Question

Many oxidation-reduction reactions can be balanced by inspection. Try to balance the following reactions by inspection. In each reaction, identify the substance reduced and the substance oxidized. a. $\mathrm{Al}(s)+\mathrm{HCl}(aq)\to {\mathrm{AlCl}}_3(aq)+{\mathrm{H}}_2(g)$ b. ${\mathrm{CH}}_4(g)+\mathrm{S}(s)\to {\mathrm{CS}}_2(l)+{\mathrm{H}}_2\mathrm{S}(\mathrm{g})$ c. ${\mathrm{C}}_3{\mathrm{H}}_8(g)+{\mathrm{O}}_2(g)\to {\mathrm{CO}}_2(g)+{\mathrm{H}}_2\mathrm{O}(l)$ d. $\mathrm{Cu}(s)+{\mathrm{Ag}}^{+}(aq)\to \mathrm{Ag}(s)+{\mathrm{Cu}}^{2+}(aq)$

Short answer

The balanced reactions and oxidized/reduced substances are as follows: a. $2\mathrm{Al}(s)+6\mathrm{HCl}(aq)\to 2{\mathrm{AlCl}}_3(aq)+3{\mathrm{H}}_2(g)$; Oxidized: $\mathrm{Al}$; Reduced: $\mathrm{HCl}$ b. ${\mathrm{CH}}_4(g)+2\mathrm{S}(s)\to {\mathrm{CS}}_2(l)+2{\mathrm{H}}_2\mathrm{S}(g)$; Oxidized: ${\mathrm{CH}}_4$; Reduced: $\mathrm{S}$ c. ${\mathrm{C}}_3{\mathrm{H}}_8(g)+5{\mathrm{O}}_2(g)\to 3{\mathrm{CO}}_2(g)+4{\mathrm{H}}_2\mathrm{O}(l)$; Oxidized: ${\mathrm{C}}_3{\mathrm{H}}_8$; Reduced: ${\mathrm{O}}_2$ d. $\mathrm{Cu}(s)+{\mathrm{Ag}}^{+}(aq)\to \mathrm{Ag}(s)+{\mathrm{Cu}}^{2+}(aq)$; Oxidized: $\mathrm{Cu}$; Reduced: ${\mathrm{Ag}}^{+}$

Step-by-step solution

Balance the equation

We need to make sure that the number of atoms of each element on the reactants side is equal to the number of atoms of the same element on the products side. - For chlorine, there are three in the product ${\mathrm{AlCl}}_3$, so we will need three $\mathrm{HCl}$ molecules on the reactant side: $2\mathrm{Al}(s)+6\mathrm{HCl}(aq)\to 2{\mathrm{AlCl}}_3(aq)+3{\mathrm{H}}_2(g)$

Identify the oxidized and reduced substances

The substance oxidized is $\mathrm{Al}$, since it loses electrons, and the substance reduced is $\mathrm{HCl}$, since it gains electrons. b. ${\mathrm{CH}}_4(g)+\mathrm{S}(s)\to {\mathrm{CS}}_2(l)+{\mathrm{H}}_2\mathrm{S}(\mathrm{g})$

Balance the equation

Again, make sure the number of atoms of each element is equal on both sides. - For carbon, sulfur, and hydrogen, we have: ${\mathrm{CH}}_4(g)+2\mathrm{S}(s)\to {\mathrm{CS}}_2(l)+2{\mathrm{H}}_2\mathrm{S}(g)$

Identify the oxidized and reduced substances

The substance oxidized is ${\mathrm{CH}}_4$, as carbon loses electrons, and the substance reduced is sulfur, since it gains electrons. c. ${\mathrm{C}}_3{\mathrm{H}}_8(g)+{\mathrm{O}}_2(g)\to {\mathrm{CO}}_2(g)+{\mathrm{H}}_2\mathrm{O}(l)$

Balance the equation

Verify that the number of atoms of each element is equal on both sides. - For carbon, hydrogen, and oxygen, we have: ${\mathrm{C}}_3{\mathrm{H}}_8(g)+5{\mathrm{O}}_2(g)\to 3{\mathrm{CO}}_2(g)+4{\mathrm{H}}_2\mathrm{O}(l)$

