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Chapter 18
Electrochemistry
| Term | Definition |
|---|---|
| Electrical Current | Is the flow of electric charge |
| Electrochemical Cell | A device in which a chemical reaction either produces of is carried out by an electrical current |
| Voltaic(Galvanic) Cell | An electrochemical cell which produces electrical current from a spontaneous chemical reaction |
| Electrolytic Cell | An electrochemical cell which uses electrical current to drive a nonspontaneous chemical reaction |
| Half-Cell | One half of an electrochemical reaction where either oxidation or reduction occurs |
| Electrotrode | A conductive surface through which electrons can enter or leave a half-cell |
| Ampere (A) | The SI unit used for electrical current. One ampere represents the flow of one coulomb(a measure of electrical charge) per second. 1 A = 1C/S... Since an electron has charge 1.602*10^-19 C, 1 A corresponds to the flow of 6.242*10^18 electrons per second |
| Potential Difference | A measure of the difference in potential energy(usually in joules) per unit of charge(coulomb). Is the driving force for electrical current. The SI unit of potential difference is the volt(V), which is equal to one joule per coulomb.. 1 V = 1 J/C |
| Electromotive Force(EMF) | The force that results in the motion of electrons due to a difference in potential |
| Cell Potential(Ecell) or Cell EMF | The potential difference between the cathode and anode in an electrochemical cell. Cell potential depends on relative tendencies of the reactants to undergo oxidation and reduction |
| Anode | The electrode in an electrochemical cell where oxidation occurs; electrons flow away from the anode. The more negatively charged electrode |
| Cathode | The electrode in an electrochemical cell where reduction occurs; electrons flow toward the cathode. The more positively charged electrode |
| Electron Flow in Electrochemical Cell.. (PE=Potential Energy) | Electrons flow from anode to cathode(from negative to positive) through wires connecting electrodes. When cells connected, electrons flow from electrode with more negative charge(greater PE) to electrode with more positive charge(less PE) |
| Salt Bridge | An inverted, U-shaped tube containing a strong electrolyte such as KNO3 that connects the two half cells, allowing the flow of ions that neutralizes the charge buildup |
| 1. Cell Diagram(Line Notation) | The electrochemical cell discussed in which Zn is oxidizes to Zn^2+ and Cu^2+ is reduced to Cu is represented as follows... Zn(S)|Zn^2+(aq)||Cu^2+(aq)|Cu(s)... |
| 2. Cell Diagram(Line Notation) | 1. We write the oxidation half-reaction on the left and the reduction on the right. A double vertical line, indicating the salt bridge, separates the two half-reactions |
| 3. Cell Diagram(Line Notation) | 2. Substances in different phases are separated by a single vertical line, which represents the boundary between the phases.. 3. For some redox reactions, the reactants and products of one or both of the half reactions may be n the same phase.. |
| 4. Cell Diagram(Line Notation) | 3.(cont.) In these cases, the reactants and products are simply separated from each other with a comma in the line diagram. Such cells use an inert electrode, such as platinum(Pt) or graphite, as the anode or cathode(or both) |
| Standard Electrode Potential | The potential of an electrode in a half-cell |
| Standard Hydrogen Electrode(SHE) | The half-cell consisting of an inert platinum electrode immersed in 1 M HCl with H2(g) at 1 atm bubbling through solution; used as standard of a cell potential zero. Arbitrarily assigned a zero. All other electrode potentials measured relative to SHE |
| Defining E°cell | The difference in voltage between the cathode(final state) and the anode(initial state)... E°cell = E°final - E°initial.. E°cell = E°cathode - E°anode |
| Potential Energy | The more negative the electrode potential is, the greater the potential energy of an electron at that electrode(because negative charge repels electrons) |
| 1. Summarizing Standard Electrode Potentials | 1. The electrode potential of the standard hydrogen electrode(SHE) is exactly zero.. 2. The electrode in any half-cell with a greater tendency to undergo reduction is positively charge relative to the SHE and therefore has a positive E° |
| 2. Summarizing Standard Electrode Potentials | 3. The electrode in any half-cell with a lesser tendency to undergo reduction(or greater tendency to undergo oxidation) is negatively charged relative to the she and therefore has a negative E° |
