In the following equation, which species is the oxidizing agent?
KCLO3 --> KCL + KCLO4
In the same are, How do you know that Cu is the reducing agent in this equation:
AgNO3 + Cu ---> Cu(NO3)2 + Ag
Two definitions cover the whole gamut of questions (actually, only one definition is needed; if you know one the other must be the opposite).
Oxidation is the loss of electrons. An easy way to remember this is LEO the lion goes grrr. LEO stands for loss electrons oxidation. Easy, huh?
(You don't need to remember the definition for reduction but it is "Reduction is the gain of electrons."
Now use oxidation states. For the first one, which is a tricky one.
One half reaction is ClO3^- ==> Cl^-
and the other is ClO3^- ==> ClO4^-
So Cl is +5 going to -1 in the first one (reduction) and +5 going to +7 (oxidation) in the second one.
For the Ag/Cu equation,
Ag goes from +1 on the left to zero on the right which is a gain of electrons.
Cu goes from zero on the left to +2 on the right which is a loss of electrons. Cu is oxidized (which makes it the reducing agent).
To determine the oxidizing and reducing agent in a chemical equation, you need to identify the changes in oxidation numbers for different elements. Here's how you can find the oxidizing and reducing agents in the given equations:
1. Equation: KCLO3 → KCL + KCLO4
In this equation, we need to determine the species that is being reduced and the one that is being oxidized.
First, let's assign oxidation numbers to each element in the equation:
K: +1
Cl: -1 in KCl and +7 in KCLO4
O: -2 in KCLO3 and -2 in KCl
The oxidation state of chlorine increases from -1 to +7, indicating that it gains electrons and is being oxidized. Therefore, KCLO3 is the oxidizing agent because it causes the oxidation of chlorine. On the other hand, KCl does not change its oxidation state, so it is not involved in the redox reaction.
2. Equation: AgNO3 + Cu → Cu(NO3)2 + Ag
To identify the reducing agent in this equation, we again assign oxidation numbers to each element:
Ag: +1
N: +5 in AgNO3 and +5 in Cu(NO3)2
O: -2 in AgNO3 and -2 in Cu(NO3)2
Cu: 0
In this case, the oxidation state of copper (Cu) changes from 0 to +2 in Cu(NO3)2, indicating that it loses electrons and is being oxidized. Therefore, copper (Cu) is the reducing agent because it causes the reduction of silver (Ag). Conversely, silver (Ag) gains electrons, going from +1 to 0 and is being reduced.
Remember, in a redox reaction, the oxidizing agent is the species that causes oxidation by accepting electrons, while the reducing agent is the species that causes reduction by donating electrons.