Which substance can act as both a reducing agent and an oxidising agent?

Answer 1

Any substance containing an atom that it is at an intermediate oxidation level compared to all its possible oxidation states.

For example in hydrogen peroxide (H2O2) oxygen is at the -1 state. However oxygen can be also at 0 or -2.
Thus under proper conditions H2O2 acts a reducing agent and thus it oxidises to O2 (0 oxidation number) or as an oxidising agent and thus it gets reduced to O with an oxidation number -2.

Answer 2

One example is carbon monoxide, CO.
It can be oxidised to form CO2 and reduced to form C.
In the former case, CO would be the reducing agent while in the latter, it is the oxidising agent.

Another example is Fe 2+.
It oxidises to form Fe 3+ and reduces to form Fe.

Similarly, any ion or compound is capable of gaining oxygen, losing hydrogen, increasing its oxidation state and losing electrons will oxidise. The reverse would result in a reduction.

Answer 3

Well that depends. Theoretically any substance chemically can be both an oxidizing and reducing agent depending on what it's reacting with. The one exception are usually the noble gases which require a lot of electrons and aggressive reaction to cause them to change valence. Generally you can look at Nitrogen being an electron donator and acceptor (oxidizing and reducing) element for several applications (nitrogen dioxide, nitrous oxide, etc.)

Answer 4

Hydrogen gas is a reducing agent when it reacts with non-metals and an oxidising agent when it reacts with metals.

Oxidising agent:
2Li(s) + H2(g) -->2LiH(s) hydrogen acts as an oxidizing agent because it accepts an electron donation from lithium, which causes Li to be oxidized.

Reducing agent:
H2(g) + F2(g) --> 2HF(g) hydrogen acts as a reducing agent because it donates its electrons to fluorine, which allows fluorine to be reduced.

Answer 5

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Read redox reactions.A good example is the reaction between hydrogen and fluorine in which hydrogen is being oxidized and fluorine is being reduced: H2 + F2 → 2 HF We can write this overall reaction as two half-reactions: the oxidation reaction: H2 → 2 H+ + 2 e− and the reduction reaction: F2 + 2 e− → 2 F− Analyzing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above). Elements, even in molecular form, always have an oxidation state of zero. In the first half-reaction, hydrogen is oxidized from an oxidation state of zero to an oxidation state of +1. In the second half-reaction, fluorine is reduced from an oxidation state of zero to an oxidation state of −1. When adding the reactions together the electrons are canceled: H2 → 2 H+ + 2 e− F2 + 2 e− → 2 F− H2 + F2 → 2 H+ + 2 F− And the ions combine to form hydrogen fluoride: 2 H+ + 2 F− → 2 HF The overall reaction is: H2 + F2 → 2 HF Displacement reactions Redox occurs in single displacement reactions or substitution reactions. The redox component of these types of reactions is the change of oxidation state (charge) on certain atoms, not the actual exchange of atoms in the compounds. For example, in the reaction between iron and copper(II) sulfate solution: Fe + CuSO4 → FeSO4 + Cu The ionic equation for this reaction is: Fe + Cu2+ → Fe2+ + Cu As two half-equations, it is seen that the iron is oxidized: Fe → Fe2+ + 2 e− And the copper is reduced: Cu2+ + 2 e− → Cu Other examples The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid: Fe2+ → Fe3+ + e− H2O2 + 2 e− → 2 OH− Overall equation: 2 Fe2+ + H2O2 + 2 H+ → 2 Fe3+ + 2 H2O The reduction of nitrate to nitrogen in the presence of an acid (denitrification): 2 NO3− + 10 e− + 12 H+ → N2 + 6 H2O Iron rusting in pyrite cubes Oxidation of elemental iron to iron(III) oxide by oxygen (commonly known as rusting, which is similar to tarnishing): 4 Fe + 3 O2 → 2 Fe2O3 The combustion of hydrocarbons, such as in an internal combustion engine, which produces water, carbon dioxide, some partially oxidized forms such as carbon monoxide, and heat energy. Complete oxidation of materials containing carbon produces carbon dioxide. In organic chemistry, the stepwise oxidation of a hydrocarbon by oxygen produces water and, successively, an alcohol, an aldehyde or a ketone, a carboxylic acid, and then a peroxide.

