Question

Assign oxidation states to each atom in each element, ion, or compound. \(\begin{array}{llll}{\text { a. } \mathrm{Cl}_{2}} & {\text { b. } \mathrm{Fe}^{3+}} & {\text { c. } \mathrm{CuCl}_{2}} \\ {\mathrm{d} . \mathrm{CH}_{4}} & {\text { e. } \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}} & {\text { f. } \mathrm{HSO}_{4}^{-}}\end{array}\)

Short answer

a. Cl (0), b. Fe (+3), c. Cu (+2), Cl (-1), d. C (-4), H (+1), e. Cr (+6), O (-2), f. H (+1), S (+6), O (-2)

Step-by-step solution

Oxidation State of Cl in Cl2

For the diatomic molecule Cl2, the oxidation state of chlorine (Cl) is 0 because it is in its elemental form.

Oxidation State of Fe in Fe3+

For the ion Fe3+, the oxidation state of iron (Fe) is +3 because it carries a charge of +3.

Oxidation State of Atoms in CuCl2

For CuCl2, since Cl is more electronegative and has an oxidation state of -1, copper (Cu) must balance the charge with a +2 oxidation state to give the compound a neutral charge.

Oxidation State of Atoms in CH4

In CH4, the carbon (C) is in the -4 oxidation state because hydrogen (H), being less electronegative, will have an oxidation state of +1 for each hydrogen atom.

Oxidation State of Atoms in Cr2O7(2-)

For the dichromate ion Cr2O7(2-), oxygen (O) has an oxidation state of -2. Since there are 7 oxygen atoms, the total charge contributed by oxygen is -14. With an overall charge of -2, the two chromium atoms (Cr) combined must have an oxidation state of +12. Therefore, each Cr atom has an oxidation state of +6 to balance the -14 from the oxygen atoms.

Oxidation State of Atoms in HSO4-

In the bisulfate ion HSO4-, the oxygen (O) is assigned an oxidation state of -2. Therefore, the four oxygen atoms contribute -8 to the charge. Since sulfur (S) is more electronegative than hydrogen but less than oxygen, it will have a positive oxidation state. Hydrogen (H) will have an oxidation state of +1. To balance the total charge of -1 for the ion, the sulfur must have an oxidation state of +6.

Key concepts

Oxidation State Rules

Understanding oxidation state rules is essential for mastering redox chemistry. An oxidation state, often called an oxidation number, is a theoretical charge that an atom would have if all bonds were ionic. These rules help in determining how electrons are distributed among atoms in a compound.

• Atoms in their elemental form (e.g., O2, Cl2) have an oxidation state of 0.
• For monoatomic ions (e.g., Na+, Al3+), the oxidation state is equal to the ion's charge.
• Oxygen typically has an oxidation state of -2, except in peroxides where it's -1.
• Hydrogen usually has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals.
• The sum of oxidation states in a neutral compound is 0, and in a polyatomic ion, it is equal to the ion's charge.

Applying these rules systematically can simplify determining oxidation states, turning a complex task into a series of manageable steps.

Determining Oxidation Numbers

Determining oxidation numbers is like solving a puzzle where the pieces are the atoms in a molecule. Start with what you know: the oxidation state of elements in their standard state is always zero, and familiar charges on common ions. For example, hydrogen is nearly always +1 and oxygen -2, with rare exceptions.

When assigning oxidation numbers, one should confer with a periodic table as it can inform on the electronegativities of elements and their common ionic states. For instance, in a compound like CuCl2, we know chlorine is more electronegative and generally has an oxidation state of -1. Since there are two chlorine atoms, this adds up to -2, requiring copper to have an oxidation state of +2 in order for the compound to be neutral.

Exercise Improvement Advice

To assist students in understanding the nuances, it would be helpful to point out specific rules being applied at each step of the solution. This way, not only will students know the oxidation states but also the reasoning behind them.

Redox Chemistry

Redox chemistry is the study of reactions involving changes in oxidation states, known as oxidation-reduction (redox) reactions. These are processes where one substance transfers electrons to another, altering the oxidation states of certain atoms.

In a redox reaction, oxidizing agents gain electrons and thus are reduced, while reducing agents lose electrons and are oxidized. Recognizing changes in oxidation numbers is vital for identifying the substances that are oxidized and reduced.

Real-world Connection

Redox reactions are fundamental to many natural and industrial processes, such as photosynthesis, metabolism, and battery operation. Grasping how to determine oxidation states is not just for academic exercises but also provides insight into the workings of the world around us. Additionally, sometimes the oxidation number can reveal if a redox reaction is even possible, determining the feasibility of such processes in practical applications.