Here are even more websites to refer to should you want to go into deeper details regarding redox equations.
http://www.science.uwaterloo.ca/~cchieh/cact/c123/balance.html A rather detailed website on how to write redox equations well, there are also quizzes at the bottom of the website of this page..
http://www.chemreview.net/HalfReactionBalancing.pdf Notes together with quizzes at the same time. Highly useful.
Here are some quizzes definitely worth trying for practice in redox.
http://www.mp-docker.demon.co.uk/as_a2/topics/redox/quiz_1.html
Friday, August 17, 2012
Oxidation and reduction:Changes in oxidation state
Oxidation is an increase in oxidation state(number) during a reaction.
Reduction is a decrease in oxidation state (number during a reaction.
The concept of oxidation state is used in several reactions are given below.
1) Reaction of metals with dilute acids
Consider the reaction of zinc with dilute sulfuric acid:
Zn(s) +H2SO4 --->ZnSO4(aq) +H2(g)
0 +1 +2 0
The oxidation of zinc increase from 0 to 2 in ZnSO4 and is oxidised. The oxidation state of hydrogen decreases from +1 in H2SO4 to 0 and so is reduced, NOTE THAT THE SULFATE ION DOES NOT CHANGE. THEREFORE THE OXIDATION NUMBER OF THE S(+6) AND O2(-2) in the ion DO NOT CHANGE.
2)Halide displacement reactions
Chlorides react with KI solution to give an ionic equation...
Cl2(aq) + 2I(ag) --> 2Cl + I2
Chlorine is reduced (oxidation number decreases) whereas the iodide ions are oxidised. (THE OPPOSITE)
Reduction is a decrease in oxidation state (number during a reaction.
The concept of oxidation state is used in several reactions are given below.
1) Reaction of metals with dilute acids
Consider the reaction of zinc with dilute sulfuric acid:
Zn(s) +H2SO4 --->ZnSO4(aq) +H2(g)
0 +1 +2 0
The oxidation of zinc increase from 0 to 2 in ZnSO4 and is oxidised. The oxidation state of hydrogen decreases from +1 in H2SO4 to 0 and so is reduced, NOTE THAT THE SULFATE ION DOES NOT CHANGE. THEREFORE THE OXIDATION NUMBER OF THE S(+6) AND O2(-2) in the ion DO NOT CHANGE.
2)Halide displacement reactions
Chlorides react with KI solution to give an ionic equation...
Cl2(aq) + 2I(ag) --> 2Cl + I2
Chlorine is reduced (oxidation number decreases) whereas the iodide ions are oxidised. (THE OPPOSITE)
Oxidation States
Oxidation states, also known by many others as valency or oxidation number, is the charge of an atom IF IT WAS TO EXIST AS AN ION. Remember that oxidation states follow a few rule...
1)Elements have an oxidation state of zero.
2)The oxidation state of a simple ion is the charge of the ion.
3)The oxidation state of some elements in their compoounds is fixed.
4) The sum of the oxidation states of the elements in an ion is equal ti tge charge if the ion
Transition metals have variable oxidation states., most of them having names including Roman numeral(I,II,III,etc.)
For example copper(II) sulfate, the copper atom has an oxidation of 2, having a charge of 2+.
In potassium manganate (VII) the manganese atom has an oxidation state of 7+.
1)Elements have an oxidation state of zero.
2)The oxidation state of a simple ion is the charge of the ion.
3)The oxidation state of some elements in their compoounds is fixed.
4) The sum of the oxidation states of the elements in an ion is equal ti tge charge if the ion
Transition metals have variable oxidation states., most of them having names including Roman numeral(I,II,III,etc.)
For example copper(II) sulfate, the copper atom has an oxidation of 2, having a charge of 2+.
In potassium manganate (VII) the manganese atom has an oxidation state of 7+.
Redox reactions involving transfer of electrons
Redox reactions involving transfer of electrons can be splitted into 2 1/2ves, with one half showing oxidation and the other showing reduction!
Two kinds of redox reactions that can be explained using a transfer of electrons are: reactions of metals with dilute acids, and displacement reactions. First there is reaction of metals with dilute acid. The ionic equation of magnesium and dilute hydrochloric acid is:
Mg(s) +2H+-->Mg2+(aq) + H2(g)
In the reaction, electrons are transferred from the magnesium atom to the hydrogen ions, the Mg atom losing electrons to form Magnesium ion and hydrogen having more electrons to form a molecule..
The 2nd kind are displacement reactions. Chlorine displaces iodine to form KI solution. Electrons are transferred from iodide ions to the chlorine. The iodide ions are oxidised through losing electrons and the chlorine gains electrons and is reduced.
