DOCSLIB.ORG

Daniell Cell

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
Daniell Cell Experiment 2 : Daniell Cell 210-512-DW: Electrochemistry Objectives • To build a complete galvanic cell (Daniell cell). • To find the EMF of this system under standard conditions. • To verify the Nernst equation. • To find by potentiometry the copper(II) ion concentration (a) in Cu2+ saturated solution and (b) in an unknown Cu2+ solution. Theory The Daniell Cell in 1836 was the starting point of modern electrochemistry because, for the first time, a reliable source of current at a precise voltage was available. It also was the first practical cell that does not generate a gas when operating. This cell was used as the first standard to defined the unit “Volt”. Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) with Eo = 1.10 V However, when the concentration of either Cu2+(aq) or Zn2+(aq) changes in the solution, the EMF of the cell (or potential) also changes according to the Nernst equation. RT [products] [Zn2+ ] E = Eo − lnQ where Q = therefore: E = 1.10 V − 0.0296 log at 25 °C. nF [reactants] [Cu2+] with R = 8.314 J.K−1.mol−1 , T = 298 K, n = 2 (electron exchanged), F = 96485 C.mol−1. The purpose of this lab is to use the Daniell cell to corroborate the Nernst law (potential vs. concentration). It will also be used to find the concentrations (activity) of copper in a Cu 2+(sat.) reference electrode as well as in an unknown sample of Cu2+. Material Chemical Tools and devices CuSO4·5H2O (249.68 g/mol) two 50 mL beakers ZnSO4 (161.45 g/mol) two 100 mL beakers CuSO4(aq) 0.100 M (for reference electrode). seven 100.0 mL volumetric flasks CuSO4(aq) unknown solution. one 10.00 mL pipette with rubber bulb 1 copper wire (10 cm x 1 mm diameter) 50 mL burette 1 zinc metallic plate (2 cm x 10 cm) pH meter used as a voltmeter 1 zinc wire (diameter = 1 mm) one coaxial cable ending with two alligator clips 6 M or concentrated HNO3 Filter paper used as a salt bridge (1.5 M KNO3) 1.5 M KNO3(aq) copper reference electrode (from exp. 1). Distilled water for all the solutions. emery paper (sandpaper) analytical scale (laboratory balance) - 1 - Laboratory setup oxidation (left) reduction (right) Fig. 2-1. Daniell cell assembly and connections. Two 50 mL beakers are used as half-cell. The salt bridge is a rolled filter paper having both extremities in contact with each solutions together. Some 1.5 M KNO3 drops are added to wet the paper at the center to ensure ionic conductivity. A pH meter is used to measure the EMF of the cell. The device set to “mV” position. It is essential to use a pH meter to read the voltage since this instrument does not require any current from the circuit to perform its measurement (high impedance). A coaxial cable ending with two alligator clips is used (ask teacher). The black cable (low potential) is connected to the electrode performing an oxidation here the zinc. The connection are made according to cell representation (electrons from the left to th right). (left = black cable) Zn | Zn2+(0.10 M) | Cu2+(0.10 M) | Cu (red cable = right) Since the oxidation reaction always happens at the left electrode, a spontaneous reaction (galvanic cell) will give a positive EMF. - 2 - Procedure 1. Make 100.00 mL of 0.10 M ZnSO4(aq). (Use a concentration as close as possible to this value and record the actual one). 2. Make 100.00 mL of 0.30 M CuSO4(aq). (Use a concentration as close as possible to this value and record the actual one). The water used should be acidic, therefore, add 5 drops of conc. HNO 3 to the distilled water to prevent any Cu(OH)2(s) formation. 3. From the 0.30 M CuSO4(aq) solution, prepare a 100.00 mL of both the following diluted copper(II) sulfate solutions: 3.0x10−2 M, 3.00x10−3 M. Record the precise concentrations. hint: use serial dilution to perform those solutions. 4. With a burette, transfer 33.3 mL of the 0.30 M CuSO4(aq) solution into a 100.0 mL volumetric flask to make a 0.10 M CuSO4(aq) solution. 5. From the 0.10 M CuSO4(aq) solution, prepare 100.00 mL of both the following diluted copper(II) sulfate solutions: 1.0x10−2 M, 1.00x10−3 M. Record these precise concentrations. 6. Put together the setup described in Fig. 2-1. 