reaction Electrochemistry
Overall Half reactions
Zn(s) + Cu (aq) Zn2+(aq) + Cu(s)
2+
Oxidation loss of electrons
Zn(s) Zn (aq) + 2e
2+
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Daniel Cell
If direct contact between the copper ions and zinc metal is avoided, this reaction can be made to do useful work. In voltaic cells direct contact is avoided by physically separating the oxidation and reduction half-reactions. The electrons are forced to flow through an external wire from the zinc half-cell to the copper halfcell. This flow of electrons through an external wire can be used to produce work. A practical voltaic cell is shown below. At the anode, zinc metal is oxidized to zinc ions. Electrons then travel through the external wire to the cathode where copper ions are reduced to copper metal. As the reaction proceeds, Zn2+ ions are produced at the anode and Cu2+ ions are consumed at the cathode. To maintain electrical neutrality, sulfate ions must flow through the salt bridge from the half-cell on the left to the half-cell on the right to balance the flow of electrons through the external wire from the anode to cathode.
Chapter18 Electrochemistry
2
2+
loss of electrons, Oxidation, Anode, -ve
LEO goes
Zn(s) Zn (aq) + 2e (electrons produced)
(oxidation and anode start with vowels)
2+
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Cell Notation for Galvanic or voltaic cell
Oxidation Reduction Cell notation
2+
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The driving force that pushes the negatively charged electrons away form the anode and pulls them toward the cathode is an electrical potential called the electromotive force (emf) or the cell potential (E) or the cell voltage. Cell potential =W = Q work done (joules) charge transferred (coulombs)
( G) Free energy (Joule, J) measure of the spontaneity of a reaction carried out at a constant temperature. If G is negative the reaction is spontaneous If G is positive the reaction is nonspontaneous If G is zero the reaction mixture is at equilibrium
G -nE +ve measurement means a reaction will be spontaneous
n = moles of electrons transferred in the reaction E = cell potential (J)
G = -nFE
F = Faraday constant the electrical charge on one mole of electrons
F = 96,500 C/mole eWhere C is the electric charge (coulomb) When 1 C of charge moves between two electrodes that differ in electrical potential by 1 volt, 1 joule of energy is released by the cell and can be used to do electrical work. Volt is potential energy per unit charge V = 1J/1C The cell potential, measured by connecting a voltmeter between the two electrodes, is a measure of the driving force behind this reaction Chapter18 Electrochemistry 5
Standard Reduction Potentials
Just as a chemical reaction can conceptually be broken into two half-reactions the cell potential can be thought to be composed of two half-cell potentials. The reaction between zinc metal and acid, for example, results from the combination of two half-reactions:
2+ -
Zn(s) Zn (aq) + 2e
(oxidation)
By arbitrarily defining the potential for the half-reaction;
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6
The standard state cell potential, Eo, is a state function. If we reverse the direction in which one of these half-cells is run, all that we have to do is reverse the sign of E :
Half cell potentials. o
Stronger oxidizing agent Half-reaction Ag+(aq) + e- Ag(s) 0.80 Fe3+(aq) + 1e- Fe2+(aq) Cu2+(aq) + 2e- Cu(s) 2H+(aq) + 2e- H2(g) Pb2+(aq) + 2e- Pb(s) Ni2+(aq) + 2e- Ni(s) Cd2+(aq) + 2e- Cd(s) Zn2+(aq) + 2e- Zn(s) 0.77 0.34 0.00 -0.13 -0.25 -0.40 7 -0.76 E (V) All potentials are tabulated as reduction potentials. Standard reduction potentials refer to measurements taken in the standard state, i.e. 1 M concentrations for all ions, 25oC, and 1 atm of pressure for all gases
Fe2+(aq) + 2e- CFe(s) -0.44 hapter18 Electrochemistry
What is the standard cell potential produced in a cell using aluminium and nickel electrodes? Write the cell notation.
2+ o
When a reaction is tripled or doubled the E does not change because, for example when we double the reaction twice as much energy and twice as much charge are transferred but the ratio of energy to charge stays the same.
