Chemistryyy (1 Viewer)

[drsabz101]

Okay, I am a bit confused on how to write redox and oxidisation equations. Also from this , how to identify which substance is the oxidising agent or reducing agent.

Can someone please go through an example and steps required.

Thanks!

 

[Shuuya]

Ok, here's an example:

Cl2 + 2HBr --> 2HCl + Br2

To be able to tell which is the reducing agent/oxidising agent, you should write out the ionic equation:

Cl2 + 2H+ + 2Br- ---> 2H+ + 2Cl- + Br2

From here it can be seen that the H+ ions are spectator ions, so the net ionic equation is:

Cl2+ 2Br- ---> 2Cl- + Br2

If you split these into the half (redox) equations you get:

Cl2 + 2e- ---> 2Cl-

2Br- ---> Br2 + 2e-

Now you have to look at the changes in charge of Cl and Br:

Cl goes from being 0 to -1, and therefore it has gained an electron (i.e has been reduced)
Br goes from being -1 to 0, and therefore it has lost an electron (i.e has been oxidised)

Now, the confusing bit is that a species that undergoes oxidation is known as the reducing agent, and the one that undergoes reduction is known as the oxidising agent.

Therefore, Cl is the oxidising agent and Br is the reducing agent.

Hope that helped! Feel free to ask if something doesn't make sense

 

[DatAtarLyfe]

Just to make it less confusing, in a redox reaction, the more reactive species undergoes oxidation, as it is more inclined to lose its electrons. So that species is donating its electrons to the other species, causing the other species to undergo reduction. Hence the species that undergoes oxidation, also induces a reduction reaction to the other species, and so is known as the reducing agent.
The converse holds true as the species undergoing reduction gains electrons from the other species, thereby causing the other species to undergo oxidation. This is why its called the oxidising agent

 

[leehuan]

Shouldn't this be under the HSC Chemistry section?

Anyway, useful acronyms:
OIL-RIG
OxidationIsLoss-ReductionIsGain

What is lost? Always electrons


Later when you do cells:
RedCat (Reduction at Cathode)
AnOx (Oxidation at Anode)

The reducing agent (aka reductant) GETS oxidised
The oxidising agent (aka oxidant) GETS reduced

 

[drsabz101]

"From here it can be seen that the H+ ions are spectator ions, so the net ionic equation is:

Cl2+ 2Br- ---> 2Cl- + Br2"

- What are spectator ions???
- Why does Cl- have the minus next to it (is that because it has to lose 1 electron to gain an full shell)
- Also so an ionic equation is basically breaking down all the compounds in the equation wit their charge (?)

 

[DatAtarLyfe]

"From here it can be seen that the H+ ions are spectator ions, so the net ionic equation is:

Cl2+ 2Br- ---> 2Cl- + Br2"

What are spectator ions???

- Also so an ionic equation is basically breaking down all the compounds in the equation wit their charge (?)

spectator ions are ions that do not gain or lose electrons throughout the redox reaction i.e. they remain the same and simply 'spectate' the reaction

yeh so net ionic equations depict the charge of each atom and ion in the reaction. It makes it easier to see which atoms gained or lose electrons, as well as to determine oxidation states

 

[drsabz101]

ohh okay. does does Cl- have the minus next to it because it has to lose 1 electron to gain an full shell

 

[DatAtarLyfe]

ohh okay. does does Cl- have the minus next to it because it has to lose 1 electron to gain an full shell

chlorine is an anion, meaning it wants to gain electrons (due to its high electro negativity). So, during ionic bonding, it gains one electron (as it already has 7 and needs one more) from the cation to gain a complete shell.
In it's natural state, chlorine is neutral charged (i.e. same number of electrons and protons) however, the moment it gains that electron to become an anion, there are more electrons then protons, and so the chlorine ion becomes negatively charged by one (due to the one extra electron). It's this negative charge that allows it to bond with the cations.