7. Zinc electrode is cleaned with some diluted HNO3 acid (≈ 0.1 M). The copper wire is refreshed by scrubbing the surface with an emery paper. Both electrodes are rinse with distilled water prior to use and placed in their respective beakers. 8. Pour about 30 mL of 0.10 M ZnSO4(aq) in the beaker with the zinc electrode; the volume of liquid does not have be precise. 9. In the same beaker, place your copper reference electrode (Lab 1) filled with Cu2+(aq, 0.10 M). 10. Connect both electrode with the alligator clips in such a way that you get a positive voltage reading (black connector = zinc electrode). Record the value. 11. Repeat this process by replacing your reference electrode by the one provided by the teacher. Ref: Cu with Cu2+(sat.). After the reading, this electrode could be rinse and stored. −3 12. Pour roughly 30 mL of 1.0x10 M CuSO4(aq) in the other beaker (copper wire electrode). 13. Place the salt bridge between your two beakers (strip of filter paper) to ensure the ionic conductivity between the two half cells. The two extremities of the paper are in contact with both solutions. Use some 1.5 M KNO3 to wet the paper at the center. Record the EMF of this cell (in volt). −3 14. While keeping the same zinc solution throughout the experiment, discard the 1.0x10 M CuSO4(aq) solution. Rinse and fill the beaker and the electrode with the next most diluted Cu2+ solution (3.0x10−3 M), and record its potential. 15. Repeat the step number 14 with the four Cu2+ solutions remaining (see data sheet 15 a to 15 d.) in the order of the least concentrated to the most concentrated. 16. Finally, repeat the step number 14 with the unknown solution of Cu2+(aq) provided by the teacher. 17. Put a thermometer in the zinc solution and record its temperature. 18. Put a zinc wire in the copper solution for 10 seconds and note your observations. 19. Put a copper wire in the zinc solution for 10 seconds and note your observations. - 3 - Experiment 2 : Daniel cell data sheet Potential Note: The electrode on the left in the cell notation is the anode. It is the negative one (black cable) in a spontaneous galvanic cell. It is important to follow the cell notation when measuring voltages. step System (cell notation) Potential / mV 10 Zn | Zn2+(0.10 M) | Cu2+(0.1 M) reference electrode 11 Zn | Zn2+(0.10 M) | Cu2+(sat.) reference electrode 13 Zn | Zn2+(0.10 M) | Cu2+(1.0x10−3 M) | Cu 14 Zn | Zn2+(0.10 M) | Cu2+(3.0x10−3 M) | Cu 15a Zn | Zn2+(0.10 M) | Cu2+(0.010 M) | Cu 15b Zn | Zn2+(0.10 M) | Cu2+(0.030 M) | Cu 15c Zn | Zn2+(0.10 M) | Cu2+(0.10 M) | Cu 15d Zn | Zn2+(0.10 M) | Cu2+(0.30 M) | Cu 16 Zn | Zn2+(0.10 M) | Cu2+(unknown solution) | Cu Other readings and observations 17 Temperature / °C : _______________ 18 zinc wire in the copper solution: 19 copper wire in the zinc solution: - 4 - Experiment 2 : Daniell cell Lab report Calculations 1. Plot a graph of: potential vs. ln([Zn2+]/[Cu2+]), with “V” on the y-scale, for the six values of Cu2+(aq). Label your axes with units (if any) and provide a meaningful title. 2. Use the linear part of the graph and calculate the slope. Provide the slope equation with correlation. 3. Use this slope to calculate the number of electrons exchanged in this process (slope = −RT/nF with R in joule.mol−1.K−1). Comment on any discrepancy. 4. Use this slope as a calibration curve to find the concentration of the Cu2+ unknown. 5. Use the Nernst law to calculate the “effective” Cu2+(sat) concentration in the reference electrode. Questions 1. In the Daniell cell, is the ion Cu2+ a product or a reactant? 2. Is the Nernst law respected over the complete range of concentration? In the case of a negative answer, provide a reason? 3. Indicate how the le Châtelier's principle applies to your electrochemical cell in this experiment. 4. The EMF value of the potential recorded in 15c should be the one of the standard potential since [Cu2+(aq)] = [Zn2+(aq)]. Is it the case? If not, How can you explain this difference? 5. Instead of using copper(II) sulfate for this experiment, what will be the effect of using instead: a) copper(II) nitrate b) copper(I) nitrate 6. Would it be possible to make this experiment without the use of a salt bridge? (using a solution containing both Cu2+ and Zn2+ together.) Comment. - 5 -.