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Ni (aq) + 2e Ni(s)
E = -0.25
Concentration Effects on Cell Potentials The Nernst Equation
The cell potential (E) for systems which are not at standard state can be calculated from the Nernst equation: where:
0 E = E cell 0 E = E cell 0 E = E cell
RT ln Q nF 2.3026 RT log Q nF 0.0592 log Q n
Eo = the standard state cell potential R = ideal gas constant (8.314 J mol-1 K-1) T = temperature in Kelvin (298.15 K) n = number of moles of electrons transferred per mole of reactant F = the Faraday constant (96 487 Coulombs) Q = the thermodynamic reaction quotient for the cell reaction
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For the zinc-copper cell:
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) n is 2 and the reaction quotient, Q, is:
[ Zn 2 +] Q= [Cu 2 +]
If [Zn2+] = [Cu2+], Q = 1, log10Q = 0 and Ecell = Eo. When the ratio of the concentration is not 1, the cell potential differs from the standard state cell potential. For example, if [Zn2+] = 0.100 M and [Cu2+] = 1.00 M, the cell potential would be:
E = Eo -
E = 1.103 - [(0.030)(-1)] = 1.133 V When a species is not converted to a gas or metal but to another ionic species care must be used in determining the expected cell potential even if the concentrations are the same.
0.0592 0.100 log10 2 1.00
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For example, a galvanic cell is made in which a copper electrode in a 0.1 M solution of Cu(NO3)2 is connected to a graphite electrode in a 0.1M solution of SnCl2 and SnCl4, what is the cell potential?
2+ 4+
Sn (aq) 2e + Sn
E = -0.13
Concentration Cells
An interesting voltaic cell can be devised in which both electrodes involve the same half-reaction, but at differing concentrations. Assume, for example, a voltaic cell in which pieces of copper metal are immersed in 1.0 M and 0.10 M solutions of Cu2+, respectively, and then connected to a voltmeter. The standard state cell potential (Eo) for this system is zero, since the half-cell reaction is the same at the anode and at the cathode. However, there will be a small, but measurable cell potential (E), since the reaction quotient is not equal to 1:
2+ o
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Cuo 2e- + Cu2+(aq) initial 0.1 M 2+ Cu (aq) + 2e- Cuo initial 1.0 M Cu(s) + Cu2+ Cu2+ + Cu(s)
Eo = -.34 Eo = 0.34 Eo = 0.00 E = 0.00 0.0592 [ 0.1] log10 [1.0] 2 E = 0.030 V
Ox Red
V
Cathode (+ve) Anode (-ve)
o 2+ o
2+
-
Cu Cu + 2e (0.1M) Copper is plating off of the electrode and the concentration is increasing in the porous cup NO Nitrate flows into porous
12
Chapter18 Electrochemistry
Equilibrium
The copper concentration cell will not run on indefinitely. We know that a closed system proceeds spontaneously towards a stat of dynamic equilibrium. When equilibrium is reached, there is no further conversion of reactants into products. When an electron transfer reaction stops there is no current flowing and the voltage becomes zero At equilibrium Ecell = 0 V Q = K (the equilibrium constant for the cell reaction.
