Identifying Oxidized And Reduced Reactants In A Single-displacement Reaction

Identifying Oxidized and Reduced Reactants in a Single-Displacement Reaction

Single-displacement reactions, also known as single replacement reactions, are a fundamental type of chemical reaction where one element replaces another element in a compound. Understanding these reactions requires a grasp of oxidation and reduction, concepts central to redox chemistry. This article will delve deep into identifying the oxidized and reduced reactants in single-displacement reactions, providing a comprehensive guide with numerous examples and explanations to solidify your understanding.

Understanding Oxidation and Reduction

Before we dive into single-displacement reactions, let's establish a firm foundation in the core concepts of oxidation and reduction. These terms, often shortened to "redox," are intimately linked: one cannot occur without the other.

Oxidation

Oxidation, in its simplest form, is the loss of electrons by an atom, ion, or molecule. When an atom is oxidized, its oxidation state (or oxidation number) increases. This means it becomes more positive (or less negative). Remember the mnemonic OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

Reduction

Reduction is the gain of electrons by an atom, ion, or molecule. When an atom is reduced, its oxidation state decreases. This means it becomes more negative (or less positive).

Oxidation States

Assigning oxidation states is crucial for identifying oxidized and reduced species. Here are some general rules:

• Free elements: The oxidation state of an element in its free (uncombined) state is always 0. For example, the oxidation state of Na in Na(s) is 0.
• Monatomic ions: The oxidation state of a monatomic ion is equal to its charge. For example, the oxidation state of Na⁺ is +1, and the oxidation state of Cl⁻ is -1.
• Hydrogen: Usually +1, except in metal hydrides (e.g., NaH) where it is -1.
• Oxygen: Usually -2, except in peroxides (e.g., H₂O₂) where it is -1, and in superoxides (e.g., KO₂) where it is -1/2.
• Group 1 elements (alkali metals): Always +1.
• Group 2 elements (alkaline earth metals): Always +2.
• Fluorine: Always -1.
• The sum of oxidation states in a neutral compound: Must equal zero.
• The sum of oxidation states in a polyatomic ion: Must equal the charge of the ion.

Single-Displacement Reactions: A Detailed Look

Single-displacement reactions follow a general pattern:

A + BC → AC + B

where A is a more reactive element than B, displacing it from the compound BC. This displacement is always accompanied by a redox reaction. Let's examine this in detail:

• A is oxidized: A loses electrons and its oxidation state increases. It becomes a part of a new compound (AC).
• B is reduced: B gains electrons and its oxidation state decreases. It is released as a free element.

Let's illustrate this with several examples:

Example 1: Reaction of Zinc with Hydrochloric Acid

Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

• Zn (Zinc): Starts with an oxidation state of 0 (free element). It loses two electrons to become Zn²⁺ (oxidation state +2) in ZnCl₂. Therefore, zinc is oxidized.
• H (Hydrogen): Starts with an oxidation state of +1 in HCl. It gains one electron each to become H₂ (oxidation state 0). Therefore, hydrogen is reduced.

In this reaction, zinc is the reducing agent (it causes the reduction of hydrogen) and hydrogen ions (from HCl) are the oxidizing agent (they cause the oxidation of zinc).

Example 2: Reaction of Copper with Silver Nitrate

Cu(s) + 2AgNO₃(aq) → Cu(NO₃)₂(aq) + 2Ag(s)

• Cu (Copper): Starts with an oxidation state of 0. It loses two electrons to become Cu²⁺ (oxidation state +2) in Cu(NO₃)₂. Therefore, copper is oxidized.
• Ag (Silver): Starts with an oxidation state of +1 in AgNO₃. It gains one electron each to become Ag (oxidation state 0). Therefore, silver is reduced.

Copper is the reducing agent, and silver ions are the oxidizing agent.

Example 3: Reaction of Chlorine with Sodium Bromide

Cl₂(g) + 2NaBr(aq) → 2NaCl(aq) + Br₂(l)

• Cl₂ (Chlorine): Starts with an oxidation state of 0. Each chlorine atom gains one electron to become Cl⁻ (oxidation state -1) in NaCl. Therefore, chlorine is reduced.
• Br (Bromine): Starts with an oxidation state of -1 in NaBr. Each bromine atom loses one electron to become Br₂ (oxidation state 0). Therefore, bromine is oxidized.

Chlorine is the oxidizing agent, and bromide ions are the reducing agent.

Example 4: More Complex Example - Reaction of Iron with Copper(II) Sulfate

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

This reaction appears simpler, but let's break down the oxidation states:

• Fe (Iron): Starts at 0 (oxidation state of elemental iron). It loses two electrons to become Fe²⁺ (oxidation state +2) in FeSO₄. Therefore, iron is oxidized.
• Cu (Copper): Starts at +2 (oxidation state of copper in CuSO₄). It gains two electrons to become Cu (oxidation state 0). Therefore, copper is reduced.

Iron is the reducing agent and copper(II) ions are the oxidizing agent.

Predicting Single-Displacement Reactions: The Activity Series

Predicting whether a single-displacement reaction will occur depends on the relative reactivity of the elements involved. This is summarized in the activity series (also known as the reactivity series), which lists metals in order of decreasing reactivity. A more reactive metal will displace a less reactive metal from its compound. Similarly, a more reactive non-metal will displace a less reactive non-metal.

For example, in the activity series, zinc is more reactive than hydrogen; therefore, zinc can displace hydrogen from acids. However, copper is less reactive than hydrogen, and thus it cannot displace hydrogen from acids.

The activity series is not absolute; reaction conditions (temperature, concentration) can influence reactivity.

Identifying Oxidized and Reduced Reactants: A Step-by-Step Approach

Here’s a systematic approach to identify oxidized and reduced reactants in a single-displacement reaction:

1. Write and balance the chemical equation: Ensure the equation is properly balanced to accurately track electron transfer.

2. Assign oxidation states: Assign oxidation states to all atoms in both reactants and products using the rules mentioned above.

3. Identify changes in oxidation states: Determine which atoms have changed their oxidation states. An increase in oxidation state indicates oxidation, while a decrease indicates reduction.

4. Identify the oxidizing and reducing agents: The species that gets reduced is the oxidizing agent (it causes the oxidation of another species), and the species that gets oxidized is the reducing agent (it causes the reduction of another species).

Conclusion

Understanding single-displacement reactions requires a thorough grasp of redox chemistry. By systematically assigning oxidation states and analyzing changes in these states, you can confidently identify the oxidized and reduced reactants and determine the oxidizing and reducing agents. The activity series serves as a valuable tool for predicting the feasibility of such reactions. Practice with various examples is key to mastering this important aspect of chemical reactions. Remember to always double-check your work and utilize the provided rules and examples to reinforce your learning and understanding. Through consistent practice and application of these principles, you'll become proficient in identifying oxidized and reduced reactants in single-displacement reactions and gain a deeper understanding of redox chemistry.