Question

Various oxidation states of few elements are mentioned. Which of the options is not correctly matched? (a) Phosphorus: \(+3\) to \(+5\) (b) Nitrogen: \(+1\) to \(+5\) (c) lodine: \(-1\) to \(+7\) (d) Chromium : \(-3\) to \(+6\)

Short answer

The incorrectly matched option for oxidation states is (d) Chromium: -3 to +6.

Step-by-step solution

Understanding Oxidation States

Oxidation states are the hypothetical charges that an atom would have if all bonds to atoms of different elements were 100% ionic. They provide a way to keep track of electrons in oxidation-reduction reactions. It is important to know the common oxidation states of elements to identify whether the given states are feasible.

Verifying Common Oxidation States for Each Element

For each of the given elements, verify their common oxidation states with known values from the periodic table. Phosphorus commonly shows oxidation states of +3 and +5, Nitrogen's oxidation states range from -3 to +5, Iodine has a range from -1 to +7, and Chromium typically shows oxidation states from +2 to +6. Any value outside these common oxidation states would be considered incorrectly matched.

Identifying the Incorrectly Matched Option

Compare the given ranges with the verified ranges for each element. Upon inspection, it is found that for Chromium the given oxidation state range is -3 to +6; however, -3 is not a commonly known or accepted oxidation state for Chromium. The common oxidation states of Chromium are +2, +3, and +6. Thus, the provided range for Chromium is incorrect.

Key concepts

Common Oxidation States

Understanding the common oxidation states of elements is a fundamental aspect of chemistry, particularly when delving into the behavior of atoms during chemical reactions. Oxidation states, often referred to as oxidation numbers, indicate the degree of oxidation of an atom in a chemical compound. How do we determine these numbers? When assigning oxidation states, the rules begin by looking at lone elements, which have an oxidation state of zero. For ions, the oxidation state is equal to the charge of the ion, for example, Na+ has an oxidation state of +1.

However, as we encounter compounds, the situation gets more complex. The oxidation states must sum up to the charge of the compound. In a molecule of water (H₂O), oxygen has an oxidation state of -2, while hydrogen has an oxidation state of +1. This concept can become intricate with transition metals or nonmetals with variable oxidation states. For example, iron can have an oxidation state of +2 or +3, significantly affecting the chemical properties of the compounds in which it participates. Recognizing and predicting common oxidation states can be crucial when balancing chemical equations or predicting reaction products.

In the given exercise, the problem revolves around identifying which of the listed oxidation states for particular elements do not match their common occurrences. Students should remember that the 'common' oxidation states are those found most frequently and are generally related to the element's position in the periodic table.

Periodic Table Elements

The periodic table is essentially a tabular display of the known chemical elements, organized by their atomic numbers, electron configurations, and recurring chemical properties. Elements are presented in order of increasing atomic number, which is the number of protons in an element's atom, and are aligned in such a way that elements that exhibit similar behavior fall into the same columns, known as groups.

Each element on the periodic table can exhibit one or several oxidation states in its compounds. These states can often be guessed based on the element's group. For instance, group 1 elements (the alkali metals) almost always have an oxidation state of +1 since they have one electron to lose. Similarly, group 17 elements (the halogens) frequently display an oxidation state of -1 because they tend to gain an electron.

Transition metals, found in groups 3 through 12, are known for having multiple possible oxidation states. This is because their d-orbitals allow for a variety of different electron configurations. An adept reader of the periodic table can anticipate the common oxidation states of an element based on its position, which is particularly useful when analyzing the feasibility of a given oxidation number in an exercise, as seen with the element chromium in the original problem.

Redox Reactions

Redox reactions are chemical reactions involving the transfer of electrons between two species. These reactions are characterized by the movement of electrons from one atom or molecule (the reducing agent) to another (the oxidizing agent). Redox reactions are essential for various biological processes and industrial applications.

Within a redox reaction, the substance that loses electrons is said to be oxidized, while the substance that gains electrons is reduced. The common oxidation states of elements play a pivotal role in determining how they interact in redox reactions. To illustrate, in a reaction, an element with a common oxidation state of +3 may be readily reduced to an oxidation state of +2 by gaining an electron.

Oxidation states are changed with respect to the standard hydrogen electrode, which is assigned an oxidation state of zero. Balancing redox reactions requires attention to the changes in oxidation states and the stoichiometry of the reaction. When students encounter a redox problem, they're often required to calculate the oxidation states of each participating atom in order to balance the overall equation. The exercise provided was about identifying incorrect oxidation states and hence plays into understanding how redox reactions work, helping one to understand both the concepts of electron transfer and the importance of correct oxidation state values in chemical reactions.