Is it necessary that if an oxidation reaction happens, then a reduction reaction has also happened?
Answers
Explanation:
Oxidation-reduction reactions (redox reactions) are reactions in which electrons are lost by an atom or ion in one reactant and gained by an atom or ion in another reactant. Although electrons are gained and lost in these reactions, the balanced equation for a redox reaction does not show the electrons that are being transferred. In order to tell whether a redox reaction has occurred or not, we need a way to keep track of electrons. The best way to do so is by assigning oxidation numbers to the atoms or ions involved in a chemical reaction.
An oxidation-reduction (redox) reaction is a type of chemical reaction that involves a transfer of electrons between two species. An oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion changes by gaining or losing an electron. Redox reactions are common and vital to some of the basic functions of life, including photosynthesis, respiration, combustion, and corrosion or rusting.
Rules for Assigning Oxidation States
The oxidation state (OS) of an element corresponds to the number of electrons, e-, that an atom loses, gains, or appears to use when joining with other atoms in compounds. In determining the oxidation state of an atom, there are seven guidelines to follow:
The oxidation state of an individual atom is 0.
The total oxidation state of all atoms in: a neutral species is 0 and in an ion is equal to the ion charge.
Group 1 metals have an oxidation state of +1 and Group 2 an oxidation state of +2
The oxidation state of fluorine is -1 in compounds
Hydrogen generally has an oxidation state of +1 in compounds
Oxygen generally has an oxidation state of -2 in compounds
In binary metal compounds, Group 17 elements have an oxidation state of -1, Group 16 elements of -2, and Group 15 elements of -3.
The sum of the oxidation states is equal to zero for neutral compounds and equal to the charge for polyatomic ion species.
Example 1 : Assigning Oxidation States
Determine the Oxidation States of each element in the following reactions:
Fe(s)+O2(g)→Fe2O3(g)
Fe2+
Ag(s)+H2S→Ag2S(g)+H2(g)
SOLUTIONS
Fe and O2 are free elements; therefore, they each have an oxidation state of 0 according to Rule #1. The product has a total oxidation state equal to 0, and following Rule #6, O has an oxidation state of -2, which means Fe has an oxidation state of +3.
The oxidation state of Fe corresponds to its charge; therefore, the oxidation state is +2.
Ag has an oxidation state of 0, H has an oxidation state of +1 according to Rule #5, S has an oxidation state of -2 according to Rule #7, and hence Ag in Ag2S has an oxidation state of +1.
Example 2 : Assigning Oxidation States
Determine the Oxidation State of the bold element in each of the following:
Na3PO3
H2PO4-
SOLUTIONS
The oxidation numbers of Na and O are +1 and -2. Because sodium phosphite is neutral, the sum of the oxidation numbers must be zero. Letting x be the oxidation number of phosphorus, 0= 3(+1) + x + 3(-2). x=oxidation number of P= +3.
Hydrogen and oxygen have oxidation numbers of +1 and -2. The ion has a charge of -1, so the sum of the oxidation numbers must be -1. Letting y be the oxidation number of phosphorus, -1= y + 2(+1) +4(-2), y= oxidation number of P= +5.
Example 3 : Identifying Reduced and Oxidized Elements
The two species that exchange electrons in a redox reaction are given special names. The ion or molecule that accepts electrons is called the oxidizing agent; by accepting electrons it causes the oxidation of another species. Conversely, the species that donates electrons is called the reducing agent; when the reaction occurs, it reduces the other species. In other words, what is oxidized is the reducing agent and what is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound, as in disproportionation reactions).
A good example of a redox reaction is the thermite reaction, in which iron atoms in ferric oxide lose (or give up) O atoms to Al atoms, producing Al2O3.
Fe2O3(s)+2Al(s)→Al2O3(s)+2Fe(l)(1)
Example 4 : Identifying Oxidized Elements
Using the equations from the previous examples, determine what is oxidized in the following reaction.
Zn+2H+→Zn2++H2(2)
SOLUTION
The oxidation state of H changes from +1 to 0, and the oxidation state of Zn