Electrochemical Changes
The terms oxidation and reduction can be used to describe several chemical reactions. For example, in ancient times, large amount of metal ores were reduced to smaller amount of pure metal, a process that frequently required the removal of oxygen. Oxidation, in this context, would refer to the addition of oxygen to create metal oxides.
Reduction oxidation reactions can also be defined from the electron transfer model. Consider a single replacement reaction of Zinc and HCL to form ZnCl2 and Hydrogen gas as shown below.
Zn(s) + 2 HCl(aq) -> ZnCl2(aq) + H2(g)
We can split this reaction into two half-reactions. A half-reaction represents what is happening to one reactant in the reaction.
Zn(s) -> Zn2+(aq) + 2 e-
2 H+(aq) + 2 e- -> H2(g)
The top half reaction shows Zn losing electrons to become Zn2+, while the 2 H+ gains 2 electrons to become H2(g).
The gain of electrons is called Reduction, and the loss of electrons is called Oxidation. Chemical reactions where exchange of electrons occurs are called redox reactions. To memorize these you can use the acronyms OIL RIG or LEO the lion says GER
Consider the reaction below:
1Fe2O3 (s) + 3CO (g) -> 2Fe (s) + 3CO2 (g)
In this reaction, CO is used to reduce Fe2O3 into pure Fe (s). We can therefore refer to CO as a reducing agent. The reducing agent donates electrons to the iron ion to reduce it into solid iron. An oxidizing agent is generally the opposite. rather than donating electrons, it takes/receives electrons.
Oxidation occurs when metals are converted into metallic compounds by reacting with nonmetals. For example:
2Mg (s) + 1O2 (g) -> 2MgO (s)
1Cu (s) + 1Br2 (g) -> 1CuBr2 (s)
Redox reactions involve the conversion or atoms to ions, and ions back to atoms simultaneously in the same reaction. Consider the following example
1Cu (s) + 2AgNO3 (aq) -> 2Ag (s) + 1Cu(NO3)2 (aq)
In this reaction, Cu (s) loses electrons (oxidation) to become Cu2+. Ag+ gain electrons (reduction) to become Ag(s).
We can summarize this redox reaction by writing it this way:
1Cu (s) + 2Ag+ (aq) -> 2Ag (s) + 1Cu+ (aq)
The Ag+ have been reduced to a metal and the Cu has been oxidized to an ion. Notice that the NO3- is not involved in the actual redox reaction. This is called the spectator ion. The two half reactions can be written as follows:
1Ag+(aq) + 1 e- -> 1Ag0(s)
1Cu0(s) -> 1Cu2+(aq) + 2 e-
Reduction: Reduction is defined as the gain of electrons or loss of oxygen, The substance that gets reduced is the oxidizing agent.
Oxidation: Oxidation is the loss of electrons or the addition of oxygen. The substance thats gets oxidized is the reducing agent.
Redox reactions involve one substance undergoing reduction and the other undergoing oxidation. The substance undergoing reduction is the oxidizing agent, and the substance undergoing oxidation is the reducing agent. As you will recall from bonding theory, different substances have differing levels in their ability to act as oxidizing or reducing agents. For example, if you have 4 ions (Ag+, Cu2+, Mg2+ and Zn2+) the activity of oxidizing agent reduces from Ag+ (highest activity) to Zn2+ (lowest activity). On the other hand, the same elements, but as metals, have a reversed order of activity as reducing agents. Therefore, Zn(s) has the highest reducing activity and Ag(s) has the lowest reducing activity.
Reactions are spontaneous if the oxidizing agent (OA) is ABOVE the reducing agent (RA) in a redox 1/2 reaction table, and non-spontaneous if the reducing agent is above the oxidizing agent.
Predicting the oxidizing agent activity of 4 ions
Predicting the reducing agent activity of 4 metals
A redox table is a series of reduction reactions, with the strongest oxidizing agent at the top left and the strongest reducing agent at the bottom right of the table.
Standard Reduction Potential, or E0, is a measure of the potential of a half-reaction to reduce at standard conditions (1M and 1atm) measured in volts. Standard reduction potentials are calculated by comparing them to H2 so that if it can reduce better than H2 it will have a positive sign, and vice versa. Below is a table showing standard reduction potentials for given half reactions.