0 cell
0V = E
0.0592 log K n
0 nE cell log k = 0.0592 V
Example: Use standard reduction potentials to calculate the equilibrium constant for the reaction below. 2 Cr3+(aq) + 3 Ni(s) 2 Cr(s) + 3 Ni2+(aq) Ni2+(aq) + 2e-0.257 Ni(s) E =
Cr3+(aq) + 3e- Cr(s) E = -0.74
0 nE cell log k = 0.0592 V
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Example: Use standard reduction potentials to calculate the equilibrium constant for the reaction below. 6 Br(aq) + Cr2O72-(aq) + 14 H+ (aq) 3 Br2(l) + 2 Cr3+(aq) + 7 H2O
2
Oxidation 3 O72-(aq) + B aq) 2e ] Reduction Cr2x [2 Br(aq)14 H+r((l) ++ 6e + 2 Cr3+(aq) + 7 H2O (l) 6 Br(aq) + Cr2O72-(aq) + 14 H+ (aq) 3 Br2(l) + 2 Cr3+(aq) + 7 H2O
E = - 1.09 1.33 E = 0.24
0 nE cell log k = 0.0592 V
Unlike acid-base equilibrium redox reactions typically go either go essentially to completion (K is very large) or almost not at all (K is very small) Summary E > 0 K > 1 Starting from standard conditions (Q=1) the reaction must go to from left to right to establish equilibrium (Q=K) Thus E > 0 under standard conditions the reaction proceeds spontaneously from left to right E < 0 K < 1 Starting from standard conditions (Q=1) the reaction must go to from right to left to establish equilibrium (Q=K) Thus E < 0 under standard conditions the reaction proceeds spontaneously from right to left Chapter18 Electrochemistry 14
Ionization energies and Oxidation Potentials
Ionization energies generally increase from left to right across a transition series, though there are some irregularities for the atoms of the first transition series. The general trend correlates with an increase in effective nuclear charge and a decrease in atomic radius.
Chapter18 Electrochemistry 15
Oxidation Potentials
Atomic radii
The second ionization energy for Cu is the highest within this group Note: potentials are the negative of the corresponding standard reduction potentials. The larger the value the more readily the metal will give up an electron. Except for copper all, the Eo values are positive, which means that the solid metal is oxidized to its aqueous cation more readily than H2 gas is oxidized to H+(aq) In other words, except for copper the first-series metals are stronger reducing agents than H2 gas and can therefore be oxidized by the H+ ion in acids like HCl that lack an oxidizing anion: M(s) + 2 H+(aq) M2+(aq) + H2(g) Eo > V (except for M = Cu)
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The standard potential for the oxidation of a metal is a composite property that depends on Go for the sublimation of the metal, the ionization energies of the metal atom and Go for the hydration of the metal ion.
o
Nevertheless the general trend in the E values correlates with the general trend in the ionization energies. The ease of oxidation of the metal decreases as the ionization energies increase across the transition series from Sc to Zn. Thus, the so-called early transition metals, those on the left side of the d-block (Sc through Mn) are oxidized most easily and are the strongest reducing agents. Batteries The automobile or lead storage battery is the best known example or a battery that can be recharged by applying an outside voltage to reverse the electrode reactions. The spontaneous discharging reactions are: Lead Storage Battery: A typical 12 volt battery consists of six individual cells connected in series. Anode: Lead grid packed with spongy lead. Pb(s) + HSO4(aq) PbSO4(s) + H+(aq) + 2 e Cathode: Lead grid packed with lead oxide. PbO2(s) + 3 H+(aq) + HSO4(aq) + 2 e PbSO4(s) + 2 H2O(l) Electrolyte: 38% by mass sulfuric acid. Cell Potential: 1.924 V Chapter18 Electrochemistry 17
Why rechargeable?
PbSO4(s) is formed on the electrode surface Battery life surface imperfections - inclusions dislodging of PbSO4.
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Zinc Dry-Cell: Also called a Leclanch cell, uses a viscous paste rather than a liquid solution. Anode: Zinc metal can on outside of cell. Zn(s) Zn2+(aq) + 2 e Cathode: MnO2 and carbon black paste on graphite. 2 MnO2(s) + 2 NH4+(aq) + 2 e Mn2O3(s) + 2 NH3(aq) + 2 H2O(l) Electrolyte: NH4Cl and ZnCl2 paste. Cell Potential: 1.5 V but deteriorates to 0.8 V with use.
Chapter18 Electrochemistry 19
Mercury Dry-Cell: Modified Leclanch cell which replaces MnO2 with HgO and uses a steel cathode. Anode: Zinc metal can on outside of cell. Zn(s) + 2 OH(aq) ZnO(s) + H2O(l) + 2 e Cathode: HgO in contact with steel. 2 HgO (s) + H2O(l) + 2 e Hg(l) + 2 OH(aq) Electrolyte: KOH, and Zn(OH)2 paste. Cell Potential: 1.3 V with small size, longer lasting, and stable current and voltage.
