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 \(\mathrm{FeSO}_{4}\), (e) \(\mathrm{Pt}\) in \(\mathrm{PtCl}_{4}\), (f) Cl in \(\mathrm{NaClO}_{4}\)

Short answer

(a) -2, (b) +4, (c) +3, (d) +2, (e) +4, (f) +7.

Step-by-step solution

Identify Oxidation Number Rules

The oxidation number is the charge an atom would have if the compound was composed of ions. Common rules include: the oxidation number of an element in its natural state is 0, and for a monatomic ion, it is equal to the ion charge. Hydrogen is generally +1, oxygen is -2, and in compounds, the sum of the oxidation numbers must equal the overall charge of the compound.

Determine Oxidation Number of N in N2H4

In hydrazine ( N_{2}H_{4} ), hydrogen's oxidation number is +1. With four hydrogens, the total positive charge is +4. The compound N_2H_4 is neutral, so the total oxidation number for nitrogen atoms is -4. Since there are two nitrogen atoms, each nitrogen has an oxidation number of -4/2 = -2.

Find Oxidation Number of N in NO2

In nitrogen dioxide ( NO_{2} ), oxygen has an oxidation number of -2. With two oxygens, the total is -4. The compound is neutral, so nitrogen must have an oxidation number of +4 to balance the -4 from oxygen.

Calculate Oxidation Number of Mn in MnCl3

Chlorine's oxidation number is -1. With three chlorine atoms, the total negative charge is -3. As the compound is neutral, manganese must have an oxidation number of +3 to balance the -3 from the chlorines.

Ascertain Oxidation Number of Fe in FeSO4

In iron(II) sulfate ( FeSO_{4} ), the sulfate ion ( SO_{4}^{2-} ) has an overall charge of -2. Since the compound is neutral, iron must have an oxidation number of +2 to balance the sulfate's -2.

Compute Oxidation Number of Pt in PtCl4

For platinum(IV) chloride ( PtCl_{4} ), each chlorine has an oxidation number of -1, totaling -4 for four chlorine atoms. The compound is neutral, so platinum has an oxidation number of +4 to offset the -4 from chlorine.

Evaluate Oxidation Number of Cl in NaClO4

In sodium perchlorate ( NaClO_{4} ), sodium's oxidation number is +1. Oxygen is -2, leading to -8 for four oxygens. The compound is neutral, and the oxidation state of chlorine must be +7 to balance the overall -7 from sodium and oxygen.

Key concepts

Chemical Compounds

Chemical compounds are substances formed when two or more elements combine through chemical bonds. These bonds can be covalent, ionic, or metallic, and they hold the atoms together in specific arrangements. Each chemical compound has a fixed composition, meaning that the ratio of each type of atom in the compound remains constant.
This constancy allows for predictable properties and behaviors.

Understanding chemical compounds is key in resolving oxidation number problems. They are represented by chemical formulas, such as \(\mathrm{H_2O}\) for water or \(\mathrm{NaCl}\) for table salt. These formulas convey the types and numbers of atoms involved.

• Covalent compounds: Atoms share electrons.
• Ionic compounds: Atoms transfer electrons, forming charged ions.
• Metallic compounds: Atoms share electron pools, offering conductivity and malleability.

Crucially, knowing whether a compound's bonds are ionic or covalent helps determine oxidation states for elements within the compound.

Oxidation State Rules

Oxidation state rules provide a system to assign charges to individual elements within chemical compounds, aiding in complex chemical equation balancing and analysis. These rules are systematic and should be applied in the following order:

• The oxidation state of a free element is always 0. For example, in \(\mathrm{N_2}\) or \(\mathrm{O_2}\), each atom has an oxidation number of 0.
• For simple ions, the oxidation state equals the ion's charge. For example, \(\mathrm{Na^+}\) has an oxidation state of +1.
• The oxidation state for oxygen is usually -2, and for hydrogen, it's usually +1.
• The oxidation state of a compound must sum to zero if the compound is neutral or to the charge of the ion.

By applying these rules, we can deduce the oxidation states of elements in various complex structures, such as figuring out that nitrogen in \(\mathrm{NO_2}\) has an oxidation state of +4 by balancing against the oxidation states of the oxygen atoms.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between two elements, affecting their oxidation states.
In these reactions, one species is oxidized (loses electrons) while the other is reduced (gains electrons).

These reactions are fundamental to many processes, including combustion, metabolism, and corrosion. They play a crucial role in biological and industrial applications.

• Oxidation: Increases oxidation state. Example: \(\mathrm{Fe^{2+} \rightarrow Fe^{3+}}\), losing an electron.
• Reduction: Decreases oxidation state. Example: \(\mathrm{Cu^{2+} + 2e^- \rightarrow Cu}\), gaining electrons.

Identifying which substance is oxidized and which is reduced involves determining changes in oxidation states. In a redox reaction equation, it's crucial to confirm that the total change in oxidation numbers is balanced, ensuring both mass and charge conservation. When tackling redox problems, remember that oxidizing agents gain electrons (get reduced) and reducing agents lose electrons (get oxidized).