Load more

Recommended publications
  • Cell Notation Practical Galvanic Cells -Batteries
    Basic Redox Vocabulary • Write reactions for each of the following: • oxidation of metallic nickel by BiO+ • reduction of Zn2+ by hydroxide ion • reaction of Fe 2+ with Hg2+ 2+ • reaction of Cd with NO2 • AgI acting as an oxidizing agent toward Sn 2+ • What’s wrong with • The oxidation of Cr by Cl- • The reduction of Co 2+ by Ag+ Cell Notation • As was noted earlier, galvanic cells normally consist of two distinct regions, one housing the oxidation half and the other the reduction half. There is a simplified notation form that allows one to represent the cell easily( text p 798- 799). • The oxidation is written on the left and the reduction on the right. starting with the anode material and ending with the cathode material. • phase boundaries represented with single vertical lines “ |” • the physical separation between the two half cells is a double v ertical line “||” if it’s a salt bridge and with a single broken vertical line, “!”, if it’s a liquid junction • within each have cell, the species are written in a reactant-product order, separated by commas if they are in the same phase. Acid/base components should be included • The electrode material may be actively participating in the redox chemistry (active electrode) or merely providing surface for the electron transfer (passive or inert electrode, usually graphite or Pt) • Represent the following as galvanic cells(assume the reactions are spontaneous as written) • Tl(s) + Cd 2+ ó Tl + + Cd(s) - 2+ 2+ • Pb(s) + MnO4 ó Pb + Mn (acid) 2+ 4+ • O2(g) + Sn ó H2O + Sn Practical Galvanic Cells -batteries • Batteries represent the most common application of the electrochemical cell.
    [Show full text]
  • Nernst Equation in Electrochemistry the Nernst Equation Gives the Reduction Potential of a Half‐Cell in Equilibrium
    Nernst equation In electrochemistry the Nernst equation gives the reduction potential of a half‐cell in equilibrium. In further cases, it can also be used to determine the emf (electromotive force) for a full electrochemical cell. (half‐cell reduction potential) (total cell potential) where Ered is the half‐cell reduction potential at a certain T o E red is the standard half‐cell reduction potential Ecell is the cell potential (electromotive force) o E cell is the standard cell potential at a certain T R is the universal gas constant: R = 8.314472(15) JK−1mol−1 T is the absolute temperature in Kelvin a is the chemical activity for the relevant species, where aRed is the reductant and aOx is the oxidant F is the Faraday constant; F = 9.64853399(24)×104 Cmol−1 z is the number of electrons transferred in the cell reaction or half‐reaction Q is the reaction quotient (e.g. molar concentrations, partial pressures …) As the system is considered not to have reached equilibrium, the reaction quotient Q is used instead of the equilibrium constant k. The electrochemical series is used to determine the electrochemical potential or the electrode potential of an electrochemical cell. These electrode potentials are measured relatively to the standard hydrogen electrode. A reduced member of a couple has a thermodynamic tendency to reduce the oxidized member of any couple that lies above it in the series. The standard hydrogen electrode is a redox electrode which forms the basis of the thermodynamic scale of these oxidation‐ reduction potentials. For a comparison with all other electrode reactions, standard electrode potential E0 of hydrogen is defined to be zero at all temperatures.