Answer 6

Hydrogen peroxide can act as a reducing agent and as an oxidising agent:

5e^- + 8^+ + MnO4^- ---> Mn^2+ + 4H2O
H2O2 ---> O2 +2H^+ 2e^-

Fe^2+ ---> Fe^3+ + e^-
2e^- + 2H^+ + H2O2 ---> 2H2O

Answer 7

Most substances can act as oxidant and reductant, under certain conditions.

Answer 8

concentrated sulphuric acid

Answer 9

BrCl is not impossible. Look it up

Answer 10

Sulphur is a non-metallic element which exists in different crystal forms known as allotropes. The main allotropes of sulphur are:



rhombic sulphur is the stable form at room temperature. The crystal shape is shown above. When it is heated slowly above 95.5 ºC, it is converted to monoclinic sulphur. Both of these forms are insoluble in water but soluble in carbon disulphide, CS2.





Crystalline forms of sulphur involve Van der Waals bonding.

When sulphur is heated above 113 ºC, it melts to form a pale yellow liquid, which becomes darker and more viscous as the temperature is increased. If liquid sulphur is heated to its boiling point of 445 ºC, and poured into cold water, so-called plastic sulphur is formed, consisting mainly of long chains of sulphur atoms, bound together by covalent bonds:



Sulphur dioxide:
Sulphur burns readily in oxygen and in air forming sulphur dioxide, SO2:



Sulphur dioxide is also formed when sulphites are treated with acids:



or when sulphuric acid is reduced with certain metals:



It is a gas with a choking smell, with a boiling point of -10.0 ºC, which dissolves in water to form the weak acid sulphurous acid, H2SO3, which in turn can give rise to salts (sulphites) containing the sulphite ion SO32-.



Sulphur dioxide is a reducing agent (which reduces ions such as Fe3+, MnO4-, and Cr2O72-:



It is also an oxidising agent oxidising H2S and certain reactive metals such as Mg:





The burning of certain types of coal which contains sulphur causes serious air pollution due to the formation of sulphur dioxide. This ultimately reacts with oxygen and water to form sulphuric acid, which falls to the ground with rain water, the so-called acid rain.



Sulphur dioxide has a boiling point of -10.0 ºC at a pressure of 101 kPa, and a melting point of -75.5 ºC. The molecule is planar, with the sulphur atom attached to two oxygen atoms, making an angle of 119º.

Sulphur trioxide:


Sulphur trioxide, SO3, is formed by the reaction of sulphur dioxide and oxygen, in the presence of catalysts such as platinum or vanadium pentoxide (V2O5). It reacts very readily with water to form sulphuric acid. In industrial practice, SO2,(from the burning of sulphur) is further oxidised catalytically by oxygen (from air), and the resulting SO3 dissolved in concentrated sulphuric acid forming pyrosulphuric acid:



This is the basis for the contact process for the industrial production of sulphuric acid.

Sulphuric acid:



Sulphuric acid, H2SO4, is one of the most important industrial chemicals. It is an oily liquid having a boiling point of 335 ºC, which evolves much heat on dilution with water.

It is a strong diprotic acid, forming salts known as sulphates:



It will turn the indicators litmus, from blue to red, and phenolphthalein from pink to colourless.

In addition, sulphuric acid can act as an oxidising and as a dehydrating agent (removing the elements of water from certain substances).

Dilute sulphuric acid is used as the electrolyte in car batteries (battery acid).

The contact process:
Millions of tons of sulphuric acid are made every year by the contact process, which converts raw sulphur, oxygen and water to sulphuric acid.