Identify the oxidized and reduced substances

The substance oxidized is ${\mathrm{C}}_3{\mathrm{H}}_8$, as it loses electrons, and the substance reduced is ${\mathrm{O}}_2$, since it gains electrons. d. $\mathrm{Cu}(s)+{\mathrm{Ag}}^{+}(aq)\to \mathrm{Ag}(s)+{\mathrm{Cu}}^{2+}(aq)$

Balance the equation

Check that the number of atoms of each element is equal on both sides. - For copper and silver, no balancing is needed because there is one atom of each element on both sides: $\mathrm{Cu}(s)+{\mathrm{Ag}}^{+}(aq)\to \mathrm{Ag}(s)+{\mathrm{Cu}}^{2+}(aq)$

Identify the oxidized and reduced substances

The substance oxidized is Cu, as it loses electrons, and the substance reduced is ${\mathrm{Ag}}^{+}$, since it gains electrons.

Key concepts

Oxidation and Reduction

Oxidation and reduction are two key concepts of chemical reactions, especially in the context of redox reactions. In simple terms, oxidation refers to the loss of electrons by a substance, while reduction is the gain of electrons. A convenient way to remember this is through the acronym 'LEO the lion says GER': Loss of Electrons is Oxidation, Gain of Electrons is Reduction.

Every redox reaction involves a transfer of electrons from one substance to another. Oxidation cannot occur without reduction happening simultaneously; the reactant that gets oxidized ends up losing electrons, which are then gained by the reactant that gets reduced. This electron transfer is what drives the chemical changes in redox reactions.

Oftentimes, identifying which substances are oxidized and which are reduced can be determined by tracking the changes in oxidation states of the elements participating in the reaction. Oxidation states indicate the hypothetical charge an atom would have if all bonds were ionic. An increase in oxidation state denotes oxidation, while a decrease indicates reduction.

Oxidizing and Reducing Agents

In a redox reaction, not only are substances oxidized and reduced, they usually act as agents that facilitate these processes for each other. Oxidizing agents are substances that oxidize other substances by gaining electrons themselves. As they receive electrons, they become reduced. Conversely, reducing agents are substances that reduce other substances by losing electrons and thereby getting oxidized themselves.

For instance, in the reaction of aluminum with hydrochloric acid, $\mathrm{Al}(s)+\mathrm{HCl}(aq)\to {\mathrm{AlCl}}_3(aq)+{\mathrm{H}}_2(g)$, the \mathrm{HCl} acts as the oxidizing agent because it accepts electrons to become hydrogen gas (\mathrm{H}_2). In contrast, aluminum (\mathrm{Al}) is the reducing agent since it donates electrons to form \mathrm{AlCl}_3. Identifying these agents is crucial because it reflects the role each reactant plays within the redox process.

Chemical Reaction Balancing

Balancing chemical reactions is a fundamental skill in chemistry, ensuring the law of conservation of mass is obeyed. During a reaction, the number of atoms of each element must remain constant from reactants to products. However, in redox reactions, it's not just about balancing the atoms but also ensuring the exchange of electrons between the oxidizing and reducing agents is accounted for.

To balance a redox reaction, one must first write the correct formulas for all reactants and products, then adjust coefficients to get the same number of atoms of each element on both sides of the equation. Sometimes, additional steps involving the separation of the reaction into half-reactions (oxidation half and reduction half) and balancing the charges with electrons are required, particularly in complex reactions.

For example, in the combustion of propane, ${\mathrm{C}}_3{\mathrm{H}}_8(g)+{\mathrm{O}}_2(g)\to {\mathrm{CO}}_2(g)+{\mathrm{H}}_2\mathrm{O}(l)$, coefficients need to be adjusted to ensure that each element is balanced and that electrons lost and gained are equal.

Electron Transfer in Redox

The essence of redox reactions is the transfer of electrons from one species to another. This movement of electrons is what distinguishes redox reactions from other types of chemical reactions. In any given redox reaction, the substance being oxidized loses a certain number of electrons, and the substance being reduced gains the same number of electrons.

Understanding the flow of electrons can help us deduce the direction of the reaction and the energy changes involved. In cases like the displacement reaction between copper and silver ions, $\mathrm{Cu}(s)+{\mathrm{Ag}}^{+}(aq)\to \mathrm{Ag}(s)+{\mathrm{Cu}}^{2+}(aq)$, we can see that copper atoms lose electrons (get oxidized), which are then acquired by the silver ions (which get reduced). Tracking electron transfer is not only pivotal for balancing redox reactions but also for understanding the broader implications like energy generation in batteries and corrosion processes.