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Redox processes include the reduction of one species and the oxidation of another. As a result, the oxidation state of the species in question must shift. These reactions are crucial for a variety of applications, including energy storage devices (batteries), photographic processing, and energy production and use in living systems, such as humans.
Reduction is a chemical reaction in which an atom receives an electron and hence lowers in mass (or reduces its oxidation number). The species’ positive feature is essentially diminished.
Oxidation is the loss of an electron by an atom, which increases its oxidation number. In other words, the species’ good characteristics are enhanced.
Oxidation
The name “oxidation” was coined because the first redox reactions to be carefully studied occurred in oxygen, with oxygen being reduced and the other species being oxidised, thus the term “oxidation reaction.” However, it was eventually discovered that this case (oxygen oxidation processes) was only one conceivable possibility. Consider the redox reaction illustrated below.
Example 1: Which species is oxidised in the following redox reaction? Which of them is being cut?
The oxidation of Al(s).
The amount of Ag+ (aq) is decreasing.
To memorise the definitions of oxidation and reduction, use the following mnemonic: The lion Leo has gone berserk.
Leo: oxidation = loss of electron(s);
Gain of electron(s) = Ger : reduction.
Oxidation State: The state of a species having a particular oxidation number. An element in the associated oxidation state has a particular oxidation number.
The Process of Assigning Oxidation Numbers
The rules for assigning oxidation numbers are stated in ascending order below.
The following are pure elements in their natural, standard state: ox. # = 0.
# = ionic charge for monatomic ions.
In compounds, F is always F (-I).
ox. # = I. Alkali metals (those in the first column of the periodic table):
Alkaline-earth metals (those in the periodic table’s second column): ox # = II.
H is almost always hydrogen (I). Metal hydrides are an exception (MHx).
In compounds, oxygen is generally always represented by the symbol O (-II). O-O and O-F are exceptions.
The total charge of a species is equal to the sum of all oxidation numbers in that species.
Examples 2:
Use the following equation to assign oxidation numbers to molecules:
Take HBrO2, for example. From Rule 7, we know that O has an oxidation number of -2 and that hydrogen is H (I). On HBrO2, the total charge is zero. When we solve for the oxidation number of Br using the equation above, we get the following result.
Guidelines for Redox Balance Equations:
Determine each species’ oxidation states.
Fill in the blanks for each half reaction:
Atoms that change oxidation state must be balanced.
Calculate the total number of electrons acquired or lost.
Use H+ (in an acidic solution) or OH- to balance charges (in basic solution).
H2O is used to balance the remaining atoms (H’s and O’s).
Using the proper factor, balance the number of electrons transmitted for each half reaction so that the electrons cancel.
Combine the two half-reactions and simplify if needed.
Example 3: Balance the redox process below.
Determine the oxidation states of the species involved in the first step.
Because the charges do not yet equal, the equation is not balanced. Each half-reaction can be used to balance the charges. Because the Cl- ions are spectator ions and do not participate in the actual redox process, they drop out.
Step 2: Make a list of half reactions.
Step 2a. Balance the atoms whose oxidation states change.
step 2b Determine the number of electrons acquired or lost in .
Three electrons are lost when aluminium goes from 0 to III. The situation is a little different with hydrogen. Hydrogen is decreasing from one to zero. This means that one electron is required for each H+ ion that interacts. Two electrons are required since the reaction involves two H+ ions.
Because the equations are balanced, steps 2c and 2d are not required.
Step 3: Ensure that the quantity of electrons transported is balanced.
Because the electrons exchanged have a common factor of 6, the above multiplication is done.
Step 4: The charges and atoms have now been balanced. To make sure, sum up all the charges and atoms on each side. Both the quantity of atoms and their charges must be balanced. Note that this is not a neutral response. Keep in mind that the solution is neutralised by the spectator ions, Cl-.
Conclusion
Redox processes include the reduction of one species and the oxidation of another. As a result, the oxidation state of the species in question must shift. These reactions are crucial for a variety of applications, including energy storage devices (batteries), photographic processing, and energy production and use in living systems, such as humans.
Reduction is a chemical reaction in which an atom receives an electron and hence lowers in mass (or reduces its oxidation number). The species’ positive feature is essentially diminished.
Oxidation is the loss of an electron by an atom, which increases its oxidation number. In other words, the species’ good characteristics are enhanced.
