-1 in your textbook. Mg2+(aq) + 2e– → Mg(s) Pd2+(aq) + 2e– → Pd(s) Additional Assessment Questions Topic 23 Answer Additional Assessment Questions Topic 23 Question 2 Calculate the cell potential to determine if each of these redox reactions is spontaneous. Additional Assessment Questions Topic 23 Question 2a 2Ag+(aq) + Co(s) → Co2+(aq) + 2Ag(s) Answer 2a + 1.08V; spontaneous Additional Assessment Questions Topic 23 Question 2b Cu(s) + Cu2+(aq) → 2Cu+(aq) Answer 2b – 0.368V; not spontaneous Help To advance to the next item or next page click on any of the following keys: mouse, space bar, enter, down or forward arrow. Click on this icon to return to the table of contents Click on this icon to return to the previous slide Click on this icon to move to the next slide Click on this icon to open the resources file. Click on this icon to go to the end of the presentation. End of Topic Summary File