Problem 50 - Chemistry: The Central Science in SI Units, Global Edition

Question

Determine the oxidation number for the indicated element in each of the following substances: (a) \(\mathrm{N}\) in \(\mathrm{N}_{2} \mathrm{H}_{4}\), (b) \(\mathrm{Nin} \mathrm{NO}_{2}\), (c) \(\mathrm{Mn}\) in \(\mathrm{MnCl}_{3}\) (d) Fe in \(\mathrm{FeSO}_{4^{\prime}}\) (e) \(\mathrm{Pt}\) in \(\mathrm{PtCl}_{4}\) (f) \(\mathrm{Cl}\) in \(\mathrm{NaClO}_{4}\).

Step-by-Step Solution

Answer

The oxidation numbers for the indicated elements in the following substances are: (a) N in \(\mathrm{N}_{2} \mathrm{H}_{4}\) is -2, (b) N in \(\mathrm{NO}_{2}\) is +4, (c) Mn in \(\mathrm{MnCl}_{3}\) is +3, (d) Fe in \(\mathrm{FeSO}_{4}\) is +2, (e) Pt in \(\mathrm{PtCl}_{4}\) is +4, and (f) Cl in \(\mathrm{NaClO}_{4}\) is +7.

Step 1: Use rule 2 and the known oxidation numbers of hydrogen

Since hydrogen typically has an oxidation number of +1, and there are four hydrogen atoms in the compound, the sum of the oxidation numbers for hydrogen is +4. Since the compound is neutral, the sum of the oxidation numbers for nitrogen must be -4. There are two nitrogen atoms, so the oxidation number of each nitrogen atom is -2. (b) Determine the oxidation number of N in \(\mathrm{NO}_{2}\)

Step 1: Use rule 2 and the known oxidation numbers of oxygen

Oxygen typically has an oxidation number of -2, and there are two oxygen atoms in the compound, so the sum of the oxidation numbers for oxygen is -4. Since the compound is neutral, the oxidation number of nitrogen must be +4. (c) Determine the oxidation number of Mn in \(\mathrm{MnCl}_{3}\)

Step 1: Use rule 2 and the known oxidation numbers of chlorine

Chlorine typically has an oxidation number of -1, and there are three chlorine atoms in the compound, so the sum of the oxidation numbers for chlorine is -3. Since the compound is neutral, the oxidation number of manganese must be +3. (d) Determine the oxidation number of Fe in \(\mathrm{FeSO}_{4}\)

Step 1: Use rule 3 and the known oxidation numbers of oxygen, sulfur, and the charge of the polyatomic ion

First, note that sulfate ion, \(SO_{4}^{2-}\), is a polyatomic ion with a charge of -2. The oxidation number of oxygen is -2, and there are four of them, contributing an oxidation number of -8. In the sulfate ion, the oxidation number of sulfur must be +6 to make the sum of the oxidation numbers equal to -2. Now, use the neutral compound rule (rule 2) for \(\mathrm{FeSO}_{4}\). The oxidation numbers of all other elements sum to -2, so the oxidation number of iron must be +2. (e) Determine the oxidation number of Pt in \(\mathrm{PtCl}_{4}\)

Step 1: Use rule 2 and the known oxidation numbers of chlorine

Chlorine typically has an oxidation number of -1, and there are four chlorine atoms in the compound, so the sum of the oxidation numbers for chlorine is -4. Since the compound is neutral, the oxidation number of platinum must be +4. (f) Determine the oxidation number of Cl in \(\mathrm{NaClO}_{4}\)

Step 1: Use rule 2 and the known oxidation numbers of sodium, oxygen, and the charge of the polyatomic ion

Sodium, an alkali metal, has an oxidation number of +1. There is an oxygen atom bonded to a more electronegative chlorine atom in the \(\mathrm{ClO}_{4}\) ion, so the oxidation number of oxygen is still -2. The compound is neutral, and the sum of the oxidation numbers for sodium, oxygen, and the other chlorine atoms in the perchlorate ion is -7, so the oxidation number of the central chlorine atom must be +7.

Key Concepts

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, are essential processes in chemistry where the oxidation state of substances changes. These reactions involve the transfer of electrons between substances.

• Oxidation is the process of losing electrons, leading to an increase in oxidation number.
• Reduction is the process of gaining electrons, leading to a decrease in oxidation number.

Both processes occur simultaneously. While one substance gets oxidized, another gets reduced. This electron exchange is what maintains the balance in redox reactions.
To identify redox reactions, look for changes in the oxidation states of the elements involved. It's also crucial to remember that redox reactions are not limited to ionic compounds. They can also occur in covalent compounds as demonstrated by the variation of oxidation numbers within these molecules.
Every chemical reaction involving the change of oxidation states is considered a redox reaction, making this concept incredibly important in both inorganic and organic chemistry.

Chemical Compounds

Chemical compounds are substances formed from two or more chemical elements bonded together. The nature of these bonds can vary, typically being ionic or covalent. Understanding chemical compounds is essential in determining the oxidation states of their constituent elements.

• Ionic compounds are formed by the complete transfer of electrons from one atom to another, resulting in oppositely charged ions.
• Covalent compounds involve the sharing of electrons between atoms.

Compounds have a fixed ratio of elements, giving them distinct properties. Recognizing the type of compound helps simplify the calculation of oxidation numbers. For instance, in ionic compounds like \(\text{NaCl}\), the oxidation numbers directly relate to the charges on the ions, with \(\text{Na}^+\) being +1 and \(\text{Cl}^-\) being -1. In covalent compounds such as \(\text{H}_2\text{O}\), oxidation states help understand how electrons are distributed within the molecule.
Understanding the properties of chemical compounds is vital in predicting their behavior during reactions, especially in determining how electron shifts influence reaction paths and outcomes.

Oxidation States

Oxidation states, or oxidation numbers, are a helpful concept in chemistry, used to track the transfer of electrons in chemical reactions. These numbers represent the theoretical charge an atom would have if all bonds were ionic.

• Positive Oxidation States: Indicate a loss of electrons. For example, \(\text{Fe}^{2+}\) has a +2 oxidation state, implying iron has lost two electrons.
• Negative Oxidation States: Indicate a gain of electrons. In \(\text{NO}_{2}\), oxygen has an oxidation state of -2, suggesting it has gained electrons.
• Zero Oxidation State: Elements in their natural, uncombined form have an oxidation state of zero. For example, \(\text{O}_2\) and \(\text{N}_2\) each have a zero oxidation state.

The rules for assigning oxidation numbers include:

• Oxygen is usually assigned an oxidation number of -2.
• Hydrogen is typically +1 when bonded to non-metals, but -1 when bonded to metals.
• Alkali metals (group 1 elements) have an oxidation number of +1, and alkaline earth metals (group 2 elements) have +2.

By using these rules, the oxidation states of elements in a compound can be determined if the compound's overall charge is known. Understanding oxidation states is crucial because they help identify electron transfer processes, helping to predict changes in chemical reactions.