For each unbalanced equation given below- write unbalanced half-reactions. identify the species oxidized and the species reduced. Id identify the oxidizing and reducing agents. (a) \(\mathrm{Ag}(s)+\mathrm{NO}_{3}^{-}(a q) \longrightarrow \mathrm{Ag}^{+}(a q)+\mathrm{NO}(g)\) (b) \(\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{O}_{2}(g)\)

In the world of chemistry, each redox reaction can be broken down into two key components: oxidation and reduction. The oxidation half-reaction specifically deals with the loss of electrons by a species. During this process, the oxidation state of the species increases.

For instance, in the equation
\[\mathrm{Ag}(s) \longrightarrow \mathrm{Ag}^{+}(aq) + e^{-}\]
silver metal (Ag) loses an electron to form the silver ion (Ag+). In this reaction, the silver, starting with an oxidation state of 0, is oxidized to an oxidation state of +1. Thus, the oxidation half-reaction reflects the half of the redox process where electron loss – and an increase in oxidation state – occurs.

It's important to always remember that during the oxidation half-reaction, electrons are produced as reactants lose them. These lost electrons are then available for the reduction half-reaction.

Opposite to oxidation, the reduction half-reaction is where a species gains electrons, leading to a decrease in its oxidation state. A simple way to remember this is the mnemonic 'OIL RIG', which stands for 'Oxidation Is Loss, Reduction Is Gain' - referring to electrons.

In our example, the nitrate ion
\[\mathrm{NO}_{3}^{-}(aq) + 3e^{-} \longrightarrow \mathrm{NO}(g) + 2\mathrm{H}_{2}\mathrm{O}(l)\]
we see that nitrate (NO3-) gains electrons to form nitrogen monoxide (NO). The oxidation state of nitrogen thereby decreases from +5 to +2. This half-reaction showcases the gain of electrons, which is the hallmark of reduction processes in redox reactions.

Moreover, the reduction half-reaction provides the electrons needed for the oxidation half-reaction to occur, thereby completing the redox cycle. It's a harmonious exchange, where one species' loss is another's gain.

In any redox reaction, we encounter two key players: the oxidizing and reducing agents. The oxidizing agent is the species that gets reduced by taking electrons from another species. Conversely, the reducing agent is the species that gets oxidized by giving away its electrons.

Let's use the earlier reaction for clarification:
Silver (Ag) is the reducing agent because it provides electrons for reduction and gets oxidized itself. On the other hand, nitrate (NO3-) acts as the oxidizing agent as it accepts electrons from silver and gets reduced.

Identifying these agents is crucial in understanding redox reactions, as they drive the process. The oxidizing agent, by accepting electrons, causes another species to oxidize, while the reducing agent, by donating electrons, causes another species to reduce.

When analyzing redox reactions, identifying the species that are oxidized and reduced is fundamental. The species oxidized is the one that loses electrons and experiences an increase in oxidation state.

In the reaction:
\[\mathrm{Ag}(s) \longrightarrow \mathrm{Ag}^{+}(aq) + e^{-}\]
silver (Ag) is oxidized as it loses an electron. Conversely, the species reduced is the one that gains electrons and undergoes a decrease in oxidation state, such as nitrate in:
\[\mathrm{NO}_{3}^{-}(aq) + 3e^{-} \longrightarrow \mathrm{NO}(g) + 2\mathrm{H}_{2}\mathrm{O}(l)\]
Understanding these concepts helps us see the electron flow within the reaction, giving us insight into the mechanics of redox processes, as well as the roles different substances play within them.

Remember, the species oxidized and reduced are often transformed into the reducing and oxidizing agents, respectively, acting to propel the redox reaction to its completion.