Chemistry: The Central Science. Chapter 20: Electrochemistry

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Gerald McCoy
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1 Chemistry: The Central Science Chapter 20: Electrochemistry Redox reaction power batteries Electrochemistry is the study of the relationships between electricity and chemical reactions o It includes the study of both spontaneous and nonspontaneous processes 20.1: Oxidation States and Oxidation-Reduction Reactions Oxidation numbers of all the elements involved in the reaction can be tracked to determine whether the reaction is a redox reaction In some reactions, the oxidation numbers change, but we cannot say that any substance literally gains or loses electron o E.g. Combustion of hydrogen to form water In this reaction, hydrogen is oxidized from 0 to +1 oxidation state and oxygen is reduced from the 0 to the -2 oxidation state Water is not an ionic substance, however, and so there is not a complete transfer of electrons from hydrogen to oxygen as water is formed o Using oxidation states is a convenient form of bookkeeping, but you should not generally equate the oxidation state of an atom with its actual charge in a chemical compound The substance the oxidizes the other substance (thus becoming reduced) is called the oxidizing agent or oxidant The substance that reduces the other substance (thus becoming oxidized) is called the reducing agent or reductant 20.2: Balancing Oxidation-Reduction Equations When balancing a redox reaction, the gains and the losses of electrons must be balanced Half-Reactions o Although oxidation and reduction must take place simultaneously, it is often convenient to consider them as separate processes o E.g. Oxidation: Reduction:

2 Equations that show either oxidation or reduction alone are called halfreactions In the overall reaction, the number of electrons lost in the oxidation half-reaction must equal to the number of electrons gained in the reduction half-reaction Balancing Equations by the Method of Half-Reactions o The use of half-reactions to balance oxidation-reduction equations usually begin with a skeleton ionic equation that shows only the substances undergoing oxidation and reduction o For balancing a redox reaction that occurs in acidic aqueous solution, the procedure is as follows: Divide the equation into two half-reactions, one for oxidation and the other for reduction Balance each half-reaction First, balance the elements other than H and O Next, balance the O atoms by adding H 2 O as needed Then, Balance the H atoms by adding H + as needed Finally, balance the charge by adding e - as needed At this point, you can check whether the number of electrons in each half-reaction equals corresponds to the changes in oxidation state Multiply the half-reactions by integers, if necessary, so that the number of electrons lost in one half-reaction equals the number of electrons gained in the other Add the two half-reactions and, if possible, simplify by canceling species appearing on both sides of the combined equation Check to make sure that atoms and charges are balanced Balancing Equations for Reactions Occurring in Basic Solution o One way to balance these reactions is to balance the half-reactions initially as if they occurred in acidic solution Then, count the H + in each half-reaction, and add the same number of OH - to each side of the half-reaction The OH - will neutralize the protons on the side containing H + and the other side ends up with OH : Voltaic Cells The energy released in a spontaneous redox reaction can be used to perform electrical work

3 o This task is accomplished through a voltaic (or galvanic) cell, a device in which the transfer of electrons takes place through an external pathway E.g. spontaneous reaction occurs when a strip of zinc is placed in contact with a solution containing Cu 2+ Zn metal is in contact with Zn 2+ (aq) in one compartment of the cell, and Cu metal is in contact with Cu 2+ (aq) in another compartment o Consequently, the reduction of the Cu 2+ can only occur by a flow of electrons through an external circuit Two solid metals that are connected by the external circuit are called electrodes o Electrode at which oxidation occurs is called the anode o Electrode at which reduction occurs is called the cathode Each compartments of a voltaic cell is called a half-cell o Anode: o Cathode: For a voltaic cell to work, the solutions in the two half-cells must remain electrically neutral As Zn is oxidized in the anode compartment, Zn 2+ enter the solution As Cu 2+ at the cathode reduces, the positive charge from the solution is removed A salt bridge serves this purpose

o A salt bridge consists of a U-shaped tube that contains an electrolyte, such as NaNO 3 (aq), whose ions will not react with other ions in the cell or with the electrode materials Anions always

4 o A salt bridge consists of a U-shaped tube that contains an electrolyte, such as NaNO 3 (aq), whose ions will not react with other ions in the cell or with the electrode materials Anions always migrate toward the anode and the cations toward the cathode In any voltaic cell the electron flow from the anode through the external circuit to the cathode A Molecular View of Electrode Processes o Redox reaction between Zn(s) and Cu 2+ (aq) lead to an increase in Zn 2+ (aq) and Cu, and a decrease in Zn(s) and Cu 2+ (aq) o In the case of the voltaic cell, the Zn atom loses two electrons and becomes a Zn 2+ (aq) in its compartment The electron travels through the wire and attached to Cu 2+ (aq), forming Cu(s) in its compartment o The redox reaction between Zn and Cu 2+ is spontaneous regardless of whether they react directly or in the separate compartments of a voltaic cell 20.4: Cell EMF under Standard Conditions The electrons flow spontaneously toward the electrode with the more positive electrical potential

The difference in potential energy per electrical charge (the potential difference) between two electrodes is measured in units of volts o C is coulomb V is volt One electron has a charge of 1.

