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PUSPITA WARDANI/ 29/ XII IPA 2
REDOX
A) DEFINITION REDOX
Redox (shorthand for reduction-oxidation reaction) describes all chemical reactions in which atoms have their oxidation number (oxidation state) changed. This can be either a simple redox process such as the oxidation of carbon to yield carbon dioxide or the reduction of carbon by hydrogen to yield methane (CH4), or it can be a complex process such as the oxidation of sugar in the human body through a series of complex electron transfer processes.
The term redox comes from the two concepts of reduction and oxidation. It can be explained in simple terms:
• Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom or ion.
• Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom or ion.
Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation and reduction properly refer to a change in oxidation number — the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation number, and reduction as a decrease in oxidation number. In practice, the transfer of electrons will always cause a change in oxidation number, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds).
B) HOW TO BALANCE
As a summary, here are the steps to follow to balance a redox equation in acidic medium (add the starred step in a basic medium):
1. Divide the equation into an oxidation half-reaction and a reduction half-reaction
2. Balance these
o Balance the elements other than H and O
o Balance the O by adding H2O
o Balance the H by adding H+
o Balance the charge by adding e-
3. Multiply each half-reaction by an integer such that the number of e- lost in one equals the number gained in the other
4. Combine the half-reactions and cancel
5. **Add OH- to each side until all H+ is gone and then cancel again**
In considering redox reactions, you must have some sense of the oxidation number (ON) of the compound. The oxidation number is defined as the effective charge on an atom in a compound, calculated according to a prescribed set of rules. An increase in oxidation number corresponds to oxidation, and a decrease to reduction. The oxidation number of a compound has some analogy to the pH and pK measurements found in acids and bases -- the oxidation number suggests the strength or tendency of the compound to be oxidized or reduced, to serve as an oxidizing agent or reducing agent. The rules are shown below. Go through them in the order given until you have an oxidation number assigned.
1. For atoms in their elemental form, the oxidation number is 0
2. For ions, the oxidation number is equal to their charge
3. For single hydrogen, the number is usually +1 but in some cases it is -1
4. For oxygen, the number is usually -2
5. The sum of the oxidation number (ONs) of all the atoms in the molecule or ion is equal to its total charge.
To balance redox reactions, assign oxidation numbers to the reactants and products to determine how many moles of each species are needed to conserve mass and charge. First, separate the equation into two half-reactions, the oxidation portion and the reduction portion. This is called the half-reaction method of balancing redox reactions or the ion-electron method. Each half-reaction is balanced separately and then the equations are added together to give a balanced overall reaction. We want the net charge and number of ions to be equal on both sides of the final balanced equation.
For this example, let's consider a redox reaction between KMnO4and HI in an acidic solution:
MnO4- + I- → I2 + Mn2+
Separate the two half reactions:
I- → I2
MnO4- → Mn2+
To balance the atoms of each half-reaction, first balance all of the atoms except H and O. For an acidic solution, next add H2O to balance the O atoms and H+ to balance the H atoms. In a basic solution, we would use OH- and H2O to balance the O and H.
Balance the iodine atoms:
2 I- → I2
The Mn in the permanganate reaction is already balanced, so let's balance the oxygen:
MnO4- → Mn2+ + 4 H2O
Add H+ to balance the 4 waters molecules:
MnO4- + 8 H+ → Mn2+ + 4 H2O
The two half-reactions are now balanced for atoms:
MnO4- + 8 H+ → Mn2+ + 4 H2O
Next, balance the charges in each half-reaction so that the reduction half-reaction consumes the same number of electrons as the oxidation half-reaction supplies. This is accomplished by adding electrons to the reactions:
2 I- → I2 + 2e-
5 e- + 8 H+ + MnO4- → Mn2+ + 4 H2O
Now multiple the oxidations numbers so that the two half-reactions will have the same number of electrons and can cancel each other out:
5(2I- → I2 +2e-)
2(5e- + 8H+ + MnO4- → Mn2+ + 4H2O)
Now add the two half-reactions:
10 I- → 5 I2 + 10 e-
16 H+ + 2 MnO4- + 10 e- → 2 Mn2+ + 8 H2O
This yields the following final equation:
10 I- + 10 e- + 16 H+ + 2 MnO4- → 5 I2 + 2 Mn2+ + 10 e- + 8 H2O
Get the overall equation by canceling out the electrons and H2O, H+, and OH- that may appear on both sides of the equation:
10 I- + 16 H+ + 2 MnO4- → 5 I2 + 2 Mn2+ + 8 H2O
Check your numbers to make certain that the mass and charge are balanced. In this example, the atoms are now stoichiometrically balanced with a +4 net charge on each side of the reaction.

