Redox Reactions or Oxidation Reduction Reactions

Redox Reactions  / Oxidation and reduction reactions


These are reactions where electrons are transferred from one species (atom, molecule or ion) to another. We can write 'half' equations to show only what happens to the species losing electrons or a different 'half' equation to show the species gaining electrons.
The whole equation is put together by making sure that the numbers of electrons are balanced in each half equation and adding them together (when the electrons will cancel out)
Oxidation
This is the name given to removal of electrons from a species - the reagent causing the loss of electrons is called the oxidising agent

Example: Mg(s) Mg2+ + 2e In this (half) equation the magnesium atom loses electrons and becomes an ion.

Reduction
This is the gain of electrons - the species donating the electrons is called the reducing agent

Example Fe3+ + 3e Fe(s) In this (half) equation the iron(III) ion gains three electrons to become an atom.

Redox reactions
Obviously the electrons leave one species and go to another. Consequently reduction has to be accompanied by oxidation and vice versa. For this reason reactions involving transfer of electrons are called reduction and oxidation or redox for short

Example 3Mg(s) + 2Fe3+ 2Fe(s)+ 3Mg2+ The electrons from the magnesium are transferred to the iron(III) ions

Summary

Loss of electrons = Oxidation Gain of electrons = Reduction	Mnemonic (memory aid) OIL-RIG
			Oxidation Is Loss Reduction Is Gain

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10.1.2: Calculate the oxidation number of an element in a compound. Oxidation numbers should be shown by a sign (+ or -) and a number, eg +7 for Mn in KMnO4.

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Oxidation number
This is the apparent valency of an atom within a compound. It is usually considered as if the element were bonded ionically to allow the apparent number of electrons gained or lost to be assessed.
The sum of all the oxidation numbers in a compound must add up to 0. By convention, the oxidation number is written as a Roman numeral in the name, eg. iron II sulphate, sulphur VI oxide.
The oxidation number of an uncombined element is always zero (0)
Calculating the oxidation number
There are some elements that virtually always have the same oxidation number and these can be used to calculate the oxidation numbers of the atoms in question.
Hydrogen, for example always has an oxidation number of -1 when bonded to a metal (more electropositive element) and +1 when bonded to a more electronegative element (non-metal). Oxygen is always -2 (except when in the form of the peroxide ion when it has an O-O bond giving it an oxidation number of -1). Group 1 and 2 metals usually have an oxidation number of 1+ and 2+ respectively.

Example - Calculate the oxidation number of sulphur in sulphuric acid H2SO4 Hydrogen = +1 oxidation number Oxygen = -2 oxidation number Therefore: (2 x H) + S + (4 x O) = 0 2 + S -8 = 0 S = 6

Example - Calculate the oxidation number of nitrogen in calcium nitrate Ca(NO3)2 Calcium is in group 2 = +2 oxidation number Oxygen = -2 oxidation number Therefore: (+2) + [(2 x N) + (6 x -2)] = 0 +2 + 2N -12 = 0 2N = 10 N = +5

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10.1.3: State and explain the relationship between oxidation numbers and the names of compounds. Oxidation numbers in names of compounds are represented by Roman numerals, eg iron(II) oxide, iron(III) oxide.

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Names of compounds
Where there is any doubt about the oxidation state of an element within a compound it is stated using Roman numerals immediately after the ambiguous element. For example Iron compounds may be iron in the oxidation state +2 or +3 - it must therefore be stated as iron II or iron III in the compound name.
In the examples above the full systematic name for sulphuric acid is sulphuric(VI) acid and calcium nitrate is calcium nitrate(V)

Example - Name the following compound - FeSO4 Oxidation state of the oxygen = -2 Oxidation state of the sulphur = +6 Therefore oxidation state of the iron = - (+6 - 8) = +2 The name of the compound FeSO4 is iron(II) sulphate

Example - Name the following compound - TiCl4 Oxidation state of the chloride = -1 Therefore oxidation state of the titanium = - (- 4) = +4 The name of the compound TiCl4 is titanium(IV) chloride

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10.1.4: Identify whether an element is oxidised or reduced in simple redox reactions, using oxidation numbers. Appropriate reactions to illustrate this can be found in topics 3 and 11. Possible examples include: iron(II) and (III), manganese(II) and (VII), chromium(III) and (VI), copper(I) and (II), oxides of sulphur and oxyacids, halogens and halide ions.

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Oxidation and reduction
As stated above, for the purposes of oxidation and reduction the oxidation number can be thought of as the apparent ionic charge of an atom within a compound. For example, in sulphuric acid the sulphur is in the VI (6+) oxidation state. For the purposes of redox we can consider that it has an ionic charge of +6 (even though it is clearly covalently bonded). This makes it easier to follow any transfer of electrons.
If the sulphur changes to an oxidation state of IV during a chemical reaction then it has gone from an apparent ionic charge of +6 to a charge of +4, i.e. it has gained two electrons (negative charges). It has therefore been reduced (gain of electrons) in the process.

