Oxidising and reducing agents
Watch this to understand oxidising and reducing agents for Higher Chemistry (and how to fix a trainline in the Canadian wilderness)
Canada is one of the most beautiful countries in the world. Nature at its most beautiful.
However, only a very small part of the country is actually inhabited. So, a lot of places are remote, difficult to access and far away from any assistance or modern technology.
But things still need to run... on time. In the middle of nowhere train tracks sometimes need to be repaired. Engineers have to use their ingenuity AND one of the most impressive chemical reactions you’ll ever see in a classroom. And it’s one of my personal favourites, the thermite reaction.
The temperature achieved is so high one of the products is liquid iron. Which is used to weld train tracks back together. A bit extreme but very efficient.
So what chemical reaction can reach such power, and can we think of an even more powerful one?
Let’s raise the temperature.
This is thinkfour
The thermite reaction is an example of a highly exothermic displacement reaction. Iron (III) oxide will react with aluminium to produce aluminium oxide and the iron we were just talking about. In this reaction, aluminium will displace Iron from its oxide. This is possible because aluminium is more reactive, meaning it will have a higher disposition to lose electrons and form ions than iron which is less reactive and will accept the electrons.
This can be shown in Ion-electron equations:
In the oxidation reaction, Aluminium will lose 3 electrons and become the aluminium ion Al3+
And in the reduction reaction, the iron (III) ion Fe3+ will gain 3 electrons to form Iron metal
Combining the two which happen simultaneously will display the overall reaction called redox reaction.
So in this example, aluminium is “pushing” iron to reduce by making these 3 electrons available. We can then call aluminium the reducing agent, a substance that accepts electrons.
The other way around the iron (III) ion is “allowing” aluminium to oxidise by accepting these 3 electrons. We can then call iron the oxidising agent, a substance that donates electrons.
To identifying which substance is the reducing agent and which one is the oxidising agent, let have a look at an example. Magnesium displaces silver from a silver nitrate.
Magnesium metal will lose 2 electrons to become magnesium ion Mg2+ and Silver ions Ag+ will gain one electron to become silver metal. The redox reaction will then be: 2Ag+(aq) + Mg(s) --> 2Ag(s) + Mg2+(aq).
The reason why an element will tend to lose or to give electrons away is related to its electronegativity. We can predict that an element with high electronegativity like the non-metals will tend to attract electron and act as oxidising agents. While on the other side of the periodic table, the alkali metals for example, are very low on the Pauling scale and will act as reducing agents and form positively charged ions.
Not only elements but compounds can be involved in redox reaction as a reducing or an oxidising agent. The whole Iron industry would not be possible without the strong reducing agent carbon monoxide. Other example can be seen across the electrochemical series.
Some of these redox reactions can be extremely exothermic like the thermite reaction. Which we’ve seen can be used to produce molten iron in the remote parts of Canada.
But there is even more... Why do you think I took the Magnesium and Silver Nitrate example? This is definitely my favourite demonstration. Just a tiny spatula, smaller than a rice grain of both reactant, in solid form are dying to react. A small drop of water to put the two together and boom! Fireworks!
This was Think Four. Thanks for watching.