If E0 is positive, the forward reaction is spontaneous; if E0 is negative the reverse reaction is spontaneous.
Some complex reactions cannot be explained using the redox theory alone. Chemists have developed a method to account for the number of electrons either lost or gained in a molecule, complex ion and also in chemical reactions. The oxidation state is defined as the net electric charge that it would have if electron pairs in covalent bonds belonged entirely to the more electronegative atom. This needs to be differentiated from oxidation number, which is a positive or negative number corresponding to the oxidation state assigned to an atom in a covalently bonded entity, based on arbitrary rules. The sum of the oxidation numbers for a compound (neutral) is zero. The sum of the oxidation numbers in a polyatomic ion is equal to the charge on the ion.
An increase in oxidation number is an oxidation. A decrease in oxidation number is a reduction. If the oxidation numbers do not change, then a redox reaction has not occurred (i.e. no transfer of electrons).
Covalent compounds are made of two nonmetals, which from the periodic table are always expected to be negative. Since covalent bonds are neutral, only the more electronegative element keeps its negative oxidation number. The other nonmetals adapt their oxidation number to keep the compound neutral.
Examples of oxidation numbers in ionic and covalent compounds.
Steps in Balancing chemical equation using oxidation numbers
In redox stoichiometry, the titrant is always a strong oxidizing or reducing agent. Two common oxidizing agents include acidic solutions of permanganate ions or dichromate ions.
Example Question:
If 6.75 mL of acidified 0.100 M KMnO4 is required to titrate 25.0 mL of FeCl2. Calculate the concentration of Fe2+.
In the previous chapter, we showed that oxidation reduction reactions involve the movement of electrons with one reactant undergoing oxidation (losing electrons) and the other reactant undergoing reduction (gaining electrons).
Spontaneous redox reactions can be made useful to generate electricity. Electrochemical cells convert this chemical energy into electrical energy that can be used to run various functions such as a battery operated flashlight. A battery consists of two or more cells connected in parallel, series or 'series and parallel' pattern. This enables the generation of a larger voltage than a single cell.
An electrochemical cell is a device capable of either generating electrical energy from chemical reactions or using electrical energy to cause chemical reactions. The electrochemical cells which generate an electric current are called voltaic cells or galvanic cells and those that generate chemical reactions, through electrolysis for example, are called electrolytic cells.
Electricity is measured using a voltmeter. Voltage is defined as the energy difference between two points on an electrical circuit.
Voltaic cells (or Galvanic cells) (named after Luigi Galvani or Alessandro Volta) are electrochemical cells that produces electricity from conductors placed in conducting solutions. Voltaic cells generally consists of two different metals connected by a salt bridge, or individual half-cells separated by a porous membrane.
An illustration of a Voltaic / Galvanic cell. (Source: Wikipedia-CC BY-SA 3.0)
Voltaic cells can either be primary or secondary cells. A primary cell is designed to be used once and discarded. A secondary cell is a rechargeable battery. In primary cells, the electrochemical reaction occurring in the cell is not reversible, the chemicals gets used up. In a secondary cell, the reaction can be reversed by running a current into the cell using a battery charger to, this regenerates the chemical reactants and can generate electricity again. A common secondary cell is the Lead-acid battery commonly used in car batteries Lead acids batteries are common because they can generate a high voltage, they are low costs, reliable and tend to last a long time.
Amperes is a measure of rate of flow (i.e. current) that passes a specific point on an electrical circuit. On the other hand, Coulombs (C) expresses total charge transferred by movement of charged particles.
A voltaic cell consists of a solid metal (an electrode) that is submerged in a solution that contains cations (+ve) of the electrode metal and anions (-ve) to balance the charge of the cations. The full cell consists of two half-cells, connected by a semi-permeable membrane or by a salt bridge that prevents the ions of the more noble metal from plating out at the other electrode. The reduction reaction occurs at the cathode (marked positive), while oxidation occurs at the Anode (marked negative). The strongest oxidizing agent present in the cell always undergoes a reduction at the cathode.