C) VOLTA CELL
A voltaic pile is a set of individual Galvanic cells placed in series. The voltaic pile, invented by Alessandro Volta in 1800, was the first electric battery. Building on Galvani's 1780s discovery of how a circuit of two metals and a frog's leg can cause the frog's leg to respond, Volta demonstrated in 1791 that when two metals and brine-soaked cloth or cardboard are arranged in a circuit they produce an electric current. In 1800 Volta literally piled up several pairs of alternating copper (or silver) and zinc discs (electrodes) separated by cloth or cardboard soaked in brine (electrolyte) to increase the electrolyte conductivity.[1] When the top and bottom contacts were connected by a wire, an electric current flowed through the voltaic pile and the connecting wire
Volta deret
“ Li K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb Bi Cu Hg Ag Pt Au ”
Spontaneous redox reaction,
To determine if a redox reaction is spontaneous, you should compute the voltage of the reaction.
• If the voltage is positive, the reaction is spontaneous
• If the voltage is negative, the reaction is not spontaneous
Example: Is the reaction below spontaneous?
Cu(s) + 2Fe+3(aq) -> Cu+2(aq) + 2Fe+2(aq)
Solution: To determine spontaneity, we need to determine the voltage. Break the reaction up into two half reactions and check the voltages of each
• Cu(s) -> Cu+2 + 2e- Eox = -0.339 V
• Fe+3(aq) +e- -> Fe+2(aq) Eox = +0.769 V
The cell voltage is thus E = Ered + Eox = -0.339 + 0.769 = +0.43 V. Since E is positive, the reaction is spontaneous.

D) ELECTROLYSIS CELL
An external electric current hooked up to an electrochemical cell will make the electrons go backward. An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes. When driven by an external voltage applied to the electrodes, the electrolyte provides ions that flow to and from the electrodes, where charge-transferring, or faradaic, or redox, reactions can take place. Only for an external electrical potential (i.e. voltage) of the correct polarity and large enough magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. The electrical energy provided undoes the effect of spontaneous chemical reactions.
Faraday's laws of electrolysis
Several versions of the laws can be found in textbooks and the scientific literature. The most-common statements resemble the following:
Faraday's 1st Law of Electrolysis - The mass of a substance altered at an electrode during electrolysis is directly proportional to the quantity of electricity transferred at that electrode. Quantity of electricity refers to the quantity of electrical charge, typically measured in coulomb.
Faraday's 2nd Law of Electrolysis - For a given quantity of electricity (electric charge), the mass of an elemental material altered at an electrode is directly proportional to the element's equivalent weight. The equivalent weight of a substance is its molar mass divided by an integer that depends on the reaction undergone by the material.
Faraday's laws can be summarized by

where
m is the mass of the substance altered at an electrode
Q is the total electric charge passed through the substance
F = 96,485 C mol-1 is the Faraday constant
M is the molar mass of the substance
z is the valency number of ions of the substance (electrons transferred per ion)

Note that M / z is the same as the equivalent weight of the substance altered.
For Faraday's first law, M, F, and z are constants, so that the larger the value of Q the larger m will be.
For Faraday's second law, Q, F, and z are constants, so that the larger the value of M / z (equivalent weight) the larger m will be.
In the simple case of constant-current electrolysis, Q = It leading to

and then to

where
n is the amount of substance ("number of moles") altered: n = m / M
t is the total time the constant current was applied.
In the more-complicated case of a variable electrical current, the total charge Q is the electric current I(τ) integrated over time τ:

Here t is the total electrolysis time. Please note that tau is used as the current I is a function of tau.[2]
E) Corrosion
Corrosion is the disintegration of an engineered material into its constituent atoms due to chemical reactions with its surroundings. In the most common use of the word, this means electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Formation of an oxide of iron due to oxidation of the iron atoms in solid solution is a well-known example of electrochemical corrosion, commonly known as rusting. This type of damage typically produces oxide(s) and/or salt(s) of the original metal. Corrosion can also refer to other materials than metals, such as ceramics or polymers. Although in this context, the term degradation is more common.
The effects of corrosion are often (ok, essentially always) undesirable. Rust is generally formed at the surface of the metal, and doesn't adhere to the surface. When rust flakes off, the underlying metal is exposed to further corrosion (a process called pitting), and eventually the structural integrity of the metal disintegrates.
The costs associated with corrosion damage are enormous. For example, in the US they add up to approximately 4.2 percent of the GNP (well over 400 billion dollars per year). These costs can be attributed to failure or deterioration of equipment, and the associated downtime, replacement and maintenance.
The overall reaction for the corrosion of iron is as follows:
H2O
4 Fe (s) + 3 O2 (aq) -> 2 Fe2O3 (s)
The preceding reaction balance does not show what is actually happening to the metal (i.e. the reaction mechanism). The reaction does not occur by a direct reaction between oxygen and iron atoms such as in a combustion reaction, but by electron transfer from the metal to water. This type of reaction is called a redox reaction. In this reaction, one of the reactants is reduced (gains an electron), and another one is oxidized. It's the same process that drives those digital clocks that run on a potato with an iron and copper electrode stuck in them.
The half-reactions of the redox reaction give a more accurate description of the reaction mechanism. The iron is first oxidized to Fe2+ (ferrous ion):
Fe -> Fe2+ + 2 e-
The ferrous ion is further oxidized to Fe3+, ferric ion:
Fe2+ -> Fe3+ + e-
The electrons that were produced by the oxidation steps are used in the reduction reaction of water:
O2 + 2 H2O + 4 e- -> 4 OH-
Combining the above reactions, and balancing them for the number of electrons being transferred, we get:
4 Fe + 3 O2 + 6 H2O -> 4 Fe3+ + 12 OH-
however, these free ferric and hydroxide ions are not actual intermediates for the reaction, but directly form a more stable complex, under the liberation of water:
-> 2 Fe2O3 + 6 H2O