| 3. Summarizing Standard Electrode Potentials | 4. The cell potential of any electrochemical cell(E°cell) is the difference between the electrode potentials of cathode and anode(E°cell = E°cathode - E°anode).. 5. E°cell is positive for spontaneous reactions and negative for nonspontaneous reactions |
| 1. Summarizing the Prediction of Spontaneous Direction for Redox Reactions | 1. The half-reaction with the more positive electrode potential attracts electrons more strongly and will undergo reduction(so substances listed at the top of Table 18.1 tend to undergo reduction, they are good oxidizing agents) |
| 2. Summarizing the Prediction of Spontaneous Direction for Redox Reactions | 2. The half-reaction with the more negative electrode potential repels electrons more strongly and will undergo oxidation(so substances listed near the bottom of Table 18.1 tend to undergo oxidation; they are good reducing agents) |
| 3. Summarizing the Prediction of Spontaneous Direction for Redox Reactions | 3. Any reduction in Table 18.1 is spontaneous when paired with the reverse of the reaction listed below it |
| Predicting Whether a Metal Will Dissolve in Acid | In general, metals whose reduction half-reaction are listed below the reduction of H^+ to H2 in Table 18.1 will dissolve in acids, while metals listed below it will not. |
| 1. Cell Potential, Free Energy, and the Equilibrium Constant Generilizations | 1. For a spontaneous reaction(one that will proceed in the forward direction when all reactants and products are in their standard states)... ΔG° is negative(<0).. E°cell is positive(>0).. K > 1 |
| 2. Cell Potential, Free Energy, and the Equilibrium Constant Generilizations | 2. For a nonspontaneous reaction(one that will proceed in the reverse direction when all reactants and products are in their standard states)... ΔG° is positive(>0).. E°cell is negative(<0).. K < 1 |
| Faraday's Constant(F) | We can quantify the charge(q) that flows in an electrochemical reaction by using Faraday's constant(F), which represents the charge in coulombs of 1 mol of electrons... F = 96485 C / mol e^- |
| Equation Relating Free Energy(ΔG°) and Cell Potential(E°cell) | ΔG° = -nFE°cell... where ΔG° is the standard change in free energy for an electrochemical reaction, n is the number of moles of electrons transferred in the balanced equation, F is Faraday's constant, and E°cell is the standard cell potential |
| Equation for Relating Cell Potential(E°cell) and the Equilibrium Constant(K) | E°cell = [(0.0592 V) / (n)] * log(K)... where E°cell is the standard cell potential, n is the number of moles of electrons transferred in the redox reaction, and K is the equilibrium constant for the balanced redox reaction at 25°C |
| Equation Relating Free Energy(ΔG°) and the Equilibrium Constant(K) | ΔG° = -RTln(K)... where K is equilibrium constant, T is temperature in Kelvins, and R is the gas constant(8.314 J/mol*K) |
| Equation for Relating Ecell(under nonstandard conditions) and E°cell | ΔG = ΔG° + RTLn(Q)... where R is that gas constant(8.314 J/mol*K), T is the temperature in kelvins, and Q is the equilibrium quotient corresponding to the nonstandard conditions |
| 1. Nernst Equation(For nonstandard conditions) | The equation relating the cell potential of an electrochemical cell to the standard cell potential and the reaction quotient |
| 2. Nernst Equation(For nonstandard conditions) | Ecell = E°cell - [(0.0592 V) / (n)] * log(Q)... where Ecell us cell potential in volts, E°cell is the standard cell potential in volts, n is the number of moles of electrons transferred in the redox reaction, and Q is the reaction quoteint |
| Nernst Equation Under Standard Conditions | Under standard conditions, Q=1, and Ecell = E°cell |
| 1. Conclusions From Nernst Equation | 1. When a redox reaction within a voltaic cell occurs under standard condistions, Q=1; therefore Ecell = E°cell |
| 2. Conclusions From Nernst Equation | 2. When a redox reaction within a voltaic cell occurs under conditions in which Q<1, the greater concentration of reactants relative to products drives the reation to the right, resulting in Ecell > E°cell |
| 3. Conclusions From Nernst Equation | 3. When a redox reaction within a voltaic cell occurs under conditions in which Q>1, the greater concentration of products relative to reactants drives the reation to the left, resulting in Ecell < E°cell |