The salt bridge device used to connect the oxidation and reduction half-cells. It maintains electrical neutrality within the internal circuit, preventing the cell from rapidly running its reaction to equilibrium. If no salt bridge is absent, one half cell would accumulate negative charge and the solution in the other half cell would accumulate positive charge as the reaction proceeded, quickly preventing further reaction, and as a result limit the production of electricity. A last bridge can be a U shaped glass tube filled with a relatively inert electrolyte, usually sodium chloride or potassium chloride. This electrolyte is often made into a gel using agar to prevent intermixing of the fluids. The electrolyte does not react with the chemicals in the cell.
A galvanic cell whose electrodes are zinc and copper submerged in zinc sulfate and copper sulfate, respectively, is known as a Daniell cell. The half reactions are:
Zinc electrode (anode) oxidation: Zn(s) -> Zn2+(aq) + 2 e-
Copper electrode (cathode): Cu2+(aq) + 2 e- -> Cu(s)
This reaction can also be driven in reverse by applying a voltage, resulting in the deposition of zinc metal at the anode and formation of copper ions at the cathode.
Inert electrodes are used in voltaic cells where the oxidizing or reducing agent is not a solid metal. They provide a location to connect a wire and a surface on which a half-reaction can occur.
Also called the cell potential, is the maximum electric potential difference of a cell i.e. between the anode and cathode. To predict the cell potential, tabulations of standard electrode potential are available. Calculations are referenced to the standard hydrogen electrode (SHE) which is arbitrarily given a potential of 0.00 V. The standard hydrogen electrode undergoes the reaction:
2 H+(aq) + 2 e- -> H2
Standard electrode potentials. (Source: Wikipedia-CC BY-SA 3.0)
The standard reduction potential is expressed as Eo. However, because it can also encompass both reduction and oxidation, the term reduction potential (Err) or oxidation potential (Eoo) are more preferred. It can be explained as the half cell with greater attraction for e- (more "+" reduction potential) takes e- from the one with lower reduction potential.
ΔEo) = Err(Cathode) - Err(Anode)
A positive cell potential (Eo > 0) indicates that the net reaction is spontaneous. This must be the case in all voltaic cells. You need to be able to determine which electrode is the cathode and which is the anode.
Corrosion is an electrochemical process in which a metal is oxidized by substances in the environment, returning the metal to an ore-like state. In general, any metal appearing below the various oxygen half-reactions in a redox table will be oxidized in our environment for example the rusting of iron.
Cathode O2(g) + 2H2O(l) + 4 e- -> 4 OH -(aq)
Anode 2[Fe(s) -> Fe2+ (aq) + 2 e-]
Net 2Fe(s) ) O2 (g) + 2H2O(l) -> Fe(OH)2 (s) .
The iron(II) hydroxide precipitate is further oxidized to eventually form rust: Fe2O3.xH2O(s).
Rust, the most common type of corrosion. (Source: Wikipedia-CC BY-SA 3.0)
1. Cathodic Protection (sacrificial anode): Cathodic protection (CP) is a technique to control the corrosion of a metal surface by making that surface the cathode of an electrochemical cell. Cathodic protection systems are most commonly used to protect steel pipelines and tanks; steel pier piles, ships, and offshore oil platforms. A sacrificial anode is a metal that is more easily oxidized than iron (such as magnesium) and connected to the iron object to be protected. An impressed current is an electric current forced to flow toward an iron object by an external potential difference which is provided by a constant power supply. This is commonly used for pipelines but similar systems can also be used for cars and boats.
2. Protective coatings
a. Galvanizing metals: Done by dipping iron in Zn. The zinc oxide outer layer is sacrificed to save iron inner part of nail. This is common in construction nails.
b. Plating, painting, and the application of enamel are the most common anti-corrosion treatments. They work by providing a barrier of corrosion-resistant material between the damaging environment and the structural material. Aside from cosmetic and manufacturing issues, there may be tradeoffs in mechanical flexibility versus resistance to abrasion and high temperature. Other metals such as tin and zinc are most commonly used. Tin adheres well to the surface of iron and the outer surface of the tin coating has a thin, strongly adhering layer of tin oxide that protects the tin.
3. Reactive coatings: If the environment is controlled (especially in recirculating systems), corrosion inhibitors can often be added to it. These chemicals form an electrically insulating or chemically impermeable coating on exposed metal surfaces, to suppress electrochemical reactions. Chemicals that inhibit corrosion include some of the salts in hard water, chromates, phosphates, polyaniline, other conducting polymers and a wide range of specially designed chemicals that resemble surfactants.
