Question

The element bromine was first prepared by oxidation of aqueous potassium bromide with solid manganese(IV) oxide. Write a balanced net ionic equation for the reaction in aqueous acidic solution. \(\left(\mathrm{Mn}^{2+}\right.\) is also formed.)

Short answer

The balanced net ionic equation is: \( \text{MnO}_2 + 4\text{H}^+ + 2\text{Br}^- \rightarrow \text{Mn}^{2+} + \text{Br}_2 + 2\text{H}_2\text{O} \).

Step-by-step solution

Write the Chemical Formulas of Reactants and Products

First, identify the reactants and products. The reactants are aqueous potassium bromide (KBr) and solid manganese(IV) oxide (MnO_2). The acidic solution provides hydrogen ions (H^+). The products include bromine (Br_2) and manganese(II) ions (Mn^{2+}).

Separate into Half-Reactions

Write the two half-reactions. The oxidation half-reaction involves the conversion of bromide ions (Br^-) to bromine (Br_2): \[ 2\text{Br}^- \rightarrow \text{Br}_2 + 2\text{e}^- \]The reduction half-reaction involves the conversion of MnO_2 to Mn^{2+} ions: \[ \text{MnO}_2 + 4\text{H}^+ + 2\text{e}^- \rightarrow \text{Mn}^{2+} + 2\text{H}_2\text{O} \]

Balance the Electrons

Ensure the number of electrons lost in the oxidation half-reaction equals the number gained in the reduction half-reaction. In this case, 2 electrons are transferred in each half-reaction, so they are already balanced.

Combine the Half-Reactions

Add the balanced half-reactions together and cancel out the electrons to get the net ionic equation:\[ \text{MnO}_2 + 4\text{H}^+ + 2\text{Br}^- \rightarrow \text{Mn}^{2+} + \text{Br}_2 + 2\text{H}_2\text{O} \]

Verify the Equation is Balanced

Check to ensure that the number of atoms and charges are balanced. Each side of the equation has 1 manganese, 2 bromine, 4 hydrogen, and 2 oxygen atoms and both sides have a net charge of +2.

Key concepts

Understanding Oxidation-Reduction Reactions

Oxidation-reduction reactions (also called redox reactions) are chemical reactions where the oxidation state of atoms changes. This process involves two key components: oxidation and reduction.
When a substance undergoes oxidation, it loses electrons. Conversely, during reduction, a substance gains electrons. These changes occur simultaneously in redox reactions.

• Oxidation: Loss of electrons leading to an increase in oxidation state.
• Reduction: Gain of electrons resulting in a decrease in oxidation state.

In the provided exercise, bromide ions are oxidized to bromine, while manganese dioxide is reduced to manganese ions. Understanding these fundamental principles will aid in determining the components of redox reactions effectively.

Breaking Down Half-Reactions

Half-reactions offer a way to represent oxidation and reduction processes individually. They help emphasize how electrons are transferred during the chemical reactions.
For every redox reaction, there will be two corresponding half-reactions:

• One representing the oxidation process, showing the loss of electrons.
• Another showcasing the reduction process, displaying the gain of electrons.

By writing separate half-reactions, we can easily examine the electron flow and balance the equation. In our current example:

• The oxidation half-reaction: \(2\text{Br}^- \rightarrow \text{Br}_2 + 2\text{e}^-\)
• The reduction half-reaction: \(\text{MnO}_2 + 4\text{H}^+ + 2\text{e}^- \rightarrow \text{Mn}^{2+} + 2\text{H}_2\text{O}\)

Observing half-reactions allows for clarity in balancing electron transfer, especially in complex reactions.

Role of an Acidic Solution

An acidic solution provides hydrogen ions \(\text{H}^+\) that are essential during certain redox reactions. These ions can influence the balance of the chemical equation.
In our exercise, the acidic environment is crucial as it supplies the needed \(\text{H}^+\) ions for the reaction:

• They participate in the reduction half-reaction, interacting with manganese dioxide to form water.

The presence of \(\text{H}^+\) not only helps balance the atoms but can also impact the stability and conditions under which the reaction occurs. When balancing reactions in acidic solutions, always ensure the \(\text{H}^+\) from the solution is accounted for in your net ionic equation.

Balancing Equations with Confidence

Balancing chemical equations, especially in redox reactions, is essential to accurately represent what transpires in the chemical process. This includes both the conservation of mass and charge.
Here's how you can approach balancing:

• Balance Mass: Ensure all atoms of each element appear equally on both sides of the equation.
• Balance Charge: Ensure the total ionic charge is the same on both sides of the equation.

In the given exercise, balancing is achieved as both sides have equal numbers of each type of atom and a net charge of +2.
Remember to verify each step carefully, ensuring that no discrepancies exist in the atoms or charge to maintain the law of conservation in chemical equations. This verification is crucial for writing correct net ionic equations, especially in redox scenarios.