Problem 114 - General Chemistry: Principles and Modern Applications

Question

Which of the following reactions are oxidationreduction reactions? (a) \(\mathrm{H}_{2} \mathrm{CO}_{3}(\mathrm{aq}) \longrightarrow \mathrm{H}_{2} \mathrm{O}(1)+\mathrm{CO}_{2}(\mathrm{g})\) (b) \(2 \mathrm{Li}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(1) \longrightarrow 2 \mathrm{LiOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\) (c) \(4 \mathrm{Ag}(\mathrm{s})+\mathrm{PtCl}_{4}(\mathrm{aq}) \longrightarrow 4 \mathrm{AgCl}(\mathrm{s})+\mathrm{Pt}(\mathrm{s})\) (d) \(2 \mathrm{HClO}_{4}(\mathrm{aq})+\mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{aq}) \longrightarrow\) \(2 \mathrm{H}_{2} \mathrm{O}(1)+\mathrm{Ca}\left(\mathrm{ClO}_{4}\right)_{2}(\mathrm{aq})\)

Step-by-Step Solution

Answer

The given redox reactions are: \n(b) \(2 Li(s) + 2 H_2O(l) \rightarrow 2 LiOH(aq) + H_2(g) \) \n(c) \(4 Ag(s) + PtCl_{4}(aq) \rightarrow 4 AgCl(s) + Pt(s) \)

Step 1: Determine Oxidation Numbers

The first step in identifying a redox reaction is to figure out the oxidation numbers of all the atoms in the reactants and products. Atom or ion oxidation numbers will represent its degree of oxidation. It's important to know the basic oxidation number rules: free elements have an oxidation number of 0, alkali metals in compounds have an oxidation number of +1 and alkaline earth metals in compounds have an oxidation number of +2.

Step 2: Analyze the first given reaction

\( H_2CO_3(aq) \rightarrow H_2O(l) + CO_2(g) \) \nIn \( H_2CO_3 \), the oxidation number of oxygen is -2 and for hydrogen is +1, therefore the carbon is +4 to balance the equation. In the products, the oxidation numbers remain the same, showing that no redox process has occurred. Hence this equation is not a redox reaction.

Step 3: Analyze the second given reaction

\(2 Li(s) + 2 H_2O(l) \rightarrow 2 LiOH(aq) + H_2(g) \) \nIn this reaction, the Li atom is oxidized from an oxidation state of 0 to +1, and the hydrogen in the water molecule is reduced from +1 to 0. Therefore, this equation is a redox reaction.

Step 4: Analyze the third given reaction

\(4 Ag(s) + PtCl_{4}(aq) \rightarrow 4 AgCl(s) + Pt(s) \) \nHere, the silver Ag is oxidized from 0 to +1 and platinum Pt is reduced from +2 to 0. Therefore, this equation is a redox reaction.

Step 5: Analyze the fourth given reaction

\(2 HClO_{4}(aq) + Ca(OH)_{2}(aq) \rightarrow 2 H_2O(l) + Ca(ClO_{4})_{2} \rightarrow (aq) \) \n In \( HClO_4 \), the oxidation number of chlorine is +7. In \( Ca(ClO_{4})_{2} \), the number for chlorine remains +7. Thus, there's no change in the oxidation state, indicating that this reaction is not a redox reaction.

Key Concepts

Identifying Redox Reactions

Oxidation-reduction, or 'redox', reactions are fundamental concepts in chemistry where the oxidation state of atoms changes through the transfer of electrons. To identify redox reactions, one must look for changes in the oxidation state of atoms in the reactants compared to the products.

For instance, in the reaction between lithium and water, lithium atoms go from a neutral state to a positive charge, losing electrons, which is oxidation. Simultaneously, hydrogen gains these electrons, going from a positive state to zero charge in the formed hydrogen gas, indicating reduction. This exchange confirms the reaction as a redox process. In contrast, if there's no change in oxidation states as seen with the reaction of carbonic acid decomposing into water and carbon dioxide, we classify it as a non-redox reaction.

Determining Oxidation Numbers

Oxidation numbers or states are assigned to elements in chemical compounds to keep track of electron transfers during chemical reactions. These numbers are based on a set of rules, such as: pure elements have an oxidation state of zero, oxidation state of alkali metals in compounds is always +1, and for alkaline earth metals, it's +2, among others.

To determine oxidation numbers, we often use these rules along with the fact that the sum of oxidation numbers in a neutral compound must equal zero. In ionic compounds, they must equal the charge of the ion. By understanding these principles, one can methodically assign oxidation states to each element in a reaction and assess whether a redox process occurs.

Chemical Reactions in Chemistry

In the diverse world of chemical reactions, classifying and understanding the nature of each reaction is essential. Chemical reactions in chemistry involve rearranging atoms to form new substances and can be exothermic or endothermic, based on energy release or absorption.

These reactions are characterized by reaction types such as synthesis, decomposition, single replacement, double replacement, and of course, redox reactions. Each type follows certain patterns and rules that help chemists predict products and comprehend the intricacies of the interactions. Redox reactions, in particular, have immense implications in energy production, biological processes, and industrial applications.

Redox Process in Chemical Equations

Incorporating the redox process into chemical equations involves identifying the substances oxidized and reduced, as well as their corresponding oxidation numbers. Analyzing the changes in these numbers from reactants to products allows us to balance the redox equation.

For complete comprehension, balancing redox equations also consists of ensuring that the number of electrons lost and gained is equal, often incorporating methods like the half-reaction or ion-electron technique for complex equations. By carefully analyzing and balancing redox equations, one can detail the electron transfers and work through predicting the products of chemical reactions.