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Nickel Cadmium Batteries are rechargeable, why? Do these batteries have a memory or is it overcharging.
2
Cd(s) + 2OH (aq) Cd(OH) (s) + 2e NiO(OH)(s) + 2H2O(l) + 2e Ni(OH)2(s) + OH (aq) The following was directly copied from http://www.batteryuniversity.com/parttwo-33.htm This is quite a good site.
The word 'memory' was originally derived from 'cyclic memory'; meaning that a nickel-cadmium battery could remember how much energy was drawn on preceding discharges. On a longer than scheduled discharge, the voltage would rapidly drop and the battery would lose power. Improvements in battery technology have virtually eliminated this phenomenon. problem The with nickel-cadmium is not so much the cyclic memory but the effects of crystalline formation. The active cadmium material is present in finely divided crystals. In a good cell, these crystals remain small, obtaining maximum surface area. With memory, the crystals grow and conceal the active material from the electrolyte. In advanced stages, the sharp edges of the crystals penetrate the separator, causing high self-discharge or electrical short. When introduced in the early 1990s, nickel-metal-hydride was promoted as being memory-free. Today, we know that this chemistry is also affected but to a lesser degree than nickel-cadmium. The nickel plate, a metal that is shared by both chemistries, is partly to blame. While nickel-metal-hydride has only the nickel plate to worry about, nickel-cadmium also includes the memory-prone cadmium plate. This is a non-scientific explanation why nickelcadmium is affected more than nickel-metal-hydride. The stages of crystalline formation of a nickel-cadmium cell are illustrated in Figure 1. The enlargements show the cadmium plate in a proper functioning crystal structure, crystalline formation after use (or abuse) and restoration.
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Newnickelcadmiumcell.Theanodeisin freshcondition.Hexagonalcadmium hydroxidecrystalsareabout1micronin crosssection,exposinglargesurfaceareato theelectrolyteformaximumperformance. Cellwithcrystallineformation.Crystalshave grownto50to100micronsincrosssection, concealinglargeportionsoftheactive materialfromtheelectrolyte.Jaggededges andsharpcornersmaypiercetheseparator, leadingtoincreasedselfdischargeor electricalshort.
Restoredcell.Afterpulsedcharge,the crystalsarereducedto3to5microns,an almost100%restoration.Exerciseor reconditionareneededifthepulsecharge aloneisnoteffective.
Figure1:Crystallineformationonnickelcadmiumcell. IllustrationcourtesyoftheUSArmyElectronicsCommandinFortMonmouth,NJ,USA.
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How to restore and prolong nickel-based batteries Crystalline formation is most pronounced if a nickel-based battery is left in the charger for days, or if repeatedly recharged without a periodic full discharge. Since most applications do not use all energy before recharge, a periodic discharge to 1 volt per cell (known as exercise) is essential to prevent memory. Nickel-cadmium in regular use and on standby mode (sitting in a charger for operational readiness) should be exercised once per month. Between these monthly exercise cycles, no further service is needed. No scientific research is available on the optimal exercise requirements of nickel-metal-hydride. Based on the reduced crystalline buildup, applying a full discharge once every three months appears right. Because of the shorter cycle life compared to nickel-cadmium, over-exercising is not recommended.
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NickelMetalHydride (NiMH): Replaces toxic Cd anode with a hydrogen atom impregnated ZrNi2 metal alloy. During oxidation at the anode, hydrogen atoms are released as H2O. Recharging reverses this reaction.
