Problem 12 - Advanced Problems in Physical Chemistry for Competitive Examinations

Question

A metal having negative reduction potential, when dipped in the solution of its own ions, has a tendency to (a) remain as metal atoms (b) become electrically positive (c) become electrically negative (d) be deposited from the solution

Step-by-Step Solution

Answer

A metal with negative reduction potential tends to lose electrons and become electrically positive (b).

Step 1: Understanding Reduction Potential

Reduction potential indicates the tendency of a species to gain electrons and get reduced. A negative reduction potential means that the species prefers to lose electrons and get oxidized rather than gaining electrons and getting reduced.

Step 2: Determining the Behavior of the Metal

As the metal has a negative reduction potential, it does not tend to gain electrons and reduce to metal atoms; rather, it will tend to lose electrons and oxidize. In this process, metal atoms are more likely to go into the solution as ions instead of coming out of the solution as atoms.

Step 3: Analyzing the Options

Given the negative reduction potential, the metal will not tend to remain as metal atoms (a) since it has a higher tendency to oxidize. When a metal loses electrons, it becomes positively charged since it loses negatively charged electrons (b). It does not become negatively charged (c) because that would require a gain of electrons, which is contrary to what happens with a negative reduction potential. As the metal prefers to lose electrons, it will not be readily deposited from the solution as atoms (d).

Key Concepts

Oxidation and Reduction

Understanding the interplay between oxidation and reduction is fundamental in chemistry, particularly when dealing with reactions that involve the transfer of electrons. These two processes are collectively known as 'redox reactions' and form the basis of many chemical processes, including electrochemical reactions.

In the simplest terms, oxidation is the loss of electrons, while reduction is the gain of electrons. It's essential to remember the mnemonic 'OIL RIG' – 'Oxidation Is Loss, Reduction Is Gain.' When a substance undergoes oxidation, its oxidation state increases, and it often becomes more positively charged because electrons are negatively charged. Conversely, when it undergoes reduction, its oxidation state decreases, leading to a more negatively charged species due to the gain of electrons.

This concept can be a bit abstract, so let's consider an example. In the exercise, a metal with a negative reduction potential indicates a propensity to lose electrons – that is, to undergo oxidation. Therefore, when this metal is placed in a solution with its ions, it's more inclined to release electrons, enhancing the ionic form rather than turning into solid metal atoms.

Understanding redox reactions also helps us predict the direction of a reaction and the stability of the products formed. This knowledge is integral to many areas of chemistry, including energy storage, corrosion, and biochemical processes.

Electrochemical Series

The electrochemical series, also known as the activity series, is a list of elements sorted by their standard reduction potentials. This list allows chemists to predict the outcome of redox reactions and determine which metals are more likely to oxidize or reduce in specific conditions.

Metals with higher reduction potentials are more likely to undergo reduction, making them good oxidizing agents. In contrast, metals with lower (or more negative) reduction potentials are more apt to lose electrons and therefore act as strong reducing agents.

• Metals at the top of the series, such as fluoride, have high tendencies to gain electrons and usually do not release them easily.
• Metals at the bottom, such as lithium, are strong electron donors and have a strong tendency to undergo oxidation.

The exercise under review involves a metal with a negative reduction potential, indicating it's likely found towards the bottom of the electrochemical series. It's prone to oxidation when in contact with its ionic solution, reinforcing its likelihood to lose electrons rather than gain them.

The series also helps us understand the feasibility of galvanic cells, which harness chemical reactions to generate electrical energy. Metals' placement on the series informs us about the direction of electron flow and the potential of the cell.

Galvanic Cells

Galvanic cells, or voltaic cells, are devices that convert chemical energy into electrical energy through spontaneous redox reactions. Two different metals, which serve as electrodes, are placed in their respective salt solutions and connected by a wire and a salt bridge.

One metal acts as the anode, where oxidation occurs, and the other as the cathode, where reduction happens. Electrons flow from the anode to the cathode through the external circuit, generating an electric current. The flow of ions in the internal circuit, through the salt bridge, helps maintain charge balance.

In the context of the exercise, a metal with a negative reduction potential would serve as the anode in a galvanic cell because it tends to lose electrons. When connected to a metal with a higher reduction potential, an electron flow would be established from the metal with negative potential (oxidation) to the one with positive potential (reduction).

The functionality of galvanic cells is rooted in the electrochemical series. This application allows us to harness the spontaneous reaction between two metals to produce electrical power. For example, these principles are applied in batteries, where galvanic cells supply power to a multitude of devices. Fully grasping the role of galvanic cells in electrochemistry helps students understand pivotal elements of energy conversion and storage.