    [Show full text]
  • Voltaic Cells
    Voltaic Cells Tro Chapter 19 – Electrochemistry 19.3 Voltaic(or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions Electric Current Flowing Directly Between Atoms Tro, Chemistry: A Molecular Approach 2 Electrochemical Cells Voltaic (Galvanic) ΔG < 0 to Electrolytic Δ G > 0 uses electrical generate electrical energy. energy to drive non-spontaneous process. Electrochemical Cells • Oxidation and reduction half-reactions are kept separate in half-cells. • Electron flow through a wire along with ion flow through a solution constitutes an electric circuit. • It requires a conductive solid electrode to allow the transfer of electrons. – Through external circuit – Metal or graphite • Requires ion exchange between the Daniell Cell two half-cells of the system. – Electrolyte Definitions Anode Salt Bridge • Electrode where oxidation always occurs • An inverted, U-shaped tube containing a • More negatively charged electrode in strong electrolyte and connecting the two voltaic cell half-cells. • Typically made of metal that is oxidized Cathode • Electrode where reduction always occurs Potential Difference • More positively charged electrode in • The difference in potential energy voltaic cell between the reactants and products. • Typically metal that is produced by reduction (Caused by an electric field resulting from the charge difference on the two If the redox reaction involves the oxidation or electrodes.) reduction of an ion to a different oxidation state, or the oxidation or reduction of a gas, we Cell Potential (Ecell or emf) may use an inert electrode. • The potential difference between the anode and the cathode in a voltaic cell. • An inert electrode is one that not does participate in the reaction but just provides a surface on which the transfer of electrons can take place.
    [Show full text]
  • Galvanic Cell Notation • Half-Cell Notation • Types of Electrodes • Cell
    Galvanic Cell Notation ¾Inactive (inert) electrodes – not involved in the electrode half-reaction (inert solid conductors; • Half-cell notation serve as a contact between the – Different phases are separated by vertical lines solution and the external el. circuit) 3+ 2+ – Species in the same phase are separated by Example: Pt electrode in Fe /Fe soln. commas Fe3+ + e- → Fe2+ (as reduction) • Types of electrodes Notation: Fe3+, Fe2+Pt(s) ¾Active electrodes – involved in the electrode ¾Electrodes involving metals and their half-reaction (most metal electrodes) slightly soluble salts Example: Zn2+/Zn metal electrode Example: Ag/AgCl electrode Zn(s) → Zn2+ + 2e- (as oxidation) AgCl(s) + e- → Ag(s) + Cl- (as reduction) Notation: Zn(s)Zn2+ Notation: Cl-AgCl(s)Ag(s) ¾Electrodes involving gases – a gas is bubbled Example: A combination of the Zn(s)Zn2+ and over an inert electrode Fe3+, Fe2+Pt(s) half-cells leads to: Example: H2 gas over Pt electrode + - H2(g) → 2H + 2e (as oxidation) + Notation: Pt(s)H2(g)H • Cell notation – The anode half-cell is written on the left of the cathode half-cell Zn(s) → Zn2+ + 2e- (anode, oxidation) + – The electrodes appear on the far left (anode) and Fe3+ + e- → Fe2+ (×2) (cathode, reduction) far right (cathode) of the notation Zn(s) + 2Fe3+ → Zn2+ + 2Fe2+ – Salt bridges are represented by double vertical lines ⇒ Zn(s)Zn2+ || Fe3+, Fe2+Pt(s) 1 + Example: A combination of the Pt(s)H2(g)H Example: Write the cell reaction and the cell and Cl-AgCl(s)Ag(s) half-cells leads to: notation for a cell consisting of a graphite cathode - 2+ Note: The immersed in an acidic solution of MnO4 and Mn 4+ reactants in the and a graphite anode immersed in a solution of Sn 2+ overall reaction are and Sn .
    [Show full text]
  • Lithium Batteries
    Batteries General planning of the „Lithium Batteries” lab for the European Master 2007/8 Warsaw University of Technology, Departament of Inorganic Chemistry and Solid State Technology, Dr Marek Marcinek Objectives: Students will: follow the development of primary and secondary lithium batteries become familiar with different types of batteries explore the applications of batteries study the major components of lithium (ion) cells learn which batteries can be recycled realize the economic and environmental advantages of using rechargeable batteries Needs: Room with avialiables 2 desks and 2 computers or Multimedia Projector Potentiostat/galvanostat with galvanic cycle mode. Suplementary materials needed for: Building the simple battery (included Volta cell) Daniell cell Leclanche Cell Commercially available variety of Li-bat (also disassembled in a controlled mode) BDS software Safety issues: Safety rules Read directions carefully before you begin any experiments. Clear an area to work. Wash your hands thoroughly after experimenting Do always wear eye protection Keep all chemicals away from your eyes and mouth Do not eat and drink in your experiment area Put all pieces of equipment away when finished using them General First aid information Indicate the person who should IMMIDIATELLY inform the teacher and ask the other person to help you out if needed. Eyes: rinse immediately with water. Remove contact lenses if wearing any. Flush eyes with water for 15 min Swallowed: Rinse mouth Drink glass full of water or milk. DO NOT INDUCE VOMITING SKIN: Flush skin thoroughly with water. In all cases, get immediate medical attention if an emergency exists. Bring the chemical container with you. 1 Materials given to students/preparation: Exemplary 5 scientific papers (or any material found related to the Battery Performance, Design, Safety, Application) for individual preparation.
