Question

(a) Write the Lewis structure of \(\mathrm{H}_{2} \mathrm{O}_{2}\) and predict the approximate \(\mathrm{H}-\mathrm{O}-\mathrm{O}\) bond angle. Which of the following ions will be oxidized by hydrogen peraxide in acidic solution: \((b\rangle \mathrm{Cu}^{+}\); (c) \(\mathrm{Mn}^{2+}\); (d) \(\mathrm{Ag}^{+}\); (e) \(\mathrm{F}^{-?}\)

Short answer

The Lewis structure for H2O2 has an O-O single bond with each oxygen atom also bound to a hydrogen atom and having two lone pairs. The H-O-O bond angle is approximately 90-120 degrees. H2O2 can oxidize \(\text{Cu}^+\), \(\text{Mn}^{2+}\), and \(\text{Ag}^+\), but not \(\text{F}^-\) in acidic solution.

Step-by-step solution

Draft the Lewis Structure for H2O2

To draw the Lewis structure for H2O2, first determine the total number of valence electrons available. Oxygen (O) has 6 valence electrons and Hydrogen (H) has 1. There are 2 oxygens and 2 hydrogens, which gives a total of 14 valence electrons (2*6 + 2*1). Arrange these electrons to form bonds between atoms: O-H single bonds and an O-O single bond. Place the remaining electrons to satisfy the octet rule for oxygen, resulting in two lone pairs on each oxygen.

Predict the H-O-O Bond Angle

Using VSEPR theory and considering the repulsive forces between the electron pairs, predict the bond angles in H2O2. The molecular geometry around each oxygen in H2O2 is bent or V-shaped due to the two lone pairs of electrons. Therefore, the approximate H-O-O bond angle is less than the tetrahedral angle (109.5 degrees) and is in the range of 90-120 degrees.

Identify the Oxidization of Specific Ions by H2O2

H2O2 acts as an oxidizing agent in acidic solution. We compare the standard electrode potentials for the half-reactions of the listed ions to the H2O2 half-reaction. Ions with a lower reduction potential compared to H2O2 will be oxidized by it. Thus, look up the standard reduction potentials for \(\text{Cu}^+\), \(\text{Mn}^{2+}\), \(\text{Ag}^+\), and \(\text{F}^-\) and compare.

Conclusion on Oxidation Capability

Without specific standard reduction potential values provided, we use general knowledge that \(\text{F}^-\) is not easily oxidized due to its high electronegativity and its position in the electrochemical series, thus excluding \(\text{F}^-\). Typically, \(\text{Cu}^+\), \(\text{Mn}^{2+}\), and \(\text{Ag}^+\) can be oxidized by H2O2, particularly when in acidic solutions.

Key concepts

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory is a model used in chemistry to predict the shape of individual molecules based upon the extent of electron-pair electrostatic repulsion. It is based on the principle that electron pairs around a central atom will organize themselves in a three-dimensional space to minimize the repulsive forces between them.

For example, in the case of hydrogen peroxide (H_{2}O_{2}), each oxygen atom has two lone pairs and forms two single bonds (one with hydrogen and the other with the other oxygen atom), leading to a bent molecular geometry. The repulsion between the lone pairs on the oxygen atoms causes the H-O-O bond angle to be less than the typical 109.5 degrees found in a tetrahedral arrangement. This results in the observed bond angle of about 90-120 degrees. Understanding the VSEPR theory is vital for predicting and explaining the shapes and bond angles in molecules.

Oxidizing Agents in Chemistry

Oxidizing agents are substances that can accept electrons during a chemical reaction, thus facilitating the oxidation of another substance. In the context of hydrogen peroxide (H_{2}O_{2}) acting as an oxidizing agent, it means that H_{2}O_{2} can gain electrons from other substances, effectively oxidizing them.

An important aspect of understanding an oxidizing agent is recognizing its ability to be reduced, as it will cause another substance to lose electrons. Substances that are easily oxidized typically have fewer electrons and a lower electrochemical potential. When analyzing whether a substance can act as an oxidizing agent, we must compare its standard electrode potential to that of other potential reactants. For instance, in the case of H_{2}O_{2}, to know which ions it can oxidize, we compare its standard electrode potential to the potentials of those ions. If the ion's potential is lower, it indicates that the ion can be oxidized by H_{2}O_{2}.

Standard Electrode Potentials

Standard Electrode Potentials, often referred to as standard reduction potentials, are intrinsic properties of elements (ions/molecules) that measure the tendency of a chemical species to acquire electrons and thereby get reduced. These potentials are measured against the standard hydrogen electrode, which is set at 0 volts by convention.

The electrode potential is essential when determining the direction and spontaneity of a redox reaction. During a redox process, the species with the higher electrode potential gets reduced, while the one with the lower potential gets oxidized. For example, comparing the standard electrode potentials of Cu^+, Mn^{2+}, Ag^+, and F^- to that of H_{2}O_{2} helps us to predict which species will be oxidized by H_{2}O_{2} in an acidic solution. The ones with electrode potentials lower than that of H_{2}O_{2} would be susceptible to oxidation.