Step 1: Melted sulphur is burned in a furnace, using air, producing sulphur dioxide, SO2.

Step 2: The SO2 gas is passed through a tower called a precipitator in order to remove dust and other impurities which might interfere with the catalyst.

Step 3: The SO2 is then washed with water, in a scrubbing tower.

Step 4: The SO2 is then dried in a drying tower.

Step 5: After passing through a heating chamber, the SO2, which is still mixed with air,is passed through a reactor. There, using vanadium pentoxide, V2O5, as catalyst, the SO2 is converted to sulphur trioxide, SO3.



Step 6: Finally, the SO3 is absorbed in concentrated sulphuric acid, giving the so-called oleum or pyrosulphuric acid. This is the diluted with water to give about 98% pure H2SO4.



Hydrogen sulphide:
Hydrogen sulphide, H2S, is the reduced form of sulphur. It is a highly poisonous gas with a smell of rotten eggs. This gas dissolves in water forming a very weakly acidic solution, from which salts known as sulphides (such as Na2S, CaS) may be derived. In turn, H2S may be released from these salts by treatment with strong acids. It is normally produced in the laboratory by reacting iron sulphide with dilute sulphuric acid.





Hydrogen sulphide always behaves as a strong reducing agent, reducing various oxidising agents such as potassium dichromate, K2Cr2O7, and potassium permanganate, KMnO4.



It can also reduce Fe(III) to Fe(II) :



In air, hydrogen sulphide burns with a bluish flame, forming water and sulphur dioxide:



Hydrogen sulphide reacts with many cations in solution to form insoluble metal sulphides. Traces of hydrogen sulphide and other sulphur compounds cause silver objects to tarnish due to the formation of a very thin layer of silver sulphide, Ag2S.


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Allotropes:
Allotropes are different physical forms of an element. For example, carbon exists as graphite and diamond, and sulphur as rhombic and monoclinic forms.




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Oxidizing and reducing agents:
A reducing agent is a substance which can supply electrons to another substance in the course of a chemical reaction.An oxidizing agent is a substance which can accept electrons from another substance in the course of a chemical reaction. (This is discussed in detail in Grade 12).




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Sulphates:
Sulphuric acid used to be commonly known as oil of vitriol, and some of its salts were (and are sometimes still) called vitriols:

Blue vitriol is copper(II) sulphate, CuSO4.5H2O

Green vitriol is iron(II) sulphate, FeSO4.7H2O

White vitriol is zinc sulphate ZnSO4.7H2O

Epsom salt is magnesium sulphate MgSO4.7H2O

Glauber salt is sodium sulphate, Na2SO4.10H2O

Plaster of Paris is calcium sulphate 2CaSO4.H2O

Gypsum is calcium sulphate CaSO4.2H2O.

All soluble sulphates react with an aqueous solution of barium chloride, BaCl2, to form an insoluble precipitate of barium sulphate, BaSO4, insoluble in dilute nitric acid. This forms a test for sulphates.






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Acids:
Acids are substances which can donate protons, H+, to basic substances. If an acid molecule can donate only one proton, it is called a monoprotic acid. Hydrochloric acid, HCl, is an example of a monoprotic acid.

Diprotic acids have molecules that are able to donate two protons. Sulphuric acid, H2SO4, and sulphurous acid, H2SO3, are examples of diprotic acids.




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Metal sulphides:
Many metals react directly with sulphur to form metal sulphides. The sulphides of heavy metals are all insoluble in water. Those of the Group IA and I IA are however soluble in water.

The sulphides of lead, copper, and silver are black. Solutions of lead, silver or copper salts yield black precipitates when exposed to hydrogen sulphide. This may be used as a test for the gas. Zinc suphide is white.

All sulphides react with acids to form a metallic salt and hydrogen sulphide. Some sulphides require concentrated acids for this to happen.