Hence, the -1 is because chlorine has gained one electron, and is now negatively charged due to an excess of one electron
(soz, got carried away with the description XDD)

 

[drsabz101]

thankyou, I get it!! It's okay I like in depth explanations

 

[DatAtarLyfe]

anytime

 

[drsabz101]

If you split these into the half (redox) equations you get:

Cl2 + 2e- ---> 2Cl-

2Br- ---> Br2 + 2e-

So what is happening in these redox equations. Like why is 2CL- on the right hand side and 2Br- on the left

 

[DatAtarLyfe]

It makes more sense if i do this: divide the top equation by two:
1/2 Cl2 + e- ----> Cl-
Forget about the half in front of the Cl, but what basically happened was that you had a diatomic molecule of chlorine (2 atoms) and it gained 1 electron from the bromine (the bromine oxidised, i'll get to that after), causing a Cl-, as now it has an excess of one electron. Thus it went through reduction.
Now with that in mind, if you go back to the original equation, the diatomic molecule gained 2 electrons, causing the each Cl in the molecule to gain one electron each. This coincided with the 2 Cl ions that were created (they weren't actually created, but actually split apart from from the diatomic molecule after the reduction reaction)
Same principle applies to the second equation, we had 2 bromine ions, each with an excess of one electron (hence together they have two electrons in excess). They then oxidised (lose/donated) their electrons to the chlorine, and then joined together to form the diatomic molecule

For redox equations, it's all about looking at them logically. In reduction equations, the species goes from being a neutral atom to becoming an anion. To become an anion, the atom must gain the electrons first, and then the product becomes an ion. Hence the atom + electron ---> anion
For oxidation, the ion becomes an atom. For it to become an atom, the ion must lose electrons. So if you write:
ion - electron ----> atom and move the electron to the otherside, you get ion----> atom + electron

 

[drsabz101]

Thankyou again for the help!

How to you know though if an atom is neutral ?

And is it true metals oxidise first

 

[drsabz101]

Feel free to post more examples guys OR practice questions that I can do
thnx

 

[drsabz101]

Okay so here's another problem I found:

for the following reaction : Aluminium with sulfuric acid
- when writing the reduction equation, why do we only o this for hydrogen gas and disregard the sulfuric acid

 

[DatAtarLyfe]

Thankyou again for the help!

How to you know though if an atom is neutral ?

And is it true metals oxidise first

Generally in equations, the atom will have a -ve or a +ve sign next to the number of electrons its gained or lost. Remember, if its positive, then its lost. If its negative, then its gained.
In more complex equations with polyatomic ions, the atoms that are stand alone (not bonded with other atoms) are neutral. The compounds are the ones that have ions.
For example in the reaction:
Mg +2HCl ----> MgCl2 + H2O
The magnesium atom on the LHS of the eqn is in its neutral state, whereas the HCl has an ion of H+ and Cl-
As an exercise, see if you can post the redox equation of this example. Also state the spectator ion

Yes that is true, metals tend to oxidise faster. This is due to a property known as electronegativity (an atoms tendency to want to take electrons, hence stronger pull of electrons)
Atoms on the LHS of the periodic table (metals) have low electronegativity, and so they donate their electrons willingly, and so metals tend to undergo oxidation.
Atoms on the RHS of the periodic table (non-metals) have high electronegativity and thus want to gain electrons, and tend to undergo reduction

 

[drsabz101]

before I write the reduction equation is the net ionic equation:

Mg + 2H+ +2Cl -1 -> Mg2+ + 2Cl- + 2H+ + O2-

 

[Shuuya]

before I write the reduction equation is the net ionic equation:

Mg + 2H+ +2Cl -1 -> Mg2+ + 2Cl- + 2H+ + O2-

DatAtarLyfe's equation should be Mg +2HCl ----> MgCl2 + H2 (not H2O, as acid+metal-->salt+hydrogen gas)

So Ionic equation is:

Mg + 2H+ + 2Cl- ---> Mg2+ + 2Cl- + H2

Hopefully you can write the net ionic equation now (i.e remove the spectator ions)

 

[drsabz101]

I was thinking the same thing

 

[Shuuya]

Okay so here's another problem I found:

for the following reaction : Aluminium with sulfuric acid
- when writing the reduction equation, why do we only o this for hydrogen gas and disregard the sulfuric acid

This is also to do with the spectator ions:

Al(s) + 3H2SO4(aq) ---> Al2(SO4)3(aq) + 3H2(g) (I find that it helps to write down states to see clear changes)

So the ionic equation is:

Al + 6H+ +3(SO4)2- ---> 2Al3+ +3(SO4)2- + 3H2

So you can see that the SO4 is the spectator ion, as it stays the same on both sides of the equation. So the net ionic equation is:

Al + 6H+ ---> 2Al3++ 3H2

From there the redox equations are:

Al ---> 2Al3+ + 6e- （oxidation)
6H+ 6e- ---> 3H2 (reduction)