4. Anodization: Aluminium alloys often undergo a surface treatment. Anodizing is very resilient to weathering and corrosion, so it is commonly used for building facades and other areas where the surface will come into regular contact with the elements.
5. Biofilm coatings: A new form of protection has been developed by applying certain species of bacterial films to the surface of metals in highly corrosive environments. This process increases the corrosion resistance substantially.
As a review, electrochemical cells are spontaneous because the strongest oxidizing agent is always above the strongest reducing agent. The cell potential is positive.
On the other hand, Electrolytic cells are non-spontaneous, the strongest oxidizing agent (SOA) is always below the strongest reducing agent (SRA) in a redox table. The cell potential is negative. Electrons are pulled from the anode to the cathode by a battery or power supply. In contrast to electrochemical cells (which consist of two half reactions in separate locations) both electrodes are in the same electrolyte. Electrolysis can therefore be summarized as the process of supplying electrical energy to force a nonspontaneous reaction.
Electrolysis of water
Electrolysis of water is the decomposition of water into oxygen and hydrogen gas (water splitting) due to the passage of an electric current. This technique can be used to make hydrogen gas, a main component of hydrogen fuel, and breathable oxygen gas, or can mix the two into oxyhydrogen, which is also usable as fuel, though more volatile and dangerous.
The reactions are as follows:
Anode: 2H2(l) -> O2 (g) + 4H+ (aq) + 4 e-
Cathode: 2H2O (g) + 2 e- -> H2 (g) + 2H- (aq)
Net: 2H2O (l) -> 2 H2 (g) + O2 (g)
A simple illustration of the electrolysis of water, to release H2 and O2 gases. (Source: Wikipedia-CC BY-SA 3.0)
The Potassium Iodide Electrolytic Cell
In this case, the strong oxidizing agent is water, while the strong reducing agent is I-.
2H2O(l) + 2e- -> H2(g) + 2 OH-(aq)
2I-(aq) -> I2 (s) + 2e-
Overall: 2H2O(l) + 2I-(aq) -> H2 (g) + 2OH- (aq) + I2 (s)
The Chloride Anomaly: For the electrolysis of solutions containing the chloride ion (e.g. brine), the chloride ion acts as the strongest reducing agent, even though the table indicates that it is not.
A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning. For example the Aluminium-Air Cell. Air is pumped into the cell and oxygen is reduced at the cathode while aluminum is oxidized at the anode. This type of cell has been developed for possible use in electric cars.
A secondary cell can be recharged using electricity to reverse the chemical reaction that occurs when electricity is produced by the cell. When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal.
The Lead acid cell is the most common example of secondary cells
Most elements occur naturally combined with other elements in compounds. These mixtures are called Ores. The production of active metals (strong reducing agents) from their minerals typically involves the electrolysis of molten compounds of the metal.
Electrorefining is the process of using an electrolytic cell to obtain high-grade metals at the cathode from an impure metal at the anode.
Electrowinning is the process of using an electrolytic cell to reduce metal cations from a molten or aqueous electrolyte at the cathode.
The most important nonmetal produced by electrolysis is chlorine. The chlor-alkali process is the electrolysis of brine (NaCl(aq)) to produce chlorine, hydrogen, and sodium hydroxide.
Illustration of refining metals (Source: Wikipedia-CC BY-SA 3.0)
If molten silver ion is reduced in the following formula:
Ag+(l) + e- -> Ag(s)
Therefore 1 mol e- -> 1 mol Ag
or Fe2+(l) + 2e- -> Fe(s)
or Al3+(l) + 3e- -> Al(s)
Faraday's laws of electrolysis: They state that the amount of material produced at an electrode (or liberated from it) during an electrochemical reaction is directly proportional to the total conducted charge or, equivalently, the average current multiplied by the total time.
Therefore, the amount of moles of e- can be calculated as;
Current = charge passing / time
I (amps) = coulombs / second
The charge on 1 mol of e- is (1.60 x 10-19 C/e-) x (6.02 x 1023 e-/mol) = 96500 C, also called 1 Faraday
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