5 The difference in potential energy per electrical charge (the potential difference) between two electrodes is measured in units of volts o C is coulomb V is volt One electron has a charge of C Electromotive force (emf) The potential difference between the two electrodes of a voltaic cell providing the driving force that pushes the electron through the circuit o The emf of a cell, denoted E cell, is also called the cell potential E cell is measured in volts so it s often referred to as cell voltage o For any cell reaction that proceeds spontaneously such as that in a voltaic cell, the cell potential will be positive Under standard conditions (25 C and 1 M for aqueous or 1 atm for gases), the emf is called the standard emf, or the standard cell potential, and is denoted E cell Standard Reduction (Half-Cell) Potentials o Standard reduction potentials (E red ) the standard electrode potentials tabulated for reduction reactions o For all spontaneous reactions at standard conditions, E cell > 0 o The reference half-reaction is the reduction of H + (aq) to H 2 (g) under standard conditions, which is assigned a standard reduction potential of exactly 0 V An electrode designed to produced this half-reaction is called a standard hydrogen electron (SHE), or the normal hydrogen electrode (NHE) An SHE consists of a platinum wire connected to a piece of platinum foil covered with finely divided platinum that serves as an inert surface for the reaction The electron is encased in a glass tube so that the hydrogen gas under standard conditions (1 atm) can bubble over the platinum The solution contains H + (aq) under standard (1 M) conditions

o Whenever we assign an electrical potential to a half-reaction, we write the reaction as a reduction o Changing the stoichiometric coefficient in a half-reaction does not affect the value of the

6 o Whenever we assign an electrical potential to a half-reaction, we write the reaction as a reduction o Changing the stoichiometric coefficient in a half-reaction does not affect the value of the standard reduction potential o The more positive the value of E red, the greater the driving force for reduction under standard conditions Strength of Oxidizing and Reducing Agents o The more positive the E red value for a half-reaction, the greater the tendency for the reactant of the half-reaction to be reduced and oxidize another species o The half-reaction with the smallest reduction potential is most easily reversed as an oxidation o Solutions of reducing agents are difficult to store for extended periods because of the ubiquitous presence of O 2, a good oxidizing agent 20.5: Free Energy and Redox Reactions

7 E: A positive value of E indicates a spontaneous process; a negative value of E indicates a nonspontaneous one The activity series consists of the oxidation reactions of the metals, ordered from the strongest reducing agent at the top to the weakest reducing agent at the bottom o E.g. Ni is oxidized and Ag + is reduced Positive value of E indicates that the displacement of silver by nickel is a spontaneous process EMF and ΔG o ΔG is the change in Gibbs free energy n is a positive number without units that represents the number of electrons transferred in the reaction F is called Faraday s constant, which is the quantity of electrical charge on one mole of electrons o A positive value of E and a negative value of ΔG both indicate that a reaction is spontaneous o When the reactants and products are all in their standard states 20.6 Cell EMF under Nonstandard Conditions As a voltaic cell is discharged, the reactants of the reaction are consumed, and the products are generated, so the concentrations of these substances change o The emf progressively drops until E = 0, at which point we say the cell is dead The Nernst Equation o This is the Nernst equation At T = 298 K the quantity RT/F equals , with the units of volts (V) o

o At E = 0, ΔG = 0 The system is at equilibrium o In general, increasing the concentration of reactants of decreasing the concentration of products increases the driving force for the reaction,

8 o At E = 0, ΔG = 0 The system is at equilibrium o In general, increasing the concentration of reactants of decreasing the concentration of products increases the driving force for the reaction, resulting in a higher emf and vice versa Concentration Cells o Cell emf depends on the concentration so a voltaic cell can be constructed using the same species in both the anode and cathode compartments as long as the concentration are different A cell based solely on the emf generated because of a difference in a concentration is called a concentration cell o E.g. Nickel Oxidation of Ni(s) occurs in the half-cell containing the more dilute solution, thereby increasing the concentration of Ni 2+ (aq) n (the number of electron being transferred) is equal to : Batteries and Fuel Cells A battery is a portable, self-contained electrochemical power source that consists of one or more voltaic cells

9 o When cells are connected in series, the battery produces a voltage that is the sum of the emfs of the individual cells Higher emfs can also be achieved by using multiple batteries in series o Some batteries are primary cells, meaning that they cannot be recharged A secondary cell can be recharged from an external power source after its emf has dropped Lead-Acid Battery o A 12-V lead-acid automotive battery consists of six voltaic cells in series, each producing 2 V The electrode reactions that occur during discharge are Because the reactants are solids, there is no need to separate the cell into anode and cathode compartments Solids are excluded from the reaction quotient Q, the relative amounts of Pb(s), PbO 2 (s), and the PbSO 4 (s) have no effect on the emf, helping the battery maintain a relatively constant emf o Lead-acid batter can be recharged During recharging, an external source of energy is used to reverse the direction of the overall cell reaction Alkaline Battery o Alkaline batteries are nonrechargeable (primary battery) o The anode of this battery consists of powdered zinc metal immobilized in a gel in contact with a concentrated solution of KOH The cathode is a mixture of MnO 2 (s) and graphite Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium-Ion Batteries o Nickel-cadmium (nicad) battery During discharge, cadmium metal is oxidized at the anode of the battery while nickel oxyhydroxide is reduced at the cathode Cadmium is a toxic heavy metal Its use increases the weight of batteries and provides an environmental hazard o Nickel-metal-hydride (NiMH) batteries Cathode reaction of NiMH is the same as that for the nickel-cadmium batteries

