Some helpful links for further information on electrochemistry:
Chemmybear
http://www.chemmybear.com/groves/apchem.html
Purdue.edu--Electrochemical Reactions
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/electro.php
Electrochemistry Dictionary and Encyclopedia
http://electrochem.cwru.edu/ed/dict.htm
Thursday, May 30, 2013
Commercial Electrolytic Processes
Some commercial processes are the production of aluminum, the electrorefining of metals, metal plating, and electrolysis of sodium chloride.
Since aluminum has such a high affinity for oxygen and is never found pure, electrolysis can be used to produce 99.5% pure aluminum by separating the oxides.
Purification of metals is another application of electrolysis and is important for other chemical processes.
Metals that readily corrode can be protected by a thin layer of a metal that resists corrosion and electrolysis can help provide a smooth, even coating of the metal.
Sodium metal is mainly produced by the electrolysis of molten sodium chloride.
Since aluminum has such a high affinity for oxygen and is never found pure, electrolysis can be used to produce 99.5% pure aluminum by separating the oxides.
Purification of metals is another application of electrolysis and is important for other chemical processes.
Metals that readily corrode can be protected by a thin layer of a metal that resists corrosion and electrolysis can help provide a smooth, even coating of the metal.
Sodium metal is mainly produced by the electrolysis of molten sodium chloride.
Electrolysis
An electrolytic cell uses electrical energy to produce chemical change. Electrolysis involves forcing a current through a cell to produce a chemical change for which the cell potential is negative (means that electrical work causes an otherwise nonspontaneous chemical reaction to occur). Electrolysis involves stoichiometry regarding how much chemical change occurs with the flow of a given current for a specified time. Plating means depositing the neutral metal on the electrode by reducing the metal ions in solution.
Electrolysis's common uses are to place a thin coating of metal onto steel and to produce pure metals such as aluminum and copper.
Electrolysis's common uses are to place a thin coating of metal onto steel and to produce pure metals such as aluminum and copper.
Corrosion
Corrosion can be viewed as the process of returning metals to their natural state, which is the ores from which they were originally obtained. Corrosion involves oxidation of the metal. The metals that have standard reduction potentials less positive than that of oxygen gas, when the half-reactions are reversed and combined with the reduction half-reaction for oxygen, the result is a positive E value. This shows that the oxidation of most metals is spontaneous.
A lot of metals form a thin oxide coating which protects their internal atoms from further oxidation. However, to actively prevent corrosion, we can apply a coating (most commonly a paint of metal plating) to protect the metal from oxygen and moisture. A process called galvanizing is used which involves coating steel with zinc to form a mixed oxide-carbonate coating. Alloying is also used to prevent corrosion, which involves mixing materials to help protect them. Another method called cathodic protection connects an active metal by a wire to the pipeline or tank to be protected.
A lot of metals form a thin oxide coating which protects their internal atoms from further oxidation. However, to actively prevent corrosion, we can apply a coating (most commonly a paint of metal plating) to protect the metal from oxygen and moisture. A process called galvanizing is used which involves coating steel with zinc to form a mixed oxide-carbonate coating. Alloying is also used to prevent corrosion, which involves mixing materials to help protect them. Another method called cathodic protection connects an active metal by a wire to the pipeline or tank to be protected.
Batteries
A battery is a galvanic cell or (more commonly) a group of galvanic cells connected in series, where the potentials of the individual cells add to give the total battery potential.'
The lead storage battery can function for several years under temperature extremes. In this battery, lead serves as the anode and lead coated with lead dioxide serves as the cathode. Both electrodes dip into an electrolyte solution of sulfuric acid. The reactions are:
Anode: Pb + HSO4- ----> PbSO4 + H+ + 2e-
Cathode: PbSO4 + HSO4- + 3H+ + 2e- ----> PbSO4 + 2H20
Because sulfuric acid is consumed as the battery discharges, the condition of the battery can be monitored by measuring the density of the sulfuric acid solution. A car with a dead battery can be jump-started by connecting its battery to the battery in a running automobile.
Also, traditional types of storage batteries require periodic "topping off" because the water in the electrolyte solution is depleted by the electrolysis that accompanies the charging process.