    [Show full text]
  • Electrochemistry –An Oxidizing Agent Is a Species That Oxidizes Another Species; It Is Itself Reduced
    Oxidation-Reduction Reactions Chapter 17 • Describing Oxidation-Reduction Reactions Electrochemistry –An oxidizing agent is a species that oxidizes another species; it is itself reduced. –A reducing agent is a species that reduces another species; it is itself oxidized. Loss of 2 e-1 oxidation reducing agent +2 +2 Fe( s) + Cu (aq) → Fe (aq) + Cu( s) oxidizing agent Gain of 2 e-1 reduction Skeleton Oxidation-Reduction Equations Electrochemistry ! Identify what species is being oxidized (this will be the “reducing agent”) ! Identify what species is being •The study of the interchange of reduced (this will be the “oxidizing agent”) chemical and electrical energy. ! What species result from the oxidation and reduction? ! Does the reaction occur in acidic or basic solution? 2+ - 3+ 2+ Fe (aq) + MnO4 (aq) 6 Fe (aq) + Mn (aq) Steps in Balancing Oxidation-Reduction Review of Terms Equations in Acidic solutions 1. Assign oxidation numbers to • oxidation-reduction (redox) each atom so that you know reaction: involves a transfer of what is oxidized and what is electrons from the reducing agent to reduced 2. Split the skeleton equation into the oxidizing agent. two half-reactions-one for the oxidation reaction (element • oxidation: loss of electrons increases in oxidation number) and one for the reduction (element decreases in oxidation • reduction: gain of electrons number) 2+ 3+ - 2+ Fe (aq) º Fe (aq) MnO4 (aq) º Mn (aq) 1 3. Complete and balance each half reaction Galvanic Cell a. Balance all atoms except O and H 2+ 3+ - 2+ (Voltaic Cell) Fe (aq) º Fe (aq) MnO4 (aq) º Mn (aq) b.
    [Show full text]
  • Units of Free Energy and Electrochemical Cell Potentials Notes on General Chemistry
    Units of free energy and electrochemical cell potentials Notes on General Chemistry http://quantum.bu.edu/notes/GeneralChemistry/FreeEnergyUnits.pdf Last updated Wednesday, March 15, 2006 9:45:57-05:00 Copyright © 2006 Dan Dill ([email protected]) Department of Chemistry, Boston University, Boston MA 02215 "The devil is in the details." Units of DG for chemical transformations There is a subtlety about the relation between DG and K, namely how DG changes when the amounts of reactants and products change. Gibbs free energy G = H - TS is an extensive quantity, that is, it depends on how much of the system we have, since H and S are both extensive (T is intensive). This means if we double the reactants and products, 2 a A V 2 b B then DG must also double, but the expression DG = RTln Q K found in textbooks does not seem to contain the amounts of reactants and products! In fact it does contain these, because theH êstoichiometricL coefficients appear as exponents in Q and K. Thereby, doubling the moles of reactants and products will change Q and K into their square, which will double DG. Here is how. DG doubled = RTln Q2 K2 = RTln Q K 2 H = 2LRTln Q K H ê L = 2 DG original H ê L This result evidently means thatH ê L H L the units of DG are energy, not energy per mole, but with the understanding that it is the change in Gibbs free energy per mole of reaction as written. The question arises, why is this not taken account by the explicit units in the equation for DG? The answer is that the appropriate unit has been omitted.