| 4. Conclusions From Nernst Equation | 4. When a redox reaction reaches equilibrium, Q=K. The redox reaction has no tendency to occur in either direction and Ecell = 0 |
| Concentration Cells | A voltaic cell in which both half-reaction are the same, but n which a difference concentration drives the current flow. |
| 1. Electron Flow in Concentration Cells | Electrons spontaneously flow from the half-cell with the lower copper ion concentration to the half-cell with the higher copper ion concentration. |
| 2. Electron Flow in Concentration Cells | The transfer of electrons from the dilute half-cell to the concentrated results in the forming of Cu^2+ ions in the dilute half-cell. The electrons flow to the concentrated cell, where they react with Cu^2+ ions and reduce them to Cu(S) |
| 3. Electron Flow in Concentration Cells | The flow of electrons has the effect of increasing the concecntration of Cu^2+ in the dilute cell and decreasing concentration of Cu^2+ in the concentrated half-cell |
| Dry-Cell Batteries | A battery that does not contain a large amount of liquid water, often using the oxidation of zinc and the reduction of MnO2 to provide the electrical current. Common batteries, such as the type you find in a flashlight |
| Alkaline Batteries | Dry-cell batteries that employ a slightly different half-reaction in a basic medium. Alkaline batteries have a longer working life and a longer shelf life than their nonalkaline counterparts |
| Lead Acid Storage Batteries | Battery in most automobiles. Consist of six electrochemical cells wired in series. Each cell produces 2 V for 12 V total. Each cell contains porous lead anode where oxidation occurs and lead(IV) oxide cathode where reduction occurs, both in sulfuric acid |
| Nickel-Cadmium(NiCad) Battery | A rechargable battery that consists of an anode composed of solid cadmium and a cathode composed of NiO(OH)(s) in a KOH solution |
| Nickel-Metal Hydride(NiMH) Battery | A rechargable battery that uses the same cathode reaction as the NiCad battery by a different anode reaction, the oxidation of hydrogens in a metal alloy |
| Lithium-Ion Battery | A rechargable battery that produces electrical current in the form of motion of lithium ions from the anode to the cathode |
| Fuel Cell | Most common is hydrogen-oxygen fuel cell, which uses a voltaic cell that used the oxidation of hydrogen and the reduction of oxygen, forming water, to provide electrical current |
| Fuel-Cell Breathalyzer | The fuel-cell breathalyzer works by oxidizing ethyl alcohol in the breath to acetic acid. The electrical current that is produced in proportional to the concentration of ethyl alcohol in the breath |
| Electrolysis | The process by which electrical current is used to drive an otherwise nonspontaneous redox reaction. Electrolytic cells use electrolysis |
| Electrolytic Cell Eletrodes | In the electrolytic cell, the anode has become the cathode(oxidation always occurs at the anode) and the cathode has become the anode |
| 1. Summarizing, Characteristics of Electrochemical Cell Types | In all electrochemical cells: 1. Oxidation occurs at the anode.. 2. Reduction occurs at the cathode.. In voltaic cells: 1. The anode is the source of electrons and has a negative charge(anode is -) |
| 2. Summarizing, Characteristics of Electrochemical Cell Types | In voltaic cells: 2. The cathode draws electrons and has a positive charge(cathode +).. In electrolytic cells: 1. Electrons are drawn away from the anode, which must be connected to the positive terminal of the external power source(anode +).. |
| 3. Summarizing, Characteristics of Electrochemical Cell Types | In electrolytic cells: 2. Electrons are forced to the cathode, which must be connected to the negative terminal of the power source(cathode -) |
| Electrolysis of Mixtures of Cations or Anions | 1. The cation that is most easily reduced(the one with the more positive electrode potential) is reduced first.. 2. The anion that is most easily oxidized(the one with the more negative electrode potential) is oxidized first |
| Electrolysis of Aqueous Solutions | The cations of active metals - those that are no t easily reduced, such as Li^+, K^+, Na^+, Mg^2+, Ca^2+, and Al^3+ - canot be reduced from aqueous solution by electrolysis because water is reduced at a lower voltage |
| Corrosion | The gradual, hearly always undesired oxidation of metals that occurs when they are exposed to oxidizing agents in the environment |
Created by:
TimChemistry2
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