Question

Which would you expect to be the stronger reducing agent, \(\mathrm{Cr}^{2+}\) or \(\mathrm{Fe}^{2+} ?\)

Short answer

We can compare the reduction potentials of the two ions to determine the stronger reducing agent. From a standard reduction potential table, we find that: Cr²⁺: \(E^o = -0.41\,\text{V}\) Fe²⁺: \(E^o = +0.77\,\text{V}\) Since the reduction potential of Cr²⁺ (\(-0.41\,\text{V}\)) is more negative than that of Fe²⁺ (\(+0.77\,\text{V}\)), Cr²⁺ is the stronger reducing agent as it can more readily donate electrons (get oxidized) in a redox reaction.

Step-by-step solution

Lookup the reduction potentials of Cr²⁺ and Fe²⁺

To compare the reduction potentials of Cr²⁺ and Fe²⁺, refer to a standard reduction potential table. Find the reduction half-reactions and their respective standard potential values, Eº. For Cr: \(\mathrm{Cr^{3+}(aq) + e^- \rightarrow Cr^{2+}(aq)}\); \(E^o = -0.41\,\text{V}\) For Fe: \(\mathrm{Fe^{3+}(aq) + e^- \rightarrow Fe^{2+}(aq)}\); \(E^o = +0.77\,\text{V}\)

Compare the reduction potentials

Now that we have the standard reduction potentials for both ions, we can compare them. A more negative reduction potential indicates a stronger reducing agent. Cr²⁺: \(E^o = -0.41\,\text{V}\) Fe²⁺: \(E^o = +0.77\,\text{V}\) Since the reduction potential of Cr²⁺ (\(-0.41\,\text{V}\)) is more negative than the reduction potential of Fe²⁺ (\(+0.77\,\text{V}\)), Cr²⁺ is the stronger reducing agent as it can more readily donate electrons (get oxidized) in a redox reaction.

Key concepts

Standard Reduction Potential

Understanding standard reduction potentials is crucial when it comes to predicting the direction of redox reactions and the strength of reducing and oxidizing agents. The standard reduction potential is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. This value is reported in volts and is determined under standard conditions, which include a temperature of 298 K, a 1 molar concentration for each ion participating in the reaction, and a pressure of 1 atmosphere for any gases that are involved.

Every chemical species has its own standard reduction potential, and by convention, the more positive the value, the greater the affinity for electrons and the stronger the oxidizing agent. Conversely, species with more negative standard reduction potentials are stronger reducing agents, as they have a greater tendency to donate electrons. When we compare different species, we can predict which one will act as the reducing agent and which as the oxidizer in a given reaction.

Redox Reactions

Redox reactions are a family of reactions that involve the transfer of electrons between species. In essence, one species will lose electrons and another will gain them. The species that loses electrons is said to be oxidized, while the species gaining electrons is reduced. These reactions are fundamental in fields ranging from electrochemistry to metabolic processes within living organisms.

Oxidation and Reduction

• Oxidation is the loss of electrons during the reaction.
• Reduction is the gain of electrons during the reaction.

A simple mnemonic to remember this is 'OIL RIG' – Oxidation Is Loss, Reduction Is Gain. In the context of redox reactions, a reducing agent donates electrons and thus is oxidized itself. Meanwhile, the oxidizing agent accepts electrons and is reduced.

Electron Transfer in Chemistry

The transaction of electrons is the cornerstone of many chemical reactions. Electron transfer is the process by which electrons are relocated from one atom, ion, or molecule to another. This transfer goes hand in hand with energy changes, as the movement of electrons can both release and store energy.

Chemical Bonds and Energy

In chemical reactions, bonds are broken and new ones are formed. This restructuring involves shifts in electron localization which can change the energy state of the species involved. For example, the release of energy in the form of light or heat during a reaction commonly results from these electron movements. Conversely, the storage of energy, like in the bonds of ATP molecules in biology, also relies on the transfer of electrons to more stable configurations. Electron transfer is indispensable to both the generation of electricity in batteries and the metabolic processes that sustain life.

Galvanic Cells

Galvanic cells, also known as voltaic cells, form the basis of batteries and are constructed to harness the energy from spontaneous redox reactions. They consist of two different metals connected by a salt bridge, with each metal in its own solution of a salt containing the same metal. The key to a galvanic cell's function is the different standard reduction potential of the two metals, which ultimately drives the flow of electrons from the metal with the lower reduction potential to the metal with the higher one.

As a result, the metal acting as the anode is oxidized (loses electrons), while the metal acting as the cathode is reduced (gains electrons), creating a current that can be conducted out of the cell and used to perform work. The standard reduction potentials of the metals used in galvanic cells provide the information necessary to predict the direction of electron flow and the cell's voltage, which is a measure of the energy difference per charge transfer.