Redox Reactions and Oxidation Number

Redox Reactions In Terms Of Electron Transfer Reactions

Redox Reactions: Redox reactions (reduction-oxidation reactions) are chemical reactions characterized by the transfer of electrons between reacting species. They involve a simultaneous increase and decrease in oxidation states.

Competitive Electron Transfer Reactions

Concept: Many redox reactions can be understood as a competition for electrons. One species has a greater tendency to lose electrons (is a better reducing agent), while another species has a greater tendency to gain electrons (is a better oxidizing agent).

Half-Reactions: Redox reactions are often represented as two separate half-reactions: an oxidation half-reaction and a reduction half-reaction. These half-reactions illustrate the transfer of electrons.

Example: Reaction between Zinc metal and Copper(II) sulfate solution:

$$Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$$

The overall ionic equation is:

$$Zn(s) + Cu^{2+}(aq) + SO_4^{2-}(aq) \rightarrow Zn^{2+}(aq) + SO_4^{2-}(aq) + Cu(s)$$

The spectator ion ($SO_4^{2-}$) can be removed to show the net ionic equation:

$$Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$$

This can be split into two half-reactions:

Oxidation Half-Reaction: $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$ (Zinc loses electrons, its oxidation state increases from 0 to +2)
Reduction Half-Reaction: $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$ (Copper(II) ion gains electrons, its oxidation state decreases from +2 to 0)

In this reaction, Zinc is the reducing agent (it gets oxidized and causes the reduction of copper ions), and Copper(II) ions are the oxidizing agent (they get reduced and cause the oxidation of zinc).

Reactivity Series and Competitive Electron Transfer: The concept of competitive electron transfer is directly related to the reactivity series of metals. A more reactive metal (higher in the series) has a greater tendency to lose electrons (is a stronger reducing agent) than a less reactive metal.

Example: If a piece of Zinc metal is placed in a Copper(II) sulfate solution, Zinc will displace Copper from the solution because Zinc is more reactive than Copper. Zinc atoms lose electrons to Copper(II) ions:

$$Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$$

However, if a piece of Copper metal is placed in a Zinc sulfate solution, no reaction will occur because Copper is less reactive than Zinc; it does not have a strong enough tendency to lose electrons to displace Zinc ions.

Balancing Redox Reactions: Balancing redox reactions often involves using these half-reactions and ensuring that the number of electrons lost in the oxidation half-reaction equals the number of electrons gained in the reduction half-reaction.

Oxidation Number

Oxidation Number (Oxidation State): The oxidation number of an element in a compound is a numerical value assigned to each atom on the basis of a set of rules. It represents the hypothetical charge an atom would have if all bonds to atoms of different elements were $100\%$ ionic, with electrons in each bond assigned to the more electronegative atom.

Rules for Assigning Oxidation Numbers:

Elements in Elemental Form: The oxidation number of an atom in its elemental form is zero. (e.g., $Na$, $O_2$, $P_4$, $S_8$, $Fe$, $Cl_2$).
Monatomic Ions: The oxidation number of a monatomic ion is equal to its charge. (e.g., $Na^+$ is +1, $Mg^{2+}$ is +2, $Cl^-$ is -1, $O^{2-}$ is -2).
Oxygen: Oxygen usually has an oxidation number of -2 in its compounds. Exceptions:
- In peroxides (e.g., $H_2O_2$), oxygen is -1.
- In superoxides (e.g., $KO_2$), oxygen is -1/2.
- When bonded to fluorine (e.g., $OF_2$), oxygen is +2.
Hydrogen: Hydrogen usually has an oxidation number of +1 when bonded to non-metals. Exceptions:
- In metal hydrides (e.g., $NaH$, $CaH_2$), hydrogen is -1.
Fluorine: Fluorine always has an oxidation number of -1 in its compounds.
Other Halogens: Other halogens (Cl, Br, I) usually have an oxidation number of -1, except when bonded to oxygen or a more electronegative halogen.
Sum of Oxidation Numbers:
- In a neutral molecule, the sum of the oxidation numbers of all atoms must be zero.
- In a polyatomic ion, the sum of the oxidation numbers of all atoms must equal the charge of the ion.

Calculating Oxidation Numbers:

Example 1: In $SO_2$, what is the oxidation number of Sulfur?

Oxygen is -2. There are two oxygen atoms, so their total oxidation state is $2 \times (-2) = -4$. The molecule is neutral, so the oxidation state of Sulfur (S) must be +4.

Example 2: In $KMnO_4$, what is the oxidation number of Manganese?

Potassium (Group 1) is +1. Oxygen is -2. Total charge of $O_4$ is $4 \times (-2) = -8$. The molecule is neutral, so (+1) + (Oxidation state of Mn) + (-8) = 0. Oxidation state of Mn = +7.

Example 3: In $SO_4^{2-}$, what is the oxidation number of Sulfur?

Oxygen is -2. Total charge of $O_4$ is $4 \times (-2) = -8$. The ion has a charge of -2. So, (Oxidation state of S) + (-8) = -2. Oxidation state of S = +6.

Limitations Of Concept Of Oxidation Number

While the concept of oxidation number is incredibly useful for identifying redox reactions and balancing them, it has certain limitations:

Hypothetical Nature: Oxidation numbers are based on the assumption of ionic bonding, which is an oversimplification. Many bonds have significant covalent character, and electron distribution is not perfectly assigned to the more electronegative atom.
Not the Actual Charge: The calculated oxidation number does not always represent the true charge on an atom within a molecule or ion. For example, in $HCl$, Chlorine is assigned -1 and Hydrogen +1, but the bond is polar covalent, not purely ionic.
Ambiguity in Some Cases: For certain elements in specific compounds or molecules with complex bonding, assigning a unique oxidation number can be ambiguous or may not accurately reflect the chemical reality (e.g., oxidation states in organometallic compounds).
Fractional Oxidation States: In some compounds containing equivalent atoms where electrons might be delocalized or shared in complex ways (e.g., $Br_3O_8$), fractional oxidation states can arise, which are mathematical averages and not representative of individual atom charges.
Does not fully explain Bonding: It does not provide a complete picture of the bonding or electron distribution within molecules, which is better described by theories like Valence Bond Theory or Molecular Orbital Theory.

Despite these limitations, the concept of oxidation number remains a powerful tool for understanding and manipulating redox reactions.