    [Show full text]
  • The Metallic World
    UNIT 11 The Metallic World Unit Overview This unit provides an overview of both electrochemistry and basic transition metal chemistry. Oxidation-reduction reactions, also known as redox reactions, drive electrochemistry. In redox reactions, one compound gains electrons (reduction) while another one loses electrons (oxidation). The spontaneous directions of redox reactions can generate electrical current. The relative reactivities of substances toward oxidation or reduction can promote or prevent processes from happening. In addition, redox reactions can be forced to run in their non-spontaneous direction in order to purify a sample, re-set a system, such as in a rechargeable battery, or deposit a coating of another substance on a surface. The unit also explores transition metal chemistry, both in comparison to principles of main-group chemistry (e.g., the octet rule) and through various examples of inorganic and bioinorganic compounds. Learning Objectives and Applicable Standards Participants will be able to: 1. Describe the difference between spontaneous and nonspontaneous redox processes in terms of both cell EMF (electromotive force) and DG. 2. Recognize everyday applications of spontaneous redox processes as well as applications that depend on the forcing of a process to run in the non-spontaneous direction. 3. Understand the basics of physiological redox processes, and recognize some of the en- zymes that facilitate electron transfer. 4. Compare basic transition-metal chemistry to main-group chemistry in terms of how ions are formed and the different types of bonding in metal complexes. Key Concepts and People 1. Redox Reactions: Redox reactions can be analyzed systematically for how many electrons are transferred, whether or not the reaction happens spontaneously, and how much energy can be transferred or is required in the process.
    [Show full text]
  • Dissociation Constants of Oxalic Acid in Aqueous Sodium Chloride and Sodium Trifluoromethanesulfonate Media to 175 °C
    View metadata, citation and similar papers at core.ac.uk brought to you by CORE provided by DigitalCommons@University of Nebraska University of Nebraska - Lincoln DigitalCommons@University of Nebraska - Lincoln Earth and Atmospheric Sciences, Department Papers in the Earth and Atmospheric Sciences of 1998 Dissociation Constants of Oxalic Acid in Aqueous Sodium Chloride and Sodium Trifluoromethanesulfonate Media to 175 °C Richard Kettler University of Nebraska-Lincoln, [email protected] David J. Wesolowski Oak Ridge National Laboratory Donald A. Palmer Oak Ridge National Laboratory Follow this and additional works at: https://digitalcommons.unl.edu/geosciencefacpub Part of the Earth Sciences Commons Kettler, Richard; Wesolowski, David J.; and Palmer, Donald A., "Dissociation Constants of Oxalic Acid in Aqueous Sodium Chloride and Sodium Trifluoromethanesulfonate Media to 175 °C" (1998). Papers in the Earth and Atmospheric Sciences. 135. https://digitalcommons.unl.edu/geosciencefacpub/135 This Article is brought to you for free and open access by the Earth and Atmospheric Sciences, Department of at DigitalCommons@University of Nebraska - Lincoln. It has been accepted for inclusion in Papers in the Earth and Atmospheric Sciences by an authorized administrator of DigitalCommons@University of Nebraska - Lincoln. J. Chem. Eng. Data 1998, 43, 337-350 337 Dissociation Constants of Oxalic Acid in Aqueous Sodium Chloride and Sodium Trifluoromethanesulfonate Media to 175 °C Richard M. Kettler*,† Department of Geology, University of Nebraska, Lincoln, Nebraska 68588-0340 David J. Wesolowski‡ and Donald A. Palmer§ Chemical and Analytical Sciences Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 37831-6110 The first and second molal dissociation constants of oxalic acid were measured potentiometrically in a concentration cell fitted with hydrogen electrodes.