Chapter18 Electrochemistry 24
Lithium Ion (Liion): The newest rechargeable battery is based on the migration of Li+ ions. Anode: Li metal, or Li atom impregnated graphite. Li(s) Li+ + e Cathode: Metal oxide or sulfide that can accept Li+. MnO2(s) + Li+(aq) + e LiMnO2(s) Electrolyte: Lithium-containing salt such as LiClO4, in organic solvent. Solid state polymers can also be used. Cell Potential: 3.0 V
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Lithium-ion batteries use a variety of cathodes and electrolytes. Common combinations use an anode of lithium (Li) ions dissolved in carbon or graphite and a cathode of lithium cobalt-oxide (LiCoO2) or lithium manganese-oxide (LiMn2O4) in an liquid electrolyte of lithium salt. Because they use a liquid electrolyte, lithium-ion batteries are limited in shape to either prismatic (rectangular) or cylindrical. The cylindrical form has a similar construction to other cylindrical rechargeable batteries, as at right. Prismatic batteries have the anode and cathode inserted into the rectangular enclosure. The image link at right illustrates this construction method. Lithium-Ion-Polymer batteries are the next stage in development and replace the liquid electrolyte with a plastic (or polymer) electrolyte. This allows the batteries to be made in a variety of shapes and sizes. The significant advantages of lithium-ion batteries are size, weight and energy density (the amount of power the battery can provide). Lithium-ion batteries are smaller, lighter and provide more energy than either nickel-cadmium or nickelmetal-hydride batteries. Additionally, lithium-ion batteries operate in a wider temperature range and can be recharged before they are fully discharged without creating a memory problem. As with most new technology, the disadvantage is pricing. Currently, lithium-ion and lithium-ion-polymer batteries are more expensive to manufacture than standard rechargeable batteries. Part of this expense is due to the volatile nature of lithium. Lithium-ion batteries are most commonly used in applications where one or more of the advantages (size, weight or energy) outweigh the additional cost, such as mobile telephones and mobile computing devices. Lithium-ion-polymer batteries are used when the battery needs to be a particular shape. http://support.radioshack.com/support_tutorials/batteries/bt-liion-main.htm Chapter18 Electrochemistry 26
2 H2(g) + 4OH(aq) 4H2O(l) + 4e O2 (g) + 2 H2O(l) + 4 e 4 OH 2 H2(g) + O2 (g) 2 H2O(l)
A hydrogen fuel cell consists of two electrodes sandwiched around an electrolyte. Oxygen passes over one electrode and hydrogen over the other, generating electricity, water and heat. Hydrogen is fed into the "anode" of the fuel cell. Oxygen (or air) enters the fuel cell through the cathode. Excited by a catalyst, the hydrogen atom splits into a proton and an electron, which take different paths to the cathode. The proton passes through the electrolyte. The electrons create a separate current that can be utilized before they return to the cathode, to be reunited with the hydrogen and oxygen in a molecule of water. Since hydrogen fuel cells rely on chemistry and not combustion, its emissions are virtually zero in comparison to the cleanest fuel combustion engines. Hydrogen fuel cells can be made in a vast quantity of sizes. They can be used to produce small amounts of electric power for devices such as personal computers, or be used to produce high voltage powers for electric power stations. Hydrogen fuel cell vehicles are an attractive alternative to regular battery-powered vehicles. They can be refueled quicker and even run longer between refueling. Copied from the web site:http://4hydrogen.com/about.html 27
Chapter18 Electrochemistry
Corrosion
Corract.swf
Tiny galvanic cell
Must have oxygen and water
2+ o
Fe(s) Fe + 2e
2 + 2
E = 0.45
o
Chapter18 Electrochemistry
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Why does salt increase the rate of corrosion?
Aluminum and titanium are more readily oxidized than iron but we use them instead of iron to stop rusting, why?
Chapter18 Electrochemistry
29
How can you help prevent corrosion?
Cathodic protection: By connecting the metal to a second metal that is more readily oxidized you are sacrificing the second metal Galvanizing is coating a metal with zinc cathodic protection plus protective ZnO2 layer Many rust proofing kits use zinc phosphate on the rusty surface (after scraping off loose rust) before fiber glassing or using sheet metal to fill holes
Sacrificial anode: connecting metal with a wire to a more readily oxidizable metal for example an underground pipeline is protected by connecting it through an insulated wire to a stake of magnesium which acts as a sacrificial anode
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AnodeFlex 1500 Impressed Current System Anodeflex 1500-1 is a long-line, flexible, cable-like anode, which is placed in continuous close proximity to the target structure. Uniform distribution of cathodic protection current is therefore achieved on applications where many conventional anode ground beds do not work. Key to the products performance is the central, conductive-polymer coated copper conductor.
