Problem 15 - Chemistry: Principles and Reactions

Question

Write a balanced net ionic equation for (a) the electrolytic decomposition of hydrogen fluoride. (b) the oxidation of iodide ion to iodine by hydrogen peroxide in acidic solution. Hydrogen peroxide is reduced to water.

Step-by-Step Solution

Answer

Question: Write balanced net ionic equations for the following reactions: (a) electrolytic decomposition of hydrogen fluoride (HF), and (b) oxidation of iodide ion (I-) to iodine (I2) by hydrogen peroxide (H2O2) in an acidic solution. Answer: a) 2HF → H2 + 2F2 b) 2I- + H2O2 + 2H+ → I2 + 2H2O

Step 1: a) Electrolytic decomposition of hydrogen fluoride

Step 1: Write the overall reaction First, we know that HF undergoes an electrolytic decomposition, which means that an electric current is used to break it down into its elements: hydrogen (H2) and fluoride (F2). HF → H2 + F2 Step 2: Write the half-reactions To write the net ionic equation, we need to identify the half-reactions involved. The oxidation half-reaction is: 2HF → H2 + 2F- + 2e- The reduction half-reaction is: 2F- + 2e- → F2 Step 3: Balance the half-reactions for mass and charge In this case, both half-reactions are already balanced for both mass and charge. If they were not balanced, we would need to adjust the coefficients accordingly. Step 4: Combine the half-reactions Now, we can combine the half-reactions to form the balanced net ionic equation: 2HF → H2 + F2 + 2F- + 2e- 2F- + 2e- → F2 ----------------------------------- Net ionic equation: 2HF → H2 + 2F2

Step 2: b) Oxidation of iodide ion to iodine by hydrogen peroxide in acidic solution

Step 1: Write the overall reaction We know that iodide ion (I-) is oxidized to iodine (I2) by hydrogen peroxide (H2O2) in an acidic solution. 2I- + H2O2 → I2 + 2H2O Step 2: Write the half-reactions To write the net ionic equation, we need to identify the half-reactions involved. The oxidation half-reaction is: 2I- → I2 + 2e- The reduction half-reaction is: H2O2 + 2H+ + 2e- → 2H2O Step 3: Balance the half-reactions for mass and charge The oxidation half-reaction is already balanced, but the reduction half-reaction needs to be balanced for both mass and charge. H2O2 + 2H+ + 2e- → 2H2O Now both half-reactions are balanced. Step 4: Combine the half-reactions Now, we can combine the half-reactions to form the balanced net ionic equation: 2I- → I2 + 2e- H2O2 + 2H+ + 2e- → 2H2O ------------------------------------ Net ionic equation: 2I- + H2O2 + 2H+ → I2 + 2H2O

Key Concepts

Electrolytic Decomposition

Electrolytic decomposition is a process where an electrical current is utilized to break down a compound. This is a fascinating concept because it involves transforming substances through electricity, specifically breaking chemical bonds to form simpler substances. For example, in the decomposition of hydrogen fluoride (HF), electrical energy is harnessed to separate it into hydrogen gas (H₂) and fluorine gas (F₂). This reaction can be split into two half-reactions, which are crucial for balancing the net ionic equation.

In the oxidation half-reaction, HF loses electrons:

• 2HF → H₂ + 2F⁻ + 2e⁻

In the reduction half-reaction, fluoride ions gain electrons:

• 2F⁻ + 2e⁻ → F₂

When these half-reactions are combined and balanced, they yield the overall net ionic equation:

• 2HF → H₂ + 2F₂

This shows us how substances interact electrochemically and how conservation of charge and mass is maintained.

Oxidation-Reduction Reactions

Oxidation-reduction reactions, or redox reactions, are a cornerstone of chemical processes. They involve the transfer of electrons between species, indicating changes in oxidation states. This interplay can either manifest through gaining or losing electrons, hence the names oxidation and reduction.

Understanding this concept can be simplified by examining the oxidation of iodide ions (I⁻) by hydrogen peroxide (H₂O₂) in an acidic solution. Here, the iodide ion loses electrons to form iodine (I₂):

• The oxidation half-reaction: 2I⁻ → I₂ + 2e⁻

Simultaneously, hydrogen peroxide acts as an oxidizing agent and gains electrons to form water:

• The reduction half-reaction: H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O

The art of balancing redox reactions lies in ensuring that the number of electrons lost and gained match correctly, leading to the net ionic equation:

• 2I⁻ + H₂O₂ + 2H⁺ → I₂ + 2H₂O

Through redox reactions, we observe chemical changes that are fundamental to energy transformation and various biological processes.

Acidic Solution

An acidic solution is one where there is an excess of hydrogen ions (H⁺). This environmental condition can greatly influence the behavior of chemical reactions, including redox reactions.

In the context of oxidation of iodide ions by hydrogen peroxide, the acidity of the solution is essential for the progression of the reaction. The presence of extra H⁺ ions, which are prominent in acidic mediums, aids in the overall balancing of the reaction:

• The reduction half-reaction already highlights the involvement of 2H⁺: H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O

The acidic environment ensures that the produced ions and molecules have the necessary H⁺ to interact and form the desired products. It can also stabilize intermediates or assist in breaking existing bonds. This attribute is crucial for fine-tuning reactions to achieve desired outcomes, such as in industrial synthesis or metabolic pathways.