    [Show full text]
  • Battery Technologies for Small Scale Embeded Generation
    Battery Technologies for Small Scale Embedded Generation. by Norman Jackson, South African Energy Storage Association (SAESA) Content Provider – Wikipedia et al Small Scale Embedded Generation - SSEG • SSEG is very much a local South African term for Distributed Generation under 10 Mega Watt. Internationally they refer to: Distributed generation, also distributed energy, on-site generation (OSG) or district/decentralized energy It is electrical generation and storage performed by a variety of small, grid- connected devices referred to as distributed energy resources (DER) Types of Energy storage: • Fossil fuel storage • Thermal • Electrochemical • Mechanical • Brick storage heater • Compressed air energy storage • Cryogenic energy storage (Battery Energy • Fireless locomotive • Liquid nitrogen engine Storage System, • Flywheel energy storage • Eutectic system BESS) • Gravitational potential energy • Ice storage air conditioning • Hydraulic accumulator • Molten salt storage • Flow battery • Pumped-storage • Phase-change material • Rechargeable hydroelectricity • Seasonal thermal energy battery • Electrical, electromagnetic storage • Capacitor • Solar pond • UltraBattery • Supercapacitor • Steam accumulator • Superconducting magnetic • Thermal energy energy storage (SMES, also storage (general) superconducting storage coil) • Chemical • Biological • Biofuels • Glycogen • Hydrated salts • Starch • Hydrogen storage • Hydrogen peroxide • Power to gas • Vanadium pentoxide History of the battery This was a stack of copper and zinc Italian plates,
    [Show full text]
  • 3297 Chapter 13.Indd
    Inert Electrodes The zinc anode and copper cathode of a Daniell cell and the silver and chromium electrodes in Figure 13.4 are all metals and can act as electrical conductors. However, some redox reactions involve oxidizing and reducing agents that are not solid metals but, instead, are dissolved electrolytes or gases and, therefore, cannot be used as electrodes. To construct a voltaic cell that will use these oxidizing and reducing agents, you have to use inert electrodes. An inert electrode is an electrode made from a material that is neither a reactant nor a product of the redox reaction. Instead, the inert electrode can carry a current and provide a surface on which redox reactions can occur. Figure 13.5 shows a cell that contains one example of an inert electrode—a platinum electrode. The complete balanced equation, net ionic equation, and half-reactions for this cell are given below. complete balanced equation: Pb(s) + 2FeCl3(aq) → 2FeCl2(aq) + PbCl2(aq) net ionic equation: Pb(s) + 2Fe3+(aq) → 2Fe2+(aq) + Pb2+(aq) oxidation half-reaction: Pb(s) → Pb2+(aq) + 2e– reduction half-reaction: Fe3+(aq) + e– → Fe2+(aq) The anode is the lead electrode. Lead atoms lose electrons that remain in the electrode while the lead(II) ions dissolve in the solution in the same way that the anode did in previous example. However, the reduction half-reaction involves iron(III) ions that accept an electron from the platinum inert electrode and become iron(II) ions. The platinum atoms in the electrode (cathode) remain unchanged. Voltmeter e- e- Anode Salt bridge Cathode (-) K+ (+) Pb Cl- Pt 2e- e- Fe3+ Pb Fe2+ Pb2+ PbCl2 FeCl3 FeCl2 Figure 13.5 This cell uses an inert electrode to conduct electrons.
    [Show full text]
  • Ch.14-16 Electrochemistry Redox Reaction
    Redox Reaction - the basics ox + red <=> red + ox Ch.14-16 1 2 1 2 Oxidizing Reducing Electrochemistry Agent Agent Redox reactions: involve transfer of electrons from one species to another. Oxidizing agent (oxidant): takes electrons Reducing agent (reductant): gives electrons Redox Reaction - the basics Balance Redox Reactions (Half Reactions) Reduced Oxidized 1. Write down the (two half) reactions. ox1 + red2 <=> red1 + ox2 2. Balance the (half) reactions (Mass and Charge): a. Start with elements other than H and O. Oxidizing Reducing Agent Agent b. Balance O by adding water. c. balance H by adding H+. Redox reactions: involve transfer of electrons from one d. Balancing charge by adding electrons. species to another. (3. Multiply each half reaction to make the number of Oxidizing agent (oxidant): takes electrons electrons equal. Reducing agent (reductant): gives electrons 4. Add the reactions and simplify.) Fe3+ + V2+ → Fe2+ + V3 + Example: Balance the two half reactions and redox Important Redox Titrants and the Reactions reaction equation of the titration of an acidic solution of Na2C2O4 (sodium oxalate, colorless) with KMnO4 (deep purple). Oxidizing Reagents (Oxidants) - 2- 2+ MnO4 (qa ) + C2O4 (qa ) → Mn (qa ) + CO2(g) (1)Potassium Permanganate +qa -qa 2-qa 16H ( ) + 2MnO4 ( ) + 5C2O4 ( ) → − + − 2+ 2+ MnO 4 +8H +5 e → Mn + 4 H2 O 2Mn (qa ) + 8H2O(l) + 10CO2( g) MnO − +4H+ + 3 e − → MnO( s )+ 2 H O Example: Balance 4 2 2 Sn2+ + Fe3+ <=> Sn4+ + Fe2+ − − 2− MnO4 + e→ MnO4 2+ - 3+ 2+ Fe + MnO4 <=> Fe + Mn 1 Important Redox Titrants
    [Show full text]