Identifying The Oxidized Substance In The Redox Reaction Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Understanding Redox Reactions

To accurately determine what substance is oxidized in the given redox reaction, it's vital to have a solid grasp of the core principles behind oxidation-reduction reactions, commonly known as redox reactions. Redox reactions are fundamental chemical processes where electrons are transferred between chemical species. This electron transfer leads to changes in the oxidation states of the participating atoms, ions, or molecules. At the heart of every redox reaction lie two simultaneous processes: oxidation and reduction. These processes are intrinsically linked, with one always occurring in tandem with the other.

Oxidation, in its essence, is the process where a substance loses electrons. This loss of electrons results in an increase in the oxidation state of the substance. The oxidation state, also known as the oxidation number, is a concept that assigns a charge to an atom in a chemical compound or ion, assuming that all bonds are ionic. It provides a way to track the electron distribution during chemical reactions. An increase in oxidation state signifies that the substance has become more positively charged, a direct consequence of losing negatively charged electrons. In chemical equations, the substance undergoing oxidation is termed the reducing agent because it donates electrons to another substance, thereby causing the reduction of that other substance. The reducing agent itself gets oxidized in the process.

On the flip side, reduction is the process where a substance gains electrons. This gain of electrons results in a decrease in the oxidation state of the substance, making it more negatively charged. The substance that gains electrons is known as the oxidizing agent. The oxidizing agent is responsible for accepting electrons from another substance, thereby causing the oxidation of that other substance. In the process of accepting electrons, the oxidizing agent itself gets reduced. To effectively identify oxidized and reduced substances, one must carefully examine the changes in oxidation states of the reactants and products in the chemical equation. The oxidation state rules provide a systematic way to assign oxidation numbers to atoms in molecules and ions, enabling the tracking of electron transfer during redox reactions. This understanding is crucial not only in chemistry but also in various fields such as biology, environmental science, and materials science, where redox reactions play pivotal roles.

Identifying Oxidation States

Identifying oxidation states is crucial for determining which substance is oxidized in a redox reaction. The oxidation state, also known as the oxidation number, represents the hypothetical charge an atom would have if all bonds were ionic. Understanding how to assign oxidation states is fundamental to grasping redox chemistry. Several rules govern the assignment of these states, ensuring a systematic approach to analyzing chemical reactions. These rules help in tracking the movement of electrons and identifying the species undergoing oxidation and reduction.

The first rule to consider is that the oxidation state of an element in its elemental form is always zero. This is because, in their elemental state, atoms have not yet formed bonds with other atoms and, therefore, have neither gained nor lost electrons. For example, the oxidation state of zinc (Zn) as a solid (Zn(s)) is 0, and the oxidation state of copper (Cu) as a solid (Cu(s)) is also 0. This baseline is essential for comparing changes in oxidation states as elements combine into compounds or ions.

Monatomic ions, which consist of a single atom with a charge, have an oxidation state equal to their charge. For instance, the oxidation state of the zinc ion ($Zn^{2+}$) is +2, reflecting its positive charge. Similarly, the oxidation state of the copper(II) ion ($Cu^{2+}$) is also +2. This rule is straightforward and provides a direct way to determine the oxidation state of ions in solution.

For neutral molecules, the sum of the oxidation states of all atoms must equal zero. This is because neutral molecules have no overall charge. To apply this rule, one must know the oxidation states of some atoms within the molecule. For example, oxygen typically has an oxidation state of -2 in compounds (except in peroxides, where it is -1, and when bonded to fluorine). Hydrogen usually has an oxidation state of +1 when bonded to nonmetals and -1 when bonded to metals. These guidelines allow us to deduce the oxidation states of other atoms in the molecule by ensuring the overall charge neutrality.

Polyatomic ions, which are ions composed of multiple atoms, have a total oxidation state equal to their overall charge. For example, in the sulfate ion ($SO_4^{2-}$), the sum of the oxidation states of sulfur and oxygen must equal -2. Knowing that oxygen typically has an oxidation state of -2, we can calculate the oxidation state of sulfur in this ion. This rule is particularly useful in complex chemical species where multiple elements are present, and the interplay of their oxidation states determines the ion’s charge.

By systematically applying these rules, chemists can accurately determine the oxidation states of atoms in various chemical species. This knowledge is essential for identifying which substances are oxidized (lose electrons, increase in oxidation state) and which are reduced (gain electrons, decrease in oxidation state) in redox reactions. Mastering these rules is a critical step in understanding and predicting the behavior of chemical reactions, particularly in fields such as electrochemistry, corrosion science, and organic chemistry.

Analyzing the Given Reaction

In the given redox reaction, $Zn(s) + Cu^{2+}(aq) ightarrow Zn^{2+}(aq) + Cu(s)$, to identify the substance that is oxidized, we need to meticulously examine the changes in oxidation states of each element involved. This careful analysis will reveal which species loses electrons (oxidation) and which gains electrons (reduction). By comparing the oxidation states of the reactants and products, we can pinpoint the substance that undergoes oxidation, thus answering the question at hand.

Starting with zinc (Zn), we observe that it begins as a solid, $Zn(s)$. As a pure element in its standard state, zinc has an oxidation state of 0. On the product side, zinc appears as a zinc ion, $Zn^{2+}(aq)$. The oxidation state of an ion is equivalent to its charge, so $Zn^{2+}$ has an oxidation state of +2. Comparing the oxidation states, zinc has changed from 0 to +2. This increase in oxidation state indicates that zinc has lost two electrons. According to the definition of oxidation, the loss of electrons signifies that zinc has been oxidized in this reaction. The process can be summarized as: $Zn(s) ightarrow Zn^{2+}(aq) + 2e^-$. This equation clearly shows the release of two electrons by the zinc atom during its transformation from the elemental state to the ionic state.

Next, let's consider copper (Cu). In the reactants, copper exists as a copper(II) ion, $Cu^{2+}(aq)$, which has an oxidation state of +2. On the product side, copper is present as solid copper, $Cu(s)$, which, as a pure element, has an oxidation state of 0. The change in oxidation state for copper is from +2 to 0, representing a decrease. This decrease in oxidation state means that copper has gained electrons. Specifically, the copper(II) ion gains two electrons to become solid copper. This process is reduction, and it can be represented by the equation: $Cu^{2+}(aq) + 2e^- ightarrow Cu(s)$. This equation illustrates the acceptance of two electrons by the copper(II) ion, leading to its reduction to elemental copper.

By systematically analyzing the changes in oxidation states, we can clearly see that zinc undergoes oxidation, while copper undergoes reduction. Zinc loses electrons and its oxidation state increases, fulfilling the definition of oxidation. Copper gains electrons, and its oxidation state decreases, fulfilling the definition of reduction. This comprehensive approach allows us to accurately identify the oxidized and reduced species in the redox reaction, which is a fundamental aspect of understanding chemical reactions and their mechanisms.

Conclusion

Based on our analysis of the redox reaction $Zn(s) + Cu^{2+}(aq) ightarrow Zn^{2+}(aq) + Cu(s)$, it is evident that zinc (Zn) is the substance oxidized. Zinc's oxidation state increases from 0 to +2, indicating a loss of electrons, which defines oxidation. Therefore, the correct answer is D. Zn.