Two kinds of redox reactions that can be explained using a transfer of electrons are: reactions of metals with dilute acids, and displacement reactions. First there is reaction of metals with dilute acid. The ionic equation of magnesium and dilute hydrochloric acid is:
Mg(s) +2H+-->Mg2+(aq) + H2(g)
In the reaction, electrons are transferred from the magnesium atom to the hydrogen ions, the Mg atom losing electrons to form Magnesium ion and hydrogen having more electrons to form a molecule..
The 2nd kind are displacement reactions. Chlorine displaces iodine to form KI solution. Electrons are transferred from iodide ions to the chlorine. The iodide ions are oxidised through losing electrons and the chlorine gains electrons and is reduced.
Oxidising and Reducing Agents
Remember in the previous reaction there was a reaction between zinc and copper (II) oxide? Copper (II) oxide does give oxygen to zinc to oxidise it. The copper (II) oxide is said to be an oxidising agent as it results in the oxidation process. Therefore it is evident to refer the zinc as the reacting agent as it reduces the copper (II) oxide by removing oxygen.
In other words, a substance that can cause oxidation is an oxidising reaction(self-explanatory from the name, I guess).
The opposite is definitely true too, which means that a substance that can cause reduction is a reducing agent(also self-explanatory from the name).
Oxidation and reduction is also described as the loss and gain of hydrogen respectively of a substance.
An example of this would be hydrogen sulfide plus chlorine....
H2S(g) + Cl2 (g) --->2HCl(g) + S(s)
This reaction has shown the hydrogen sulfide to lose hydrogen, hence it is oxidised. The chlorine gas has gained the hydrogen and resulted in the reduction of chlorine. The chlorine has caused the oxidation of the hydrogen sulfide, hence being the oxidising agent..
However, this was not the only few definitions for oxidation and reduction!!
Oxidation is regarded as the loss of electrons of a substance, and reduction is seen as the direct opposite of it--the gain of electrons by a substance.This was when chemists were able to use the knowledge acquired through atomic structure to include reactions with the exclusion of O2 and H2.
A simple example of this kind of equation is the common reaction of sodium metal with chlorine gas---that's right our sodium chloride!:)
2Na(s) + Cl2(g) -->2NaCl
Electrons are transferred from the sodium atoms to the chlorine atoms to form sodium ions and chloride ions via a vice versa process. The chlorine gains electrons, so it is reduced. The Na loses electrons, hence being oxidised. The chlorine gains electrons, so it is reduced.
In other words, a substance that can cause oxidation is an oxidising reaction(self-explanatory from the name, I guess).
The opposite is definitely true too, which means that a substance that can cause reduction is a reducing agent(also self-explanatory from the name).
Oxidation and reduction is also described as the loss and gain of hydrogen respectively of a substance.
An example of this would be hydrogen sulfide plus chlorine....
H2S(g) + Cl2 (g) --->2HCl(g) + S(s)
This reaction has shown the hydrogen sulfide to lose hydrogen, hence it is oxidised. The chlorine gas has gained the hydrogen and resulted in the reduction of chlorine. The chlorine has caused the oxidation of the hydrogen sulfide, hence being the oxidising agent..
However, this was not the only few definitions for oxidation and reduction!!
Oxidation is regarded as the loss of electrons of a substance, and reduction is seen as the direct opposite of it--the gain of electrons by a substance.This was when chemists were able to use the knowledge acquired through atomic structure to include reactions with the exclusion of O2 and H2.
A simple example of this kind of equation is the common reaction of sodium metal with chlorine gas---that's right our sodium chloride!:)
2Na(s) + Cl2(g) -->2NaCl
Electrons are transferred from the sodium atoms to the chlorine atoms to form sodium ions and chloride ions via a vice versa process. The chlorine gains electrons, so it is reduced. The Na loses electrons, hence being oxidised. The chlorine gains electrons, so it is reduced.
Oxidation and Reduction: Introduction to redox processes
Oxidation is the process of a substance gaining oxygen in a chemical reaction. For example, magnesium burns in oxygen, with the equation for this reaction being:
2Mg(s) + O2-->2MgO(s)
This shows that magnesium has been oxidised as it gains oxygen to become magnesium oxide.
There is also a reverse process for oxidation: yes, the loss of oxygen!! Such a process is called reduction. A substance that loses oxygen in a reaction is said to be reduced. Another such example would be zinc reacting with copper (II) oxide. The equation for the reaction is :
Zn(s) + CuO (s) --> ZnO(s) + Cu(s)
In the reaction, the copper (II) oxide reacts and loses oxygen, therefore being called "reduced".
Such two processes usually takes place together! In the reaction for zinc reacting with copper (II) oxide,
the zinc gets oxidised whilte the copper (II) oxide loses oxygen. The name for such reactions are called redox reactions.
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