Indicate whether each reaction is a redox reaction. If it is, which atom gets oxidized and which atom gets reduced? Consult the shortcut rules. \(\mathrm{P}_{4}+6 \mathrm{Br}_{2} \rightarrow 4 \mathrm{PBr}_{3}\) (Hint: \(\mathrm{Br}\) is more electronegative than P.)

A redox reaction involves the transfer of electrons, where one species gets reduced (gains electrons) and the other species gets oxidized (loses electrons). Analyze the given chemical reaction to see if there is any transfer of electrons taking place: Reaction: \(\mathrm{P}_{4}+6 \mathrm{Br}_{2} \rightarrow 4 \mathrm{PBr}_{3}\) According to the hint, Br is more electronegative than P, meaning that Br has a greater tendency to attract electrons. This suggests that a transfer of electrons may be occurring in this reaction.

To decide whether a redox reaction is occurring, we need to determine the oxidation number of each element before and after the reaction. This will help us identify if any electron transfers are happening. The reaction again: \(\mathrm{P}_{4}+6 \mathrm{Br}_{2} \rightarrow 4 \mathrm{PBr}_{3}\) Before reaction: 1. P in \(\mathrm{P}_{4}\) has an oxidation state of 0 since it is an uncombined element. 2. Br in \(\mathrm{Br}_{2}\) has an oxidation state of 0 since it is an uncombined element. After reaction: 1. P in \(\mathrm{PBr}_{3}\) has an oxidation state of +3 (for every 1 P atom is bonded to 3 Br atoms, each with an oxidation state of -1). 2. Br in \(\mathrm{PBr}_{3}\) has an oxidation state of -1 (each Br atom has gained 1 electron from P).

Now that we have identified the oxidation states of each element before and after the reaction, we can determine which element is being oxidized and which one is reduced: 1. P in \(\mathrm{P}_{4}\) has an initial oxidation state of 0 and ends up with an oxidation state of +3 in \(\mathrm{PBr}_{3}\). This means that P has lost 3 electrons and is being oxidized. 2. Br in \(\mathrm{Br}_{2}\) has an initial oxidation state of 0 and ends up with an oxidation state of -1 in \(\mathrm{PBr}_{3}\). This means that Br has gained one electron and is being reduced.

Since there has been a transfer of electrons happening in this reaction (P loses electrons and Br gains electrons), we can confirm that this reaction is a redox reaction.

The given reaction, \(\mathrm{P}_{4}+6 \mathrm{Br}_{2} \rightarrow 4 \mathrm{PBr}_{3}\), is indeed a redox reaction. The P atom gets oxidized, and the Br atom gets reduced.