Another common battery is the dry cell battery which has an acid version and an alkaline version. In the acid version, the battery contains a zinc inner case that acts as the anode and a carbon rod in contact with a paste of solid MnO2, solid NH4Cl, and carbon that acts as the cathode. In the alkaline version of the battery, the solid NH4Cl is replaced with KOH or NaOH. The alkaline version lasts longer because the zinc anode corrodes less rapidly under basic conditions than acidic conditions.
There are also silver cells (has a Zn anode and a cathode with Ag2O as the oxidizing agent), mercury cells (Zn anode and a cathode with HgO as the oxidizing agent).
A fuel cell is a galvanic cell for which the reactants are continuously supplied.
The lead storage battery can function for several years under temperature extremes. In this battery, lead serves as the anode and lead coated with lead dioxide serves as the cathode. Both electrodes dip into an electrolyte solution of sulfuric acid. The reactions are:
Anode: Pb + HSO4- ----> PbSO4 + H+ + 2e-
Cathode: PbSO4 + HSO4- + 3H+ + 2e- ----> PbSO4 + 2H20
Because sulfuric acid is consumed as the battery discharges, the condition of the battery can be monitored by measuring the density of the sulfuric acid solution. A car with a dead battery can be jump-started by connecting its battery to the battery in a running automobile.
Also, traditional types of storage batteries require periodic "topping off" because the water in the electrolyte solution is depleted by the electrolysis that accompanies the charging process.
Another common battery is the dry cell battery which has an acid version and an alkaline version. In the acid version, the battery contains a zinc inner case that acts as the anode and a carbon rod in contact with a paste of solid MnO2, solid NH4Cl, and carbon that acts as the cathode. In the alkaline version of the battery, the solid NH4Cl is replaced with KOH or NaOH. The alkaline version lasts longer because the zinc anode corrodes less rapidly under basic conditions than acidic conditions.
There are also silver cells (has a Zn anode and a cathode with Ag2O as the oxidizing agent), mercury cells (Zn anode and a cathode with HgO as the oxidizing agent).
A fuel cell is a galvanic cell for which the reactants are continuously supplied.
Dependence of Cell Potential on Concentration
In previous sections when describing cells, we have assumed standard conditions (all concentrations 1 M). If concentrations of any substance are greater than 1 M, we can answer qualitatively with Le Chatelier's principle. An increase in concentration of a reactant will favor the forward reaction and thus increase the driving force on the electrons, which increases the cell potential. An increase in concentration of a product will oppose the forward reaction and thus decrease the driving force on the electrons, which decreases the cell potential.
A certain type of galvanic cell called a concentration cell can be constructed. This is a cell in which both compartments have the same components but at different concentrations. To solve this problem, we need to recognize that nature will try to equalize the concentrations in the compartments and transfer electrons from the compartment of lower concentration to the compartment of higher concentration.
Another equation that is necessary is the Nernst equation, which is:
E = E° - RT/nF ln (Q)
A certain type of galvanic cell called a concentration cell can be constructed. This is a cell in which both compartments have the same components but at different concentrations. To solve this problem, we need to recognize that nature will try to equalize the concentrations in the compartments and transfer electrons from the compartment of lower concentration to the compartment of higher concentration.
Another equation that is necessary is the Nernst equation, which is:
E = E° - RT/nF ln (Q)
Wednesday, May 29, 2013
Cell Potential, Electrical Work, and Free Energy
The work that can be accomplished when electrons are transferred through a wire depends on the "push" (the thermodynamic driving force) behind the electrons. This driving force (the emf) is defined in terms of a potential difference (in volts) between two points in the circuit.
emf=potential difference (V) = work (J) / charge (C)
The term "work" is viewed from the point of view of the system. Work flowing out of the system is a minus sign and when a cell produces a current, the cell potential is positive and the current can be used to do work. Thus, cell potential and work (w) have opposite signs.
