Question

Which of the following are oxidation-reduction reactions? Explain. a. \(\mathrm{PCl}_{3}+\mathrm{Cl}_{2} \rightarrow \mathrm{PCl}_{5}\) b. \(\mathrm{Cu}+2 \mathrm{AgNO}_{3} \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{Ag}\) c. \(\mathrm{CO}_{2}+2 \mathrm{LiOH} \rightarrow \mathrm{Li}_{2} \mathrm{CO}_{3}+\mathrm{H}_{2} \mathrm{O}\) d. \(\mathrm{FeCl}_{2}+2 \mathrm{NaOH} \rightarrow \mathrm{Fe}(\mathrm{OH})_{2}+2 \mathrm{NaCl}\) e. \(\mathrm{MnO}_{2}+4 \mathrm{HCl} \rightarrow \mathrm{Cl}_{2}+2 \mathrm{H}_{2} \mathrm{O}+\mathrm{MnCl}_{2}\)

Short answer

Reactions (a), (b), and (e) are oxidation-reduction reactions, as there is a transfer of electrons, indicated by the change in oxidation numbers of the elements involved. Reactions (c) and (d) are not redox reactions as there is no change in oxidation numbers for any elements.

Step-by-step solution

Reaction (a)

: \(PCl_3 + Cl_2 \rightarrow PCl_5 \\\) Assign oxidation numbers to the elements in reactants and products: \( P: +3 \rightarrow +5 \\ Cl: -1 \rightarrow -1 \\ Cl_2: 0 \rightarrow -1\) The oxidation number of P increases from +3 to +5, so it is oxidized. Cl in \(Cl_2\) decreases from 0 to -1, so it is reduced. Since both oxidation and reduction occur, reaction (a) is a redox reaction.

Reaction (b)

: \(Cu + 2AgNO_3 \rightarrow Cu(NO_3)_2 + 2Ag\) Assign oxidation numbers to the elements in reactants and products: \(Cu: 0 \rightarrow +2 \\ Ag: +1 \rightarrow 0\) The oxidation number of Cu increases from 0 to +2, so it is oxidized. The oxidation number of Ag decreases from +1 to 0, so it is reduced. Since both oxidation and reduction occur, reaction (b) is a redox reaction.

Reaction (c)

: \(CO_2 + 2LiOH \rightarrow Li_2CO_3 + H_2O\) Assign oxidation numbers to the elements in reactants and products: \(C: +4 \rightarrow +4 \\ O: -2 \rightarrow -2 \\ Li: +1 \rightarrow +1 \\ H: +1 \rightarrow +1\) There is no change in oxidation numbers for any element. Therefore, reaction (c) is not a redox reaction.

Reaction (d)

:: \(FeCl_2 + 2NaOH \rightarrow Fe(OH)_2 + 2NaCl\) Assign oxidation numbers to the elements in reactants and products: \(Fe: +2 \rightarrow +2 \\ Cl: -1 \rightarrow -1 \\ Na: +1 \rightarrow +1 \\ H: +1 \rightarrow +1 \\ O: -2 \rightarrow -2\) There is no change in oxidation numbers for any element. Therefore, reaction (d) is not a redox reaction.

Reaction (e)

:: \(MnO_2 + 4HCl \rightarrow Cl_2 + 2H_2O + MnCl_2\) Assign oxidation numbers to the elements in reactants and products: \(Mn: +4 \rightarrow +2 \\ O: -2 \rightarrow -2 \\ H: +1 \rightarrow +1 \\ Cl: -1 \rightarrow 0\) The oxidation number of Mn decreases from +4 to +2, so it is reduced. The oxidation number of Cl increases from -1 to 0, so it is oxidized. Since both oxidation and reduction occur, reaction (e) is a redox reaction. In conclusion, reactions (a), (b), and (e) are oxidation-reduction reactions while reactions (c) and (d) are not.

Key concepts

Oxidation Numbers

Understanding oxidation numbers is crucial for identifying redox reactions and balancing chemical equations. Oxidation numbers, often referred to as oxidation states, indicate the degree of oxidation of an atom in a chemical compound. Simplistically, it's a hypothetical charge that an atom would have if all bonds to atoms of different elements were completely ionic.

For example, in the reaction \(PCl_3 + Cl_2 \rightarrow PCl_5\), phosphorus (P) changes its oxidation state from +3 in PCl_3 to +5 in PCl_5. Chlorine (Cl) in PCl_3 maintains an oxidation state of -1, while Cl_2, being diatomic, starts with an oxidation number of 0 before gaining electrons to become -1 in PCl_5. This shift in oxidation numbers confirms that a transfer of electrons has occurred, classifying this as a redox reaction.

To determine oxidation numbers, certain rules are followed: Elements in their elemental form have an oxidation number of 0, for monoatomic ions, the oxidation number is the same as the charge, and for molecules, the more electronegative element is assigned a negative oxidation number while the less electronegative element is assigned a positive one. Sum of oxidation numbers in a neutral compound is always zero; for polyatomic ions, the sum equals the ion's charge.

Redox Reaction Identification

Identification of redox reactions hinges on the concept of oxidation numbers. A redox (short for oxidation-reduction) reaction occurs when there is a transfer of electrons between atoms, resulting in changes to their oxidation states. Look at the reaction \(Cu + 2AgNO_3 \rightarrow Cu(NO_3)_2 + 2Ag\). Here, copper (Cu) starts off with an oxidation number of 0 and ends up with +2, while silver (Ag) goes from +1 to 0. This exchange characterizes a redox reaction since Cu loses electrons (oxidation) and Ag gains electrons (reduction).

Redox reactions are everywhere; they occur in batteries, during photosynthesis, and are even responsible for the rusting of iron. To identify a redox reaction, one must look for a change in oxidation numbers of the elements involved. If at least one element is oxidized while another is reduced, a redox reaction has taken place. It's also vital to note that oxidation and reduction always occur together; if an atom loses electrons, another must gain them.

Chemical Equation Balancing

Balancing chemical equations is a fundamental skill in chemistry, ensuring the adherence to the law of conservation of mass. To balance an equation, one must ensure that the same number of atoms for each element is present on both sides of the reaction. Consider the redox reaction \(MnO_2 + 4HCl \rightarrow Cl_2 + 2H_2O + MnCl_2\). In the initial equation, oxygen and hydrogen were not balanced. By adjusting the coefficients, we made sure that there are four hydrogens and two oxygens on both sides.

To balance a redox reaction, it's helpful to break it into half-reactions, one for oxidation and one for reduction. Then, balance the mass and charge of each half-reaction by adding H₂O, H⁺, and electrons as needed. After balancing both half-reactions, they can be combined to give the balanced overall reaction. This method is particularly valuable for complex redox reactions involving multiple elements undergoing changes in oxidation numbers.