Examples 2FeCl2 + Cl2 2FeCl3 The iron changes from 2+ to 3+ and is therefore oxidised (removal of electrons) The chlorine gains an electron to go from 0 to -1 and is therefore reduced (addition of electrons) Zn + CuSO4 Cu + ZnSO4 The zinc changes from oxidation state 0 to +2 (removal of electrons) it is oxidised  The copper changes from 2+ to 0 and is oxidised and is therefore reduced (addition of electrons) Cr2O72- + 3SO2 + 2H+ 2Cr3+ + 3SO42- + H2O The chromium changes from +6 to +3 and is therefore reduced (gain of electrons) The sulphur changes from +4 to +6 and therefore loses electrons = oxidation (loss of electrons) 2KI + Br2 2KBr + I2 The iodide ions (oxidation number = -1) change to iodine (oxidation number = 0) : oxidation Bromine (element, oxidation number = 0) changes to bromide ions (oxidation number = -1) : reduction 5Fe2+ + MnO4- + 8H+ 5Fe3+ + Mn2+ + 4H2O The iron changes from 2+ to 3+ and is oxidised (removal of electrons) The manganese atom changes from +7 to +2 and is therefore reduced (addition of electrons)

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10.1.5: Define the terms oxidising agent and reducing agent.

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Oxidising agents
These are the chemicals that cause the oxidation in a redox reaction. We call the reacting compounds in a reaction the reagents (short form of the words reacting agents).
We consider that the removal of electrons from a species is oxidation and these electrons have to be taken away by another compound or species. This species that attracts the electrons is said to be the oxidising agent i.e. the reagent that causes the oxidation.
Reducing agents
Similarly the reagent that causes reduction in a redox reaction is said to be the reducing agent.
The oxidising agent takes the electron and is itself reduced, the reducing agent loses the electrons and is itself oxidised.

2KI	+ Br2		2KBr	+ I2
Iodide ions get oxidised	Bromine gets reduced
Iodide - reducing agent	Bromine - oxidising agent
Cr2O72-	+ 3SO2 + 2H+		2Cr3+	+ 3SO42- + H2O
Chromium VI gets reduced	Sulphur IV gets oxidised
Chromium VI oxidising agent	Sulphur IV reducing agent

Types of Chemical Reactions

Most reactions in chemistry take place with water as the solvent. Solutions where water is the solvent are called aqueous solutions. Water is known to be a polar molecule because the unequal distribution of charge caused by the electronegative oxygen. This polarity gives the water the ability to dissolve ions and other polar substance.
A useful method for characterizing an aqueous solution is its electrical conductivity. If a solution conducts electricity well, it is considered a strong electrolyte. If it only conducts slightly, it is considered a weak electrolyte; if it doesn’t conduct, it’s a nonelectrolyte.
Strong electrolytes are substances that are completely ionized in water. Such examples are strong acids, strong bases, and soluble salts. Acids ionize into H⁺ and A^-; bases ionize into OH^- and X⁺.
A weak electrolyte is a substance that only slightly ionizes when added to water. These substances include weak acids, weak bases, and slight soluble salts. Weak acids are those that only dissociate slightly into H⁺ and A^-. Similarly, weak bases are those that dissociate only slightly into OH^- and X⁺.
Nonelectrolytes are substances that dissolve in water, but that don’t break up into ions. These substances are mostly polar molecules; the reason they don’t conduct electricity is that no ions are formed
Solution Reactions
One type of reaction is called a precipitation reaction. This occurs when two solutions are mixed resulting, and a solid or precipitate forms. The precipitate contains ions that when combined are insoluble with water. However, these individual ions do dissolve in water. So, the result of mixing these ions is an insoluble solid. One example of a precipitation reaction is when a solution containing Ca²⁺ and a solution containing SO₄^2-; the result of mixing these the formation of solid CaSO₄.
Another type of reaction is an acid-base reaction. An acid-base reaction is one when the net reaction is the combination of a proton and hydroxide to form water. The acid base reaction is also called a neutralization reaction. When performing calculations fo an acid-base reaction follow these simple steps:

1. List all the species before any reaction occurs and decide what reaction will occur.
2. Write the balance net-ionic equation
3. Calculate the moles of reactants
4. Determine the limiting reactant where appropriate
5. Calculate the moles of the required reactant or product.
6. Convert to grams or volume, as required.

The final type of aqueous reaction is an oxidation-reduction reaction. An oxidation-reaction is characterized by the transfer of one or more electrons. These reactions are often used for energy production; in fact, these reactions are often used in the human body to provide energy.
The concept of oxidation states provides a means for keeping track of the movement of electrons in a redox reaction. The oxidation states can be assigned according to these rules:

1. The oxidation state of an atom in an element is 0.
2. The oxidation state for a single atom ion is the charge of the ion.
3. Oxygen is assigned the oxidation state of –2 in covalent compounds, except in peroxides where oxygen is assigned a –1 state.
4. In covalent compound hydrogen is assigned a +1 state.
5. In a covalent compound, fluorine is always –1 state.
6. The sum of the oxidation states must equal to overall charge of the molecule.

When balancing oxidation it is convenient to divide the reaction into two half-reactions; one reaction involves the oxidation, the other the reduction. Then, follow these steps if it’s in acid:

1. write the half reactions
2. for each half reaction balance the element except H and O, balance the O with water, balance the H with H+, and balance the charge with electrons.
3. If necessary, multiple the reaction by an integer to equalize the number of electrons.
4. Add the half reaction

If the reaction occurs in base repeat the above steps, but after balancing the hydrogen with H⁺, add hydroxides to cancel out the H⁺’s. Then, continue the above steps as before.