Durichlor 51 Button and Bullet
Button and bullet anodes are designed specifically for use in water environments where a small and secure anode is preferred. Button anodes are most often used on locks, dams, and other structures where a low profile anode is required. The bullet anode, which is mounted on a metal rod, is ideally suited for use in condenser water boxes. Both anode types operate efficiently in fresh, brackish, and saltwater environments.
ELGARD 100 Ribbon Mesh
ELGARD Anode Ribbon Mesh is a key component for Cathodic Protection systems in reinforced concrete structures. It is composed of a precious metal oxide catalyst sintered to an expanded titanium mesh substrate.
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Pier, Piling, and Ballast Tank
Cast Aluminum Anodes
Aluminum anodes are used more often than any other type of galvanic material in saltwater environments due to their relatively low consumption rate of approximately 7 lbs. per amp.-year. To meet the demand for this effective anode material, Corrpro manufactures an extensive array of aluminum anodes in two different alloys.
Brochure
Soil and Ribbon
Cast and Extruded Zinc Anodes
Zinc anodes are widely utilized and have proven to be an effective choice for preventing corrosion in select soils and brackish waters. For these environments, Corrpro offers zinc anodes made to the ASTM B-418, Type II alloy standard.
Mixed Metal Oxide Anode Technology
ELTECH LIDA anodes are comprised of a titanium substrate with a mixed metal oxide coating. The mixed metal oxide is a crystalline, electrically conductive coatin g that activates the titanium and enables it to function as an anode. This coating has an extremely low consumption rate, measured in terms of milligrams per year. As a result of this low consumption rate, the anode dimensions remain nearly constant during the life of the anode, providing constant levels of performance for the duration of the anode design life. Due to the ductility of the titanium substrate, a wide range of anode shapes suitable to the structure to be protected are possible, such as wire, rod, tubular, disk and mesh configurations.
Electrolytic cells 18.11 Read over for interest but will not be tested.
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Electrolysis and Electrolytic Cells
Electric current is used to drive a nonspontaneous reaction An electrolytic cell converts electrical energy to chemical energy Drives the reaction away from equilibrium.
Electrolysis of molten Sodium Chloride
The battery serves as an electron pump, pushing electrons into the one electrode and pulling them out of the other. The negative electrode attracts Na+ cations, which combine with the electrons supplied by the battery and are thereby reduced to liquid sodium metal The positive electrode attracts Cl- anions, which replenish the electrons removed by the battery and are thereby oxidized to chlorine gas.
2 -
Anode (Oxidation)
2Cl (l) Cl (g) + 2e
+ -
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As in the galvanic cell, the anode is the electrode where oxidation takes place, and the cathode is the electrode where reduction takes place, The sign of the electrodes, however, are opposite for the two kinds of cells. In a galvanic cell, the anode is considered negative because it supplies electrons to the external circuit, but in an electrolytic cell, the anode is considered positive because electrons are pulled out of it by the battery.
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Example: A copper sulfate solution, CuSO4(aq), is electrolyzed for 7.00 minutes using an external current of 0.60 amperes. (1 ampere = 1 coulomb per second) What mass of Cu(s) is produced? 1 mole e = 96485 C F = 96 485 C/mol The cathode reaction is: Cu2+(aq) + 2e Cu(s) E = 0.339 V
The battery draws electrons away from the anode. If the anode is chosen to be an inert material like platinum, the what reaction is supply these electrons? The oxidation of water. 2 H2O(l) O2(g) + 4 H+(aq) + 4 e E = - 1.23 V
Without the battery, Ecell = -0.891 V. In order for Cu2+ to be reduced, we need a battery with a voltage significantly greater than 0.891 V Total charge transfer: Q =
# moles e transferred =
# mol Cu = Mass of Cu = Chapter18 Electrochemistry 35