E = -w (work) / q (charge)
and wmax = -qEmax
However, because current must always flow to obtain work and when current flows energy is wasted, the maximum work is not obtained. In any real, spontaneous process some energy is always wasted--the actual work realized is always less than the calculated maximum. So the actual work done is:
w = -qE
where E represents the actual potential difference at which the current flowed and q is the quantity of charge in coulombs transferred. The faraday is the charge on 1 mole of electrons and is abbreviated F. It has the value of 96,485 coulombs of charge per mole of electrons. Thus q equals the number of moles of electrons times the charge per mole of electrons:
q = nF
When we want to relate the potential to free energy, we see that the change in free energy equals the maximum useful work obtainable from that process:
ΔG = -nFE
which states that the maximum cell potential is directly related to the free energy difference between the reactants and the products in the cell.
emf=potential difference (V) = work (J) / charge (C)
The term "work" is viewed from the point of view of the system. Work flowing out of the system is a minus sign and when a cell produces a current, the cell potential is positive and the current can be used to do work. Thus, cell potential and work (w) have opposite signs.
E = -w (work) / q (charge)
and wmax = -qEmax
However, because current must always flow to obtain work and when current flows energy is wasted, the maximum work is not obtained. In any real, spontaneous process some energy is always wasted--the actual work realized is always less than the calculated maximum. So the actual work done is:
w = -qE
where E represents the actual potential difference at which the current flowed and q is the quantity of charge in coulombs transferred. The faraday is the charge on 1 mole of electrons and is abbreviated F. It has the value of 96,485 coulombs of charge per mole of electrons. Thus q equals the number of moles of electrons times the charge per mole of electrons:
q = nF
When we want to relate the potential to free energy, we see that the change in free energy equals the maximum useful work obtainable from that process:
ΔG = -nFE
which states that the maximum cell potential is directly related to the free energy difference between the reactants and the products in the cell.
Standard Reduction Potentials
The reaction in a galvanic cell is always a redox reaction that can be broken down into two half-reactions. We can determine the cell potential by assigning a potential to each half-reaction and summing them.
Standard reduction potentials are the E° values corresponding to reduction half-reactions with all solutes at 1 M and all gases at 1 atm.
Combining the half-reactions to get a balanced redox reaction often requires one of the reactions to be reversed (in order to cancel out the electron transfers and since redox reactions must involve a substance being oxidized and a substance being reduced). Also, the half-reaction with the largest possible potential (the reduction) will not be the reaction that is reversed. The other half-reaction will be forced to run in reverse (the oxidation). During this manipulation, the net potential of the cell will be the difference between the two. (cathode - anode) Also, because subtraction means "change the sign and add", we will change the sign of the oxidation (anode) reaction when we reverse it and add it to the reduction (cathode) reaction.
Another possible manipulation is multiplying the half-reaction by integers to make the number of electrons lost equal the number gained. However, when this is done, the value of E° is not changed. Since standard reduction potential is an intensive property, the potential is not multiplied by the integer.
Another question is which reaction must run in reverse. This can be answered by considering the sign of the potential of a working cell. A cell will always run spontaneously in the direction that produces a positive cell potential.
A complete description of a galvanic cell usually contains these four items:
-the cell potential
-the direction of electron flow (obtained by inspecting the half-reactions and using the direction that gives a positive Ecell)
-designation of the anode and cathode
-the nature of each electrode and the ions present in each compartment
Standard reduction potentials are the E° values corresponding to reduction half-reactions with all solutes at 1 M and all gases at 1 atm.
Combining the half-reactions to get a balanced redox reaction often requires one of the reactions to be reversed (in order to cancel out the electron transfers and since redox reactions must involve a substance being oxidized and a substance being reduced). Also, the half-reaction with the largest possible potential (the reduction) will not be the reaction that is reversed. The other half-reaction will be forced to run in reverse (the oxidation). During this manipulation, the net potential of the cell will be the difference between the two. (cathode - anode) Also, because subtraction means "change the sign and add", we will change the sign of the oxidation (anode) reaction when we reverse it and add it to the reduction (cathode) reaction.
Another possible manipulation is multiplying the half-reaction by integers to make the number of electrons lost equal the number gained. However, when this is done, the value of E° is not changed. Since standard reduction potential is an intensive property, the potential is not multiplied by the integer.
Another question is which reaction must run in reverse. This can be answered by considering the sign of the potential of a working cell. A cell will always run spontaneously in the direction that produces a positive cell potential.
A complete description of a galvanic cell usually contains these four items:
-the cell potential
-the direction of electron flow (obtained by inspecting the half-reactions and using the direction that gives a positive Ecell)
-designation of the anode and cathode
-the nature of each electrode and the ions present in each compartment
Subscribe to:
Posts (Atom)