Consider the following list of reagents Acidified \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\), alkaline \(\mathrm{KMnO}_{4}, \mathrm{CuSO}_{4}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Cl}_{2}, \mathrm{O}_{3}, \mathrm{~F} \mathrm{CCl}_{3}, \mathrm{HNO}_{3}\) and \(\mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}\). The total number of reagcnts that can oxidise aqueous iodide to iodine is (The answer to the question is a single-digit integer, ranging from 0 to 9)

Understanding the standard electrode potential is crucial when studying redox reactions in chemistry. It is a measure of the tendency of a chemical species to be reduced under standard conditions, which are typically 25°C, 1 M concentration for all solutes, a 1 atm pressure for any gases involved, and a pH of 7 for reactions in aqueous solutions. The potential is measured in volts (V) and is referenced against the standard hydrogen electrode (SHE), which is assigned a potential of 0 V.

For instance, in the context of the exercise provided, substances like acidified \(K_2Cr_2O_7\) and alkaline \(KMnO_4\) have higher electrode potentials compared to the iodide ion \(I^-\), which means they can act as strong oxidizing agents and have the ability to oxidize iodide to iodine. The standard electrode potentials are typically found in electrochemical series charts—a higher positive value indicates a stronger oxidizing agent, capable of accepting electrons more readily. Knowing these values helps us predict which substances can oxidize others in a chemical reaction.

Redox equations represent the chemical changes that occur during oxidation-reduction (redox) reactions, where the transfer of electrons between species takes place. Oxidation involves the loss of electrons, while reduction involves the gain of electrons. In a balanced redox equation, the number of electrons lost must equal the number of electrons gained to maintain electrical neutrality.

For example, when \(KMnO_4\) acts as an oxidizing agent, it will gain electrons and be reduced, the products of this change must reflect this gain in electrons. Similarly, when \(I^-\) ions are oxidized to \(I_2\), they lose electrons, which should be accounted for in the redox equation. The ability to write and balance redox equations is fundamental in understanding how different reagents, like the ones listed in the exercise, interact with each other during redox reactions.

Oxidation states (also called oxidation numbers) are hypothetical charges that an atom would have if all bonds to atoms of different elements were completely ionic. These numbers help in determining how electrons are distributed in molecules and ions, and they play a key role in identifying which atoms are oxidized or reduced during redox reactions.

Determining oxidation states is straightforward for monoatomic ions: the oxidation state is equal to the charge of the ion. However, for molecules or polyatomic ions, the sum of the oxidation states of all atoms must equal the charge of the molecule or ion. In the original exercise, recognizing that \(I^-\) has an oxidation state of -1 helps us see that it can be oxidized to iodine \(I_2\), where each iodine atom has an oxidation state of 0. This implies each iodide ion would need to lose an electron to achieve this state change. Recognizing and calculating oxidation states is therefore essential when predicting the outcome of redox reactions and understanding the behavior of various chemical species.