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The phosphate ion, denoted as \(\mathrm{PO}_{4}^{3-}\), is a polyatomic ion consisting of one phosphorus (P) atom and four oxygen (O) atoms. This ion carries an overall charge of -3, which impacts the distribution of oxidation numbers within it. It's a key player in many biological processes, including energy transfer in cells, often found as part of adenosine triphosphate (ATP). Understanding its structure and oxidation states is important in chemistry when dealing with various reactions and compounds.In this ion, the combined oxidation numbers of all the atoms must equal the ion's total charge. For the phosphate ion, this balance is a key principle that helps in calculating individual oxidation states of the atoms involved.

Typically, oxygen has an oxidation state of -2 in most compounds. This is because oxygen is highly electronegative and tends to attract electrons during bonding. In the case of the phosphate ion \(\mathrm{PO}_{4}^{3-}\), there are four oxygen atoms. Therefore, the cumulative oxidation number contributed by the oxygen atoms is:- Each oxygen: \(-2\) - Total for four oxygen atoms: \(-2 \times 4 = -8\) This large negative value from the oxygen atoms contributes significantly to the overall charge of the ion and helps us solve for other atoms' oxidation states within the ion.

To find the oxidation state of phosphorus in the \(\mathrm{PO}_{4}^{3-}\) ion, we must consider the overall charge balance dictated by the sum of oxidation states. The equation that represents this balance is:\[ x + (-8) = -3 \]Where \(x\) is the oxidation number of phosphorus. Solving this equation, we find:- \(x = -3 + 8 = +5\) This indicates that phosphorus has an oxidation state of +5 in the phosphate ion. This positive value balances out the negative contributions from the oxygens, resulting in the net ion charge of -3. Understanding these rules of oxidation states helps predict reactivity and interactions in chemistry.