10 The anode consists of a metal alloy that has the ability to absorb hydrogen ions During the oxidation at the anode, the hydrogen atoms lose electrons, and the resulting H + ions react with OH - ions to form H 2 O Due to the robustness of the batteries toward discharge and recharge, the batteries can last up to 8 years o Lithium-ion (Li-ion) battery Lithium is a very light element and therefore achieve a greater energy density the amount of energy stored per unit mass than nickelbased batteries It is based on the ability of Li + ions to be inserted into and removed from certain layered solids Hydrogen Fuel Cells o The direct production of electricity from fuels by a voltaic cell could, in principle, yield a higher rate of conversion of the chemical energy of the reaction Voltaic cells that perform this conversion using conventional fuels, such as H 2 and CH 4 are called fuel cells Strictly speaking, fuel cells are not batteries o In the fuel cell for the reaction of hydrogen and oxygen, the anode and cathode are separated by a thin polymer Protons are able to pass through these polymers but electrons cannot 20.8: Corrosion Corrosion reactions are spontaneous redox reactions in which a metal is attacked by some substance in its environment and converted to an unwanted compound For nearly all metals, oxidation is a thermodynamically favorable process in air at room temperature o When oxidation process is not inhibited in some way, it can be very destructive to whatever object is made from the metal o Oxidation can form an insulating protective oxide layer that prevents further reaction of the underlying metal Corrosion of Iron o Rusting of iron requires both oxygen and water Other factors such as the ph of the solution, the presence of salts, contact with metal more difficult to oxidize than iron, and stress on the iron can accelerate rusting

o Corrosion of iron is electrochemical in nature Electrons can move through the metal from a region where oxidation occurs to another region where reduction occurs Preventing the Corrosion of Iron o

11 o Corrosion of iron is electrochemical in nature Electrons can move through the metal from a region where oxidation occurs to another region where reduction occurs Preventing the Corrosion of Iron o Iron is often covered with a coat of paint or another metal such as tin or zinc to protect its surface against corrosion If the coating is broken and the iron is exposed to oxygen and water, corrosion will begin o Galvanized iron, which is iron coated with a thin layer of zinc, uses the principles of electrochemistry to protect the iron from corrosion even after the surface coat is broken The Zn(s) is easier to oxidize than Fe(s) Thus, even if the zinc coat is broken, the zinc will serves as the anode and is corroded instead of iron o Protecting a metal from corrosion by making it the cathode in an electrochemical cell is known as cathodic protection The metal that oxidized while protecting the cathode is called the sacrificial anode 20.9: Electrolysis Electrical energy can be used to cause nonspontaneous redox reactions to occur o Such processes, which are driven by an outside source of electrical energy, are called electrolysis reactions and take place in electrolytic cells An electrolytic cell consists of two electrodes in a molten salt or a solution A battery or some other source of direct electrical current acts as an electron pump, pushing electrons into one electrode and pulling them from the other

12 The electrode of the electrolytic cell that is connected to the negative terminal of the voltage source is the cathode of the cell Several practical applications of electrochemistry are based on active electrodes electrodes that participate in the electrolysis process o Electroplating, for example, uses electrolysis to deposit a thin layer of one metal on another metal to improve beauty or resistance to corrosion Quantitative Aspects of Electrolysis o For any half-reaction, the amount of a substance that is reduced or oxidized in an electrolytic cell is directly proportional to the number of electrons passed into the cell o A coulomb is the quantity of charge passing a point in a circuit in 1 s when the current is 1 ampere (A) Coulombs = amperes seconds o Electrons can be thought of as reagents in electrolysis reactions Electrical Work o -w max means that a voltaic cell does work on its surrounding o E ext > E cell is needed to bring about a nonspontaneous electrochemical process o When an external potential E ext is applied to a cell, the surroundings are doing work on the system n is the number of moles of electrons forced into the system by the external potential o n F is the total electrical charge supplied to the system by the external source of electricity o watt (W) is a unit of electrical power Watt-second is a joule Kilowatt-hour (kwh) is equal to J

Oxidation-Reduction Review. Electrochemistry. Oxidation-Reduction Reactions. Oxidation-Reduction Reactions. Sample Problem.

Oxidation-Reduction Review. Electrochemistry. Oxidation-Reduction Reactions. Oxidation-Reduction Reactions. Sample Problem. 1 Electrochemistry Oxidation-Reduction Review Topics Covered Oxidation-reduction reactions Balancing oxidationreduction equations Voltaic cells Cell EMF Spontaneity of redox reactions Batteries Electrolysis