Question

Indicate whether the following balanced equations involve oxidation-reduction. If they do, identify the elements that undergo changes in oxidation number. (a) \(\mathrm{PBr}_{3}(l)+3 \mathrm{H}_{2} \mathrm{O}(l)--\rightarrow \mathrm{H}_{3} \mathrm{PO}_{3}(a q)+3 \mathrm{HBr}(a q)\) (b) \(\mathrm{NaI}(a q)+3 \mathrm{HOCl}(a q)-\cdots \rightarrow \mathrm{NaIO}_{3}(a q)+3 \mathrm{HCl}(a q)\) (c) \(3 \mathrm{SO}_{2}(g)+2 \mathrm{HNO}_{3}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)--\rightarrow\) $$ 3 \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{NO}(g) $$ (d) \(2 \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{NaBr}(s)--\rightarrow\) \(\mathrm{Br}_{2}(l)+\mathrm{SO}_{2}(g)+\mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)\)

Short answer

(a) Not a redox reaction - no changes in oxidation numbers. (b) Redox reaction - iodine is oxidized, chlorine is reduced. (c) Redox reaction - sulfur is oxidized, nitrogen is reduced. (d) Redox reaction - sulfur is reduced, bromine is oxidized.

Step-by-step solution

(a) PBr3 + 3 H2O → H3PO3 + 3 HBr

First, assign oxidation numbers to each element in the reactants and products. The general rules for assigning oxidation numbers claim that hydrogen has an oxidation number of +1 in its compounds and oxygen has an oxidation number of -2 in its compounds (exceptions do exist). Other elements' oxidation numbers can be deduced from there. Assigning oxidation numbers: P: +3 (in PBr3) --> +3 (in H3PO3) Br: -1 (in PBr3) --> -1 (in HBr) H: +1 (in H2O and HBr) O: -2 (in H2O and H3PO3) Since there are no changes in the oxidation numbers of any elements, the given reaction is not a redox reaction.

(b) NaI + 3 HOCl → NaIO3 + 3 HCl

Assign oxidation numbers to each element in the reactants and products: Na: +1 (in NaI and NaIO3) I: -1 (in NaI) --> +5 (in NaIO3) O: -2 (in HOCl and NaIO3) H: +1 (in HOCl and HCl) Cl: +1 (in HOCl) --> -1 (in HCl) Iodine's oxidation number changes from -1 to +5 while chlorine's changes from +1 to -1. Since there are changes in oxidation numbers, the given reaction is a redox reaction. Iodine is oxidized, and chlorine is reduced.

(c) 3 SO2 + 2 HNO3 + 2 H2O → 3 H2SO4 + 2 NO

Assign oxidation numbers to each element in the reactants and products: S: +4 (in SO2) --> +6 (in H2SO4) O: -2 (in all compounds) H: +1 (in HNO3, H2O, and H2SO4) N: +5 (in HNO3) --> +2 (in NO) Sulfur's oxidation number changes from +4 to +6 while nitrogen's changes from +5 to +2. Since there are changes in oxidation numbers, the given reaction is a redox reaction. Sulfur is oxidized, and nitrogen is reduced.

(d) 2 H2SO4 + 2 NaBr → Br2 + SO2 + Na2SO4 + 2 H2O

Assign oxidation numbers to each element in the reactants and products: H: +1 (in H2SO4 and H2O) S: +6 (in H2SO4) --> +4 (in SO2) and +6 (in Na2SO4) O: -2 (in all compounds) Na: +1 (in NaBr and Na2SO4) Br: -1 (in NaBr) --> 0 (in Br2) Sulfur's oxidation number changes from +6 to +4, and bromine's changes from -1 to 0. Since there are changes in oxidation numbers, the given reaction is a redox reaction. Sulfur is reduced and bromine is oxidized.

Key concepts

Oxidation Numbers

To understand oxidation-reduction reactions, it is essential to comprehend the concept of oxidation numbers. An oxidation number, sometimes called an oxidation state, is a hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. Oxidation numbers help track how many electrons are lost or gained during a chemical reaction.

Here are some fundamental rules to assign oxidation numbers:

• Elements in their elemental form always have an oxidation number of zero. For example, in \( ext{O}_2\) or \( ext{N}_2\), the oxidation number is 0.
• For monoatomic ions, the oxidation number is equal to the ion's charge. For instance, for \( ext{Na}^+\), the oxidation number is +1.
• Hydrogen typically has an oxidation number of +1, and oxygen typically has an oxidation number of -2 in their compounds, although exceptions occur.
• In a compound or a neutral molecule, the sum of oxidation numbers equals zero, while in polyatomic ions, the sum equals the ion's charge.

Grasping these basics will aid in identifying whether a reaction involves a redox process, as evident shifts in oxidation numbers usually hint at electron transfers between atoms.

Redox Reaction Identification

Oxidation-reduction, or redox reactions, play a crucial role in chemistry as these processes involve the transfer of electrons between chemical species. Identifying a redox reaction involves checking if there's a change in oxidation states of atoms involved.

A redox reaction consists of two main processes:

• Oxidation: The process where an atom, ion, or molecule loses electrons, resulting in an increase in oxidation number. For example, in reaction (b) from the problem, iodine's oxidation number changes from -1 (in NaI) to +5 (in NaIO3), indicating oxidation.
• Reduction: Conversely, reduction involves the gain of electrons, resulting in a decrease in oxidation number. In the same example, chlorine undergoes a reduction, as its oxidation number diminishes from +1 (in HOCl) to -1 (in HCl).

By evaluating the changes in oxidation numbers, one can determine the atoms or elements undergoing oxidation and reduction, making it possible to classify the reaction as a redox process. Clearly identifying these shifts can reveal the intricate dances of electrons that drive chemical reactions.

Balancing Chemical Equations

Balancing chemical equations is a critical step in analyzing and performing redox reactions, ensuring that the Law of Conservation of Mass is upheld. This means the same number of each type of atom must be present on both sides of the equation.

When balancing redox reactions, a specific approach needs to be followed:

• First, break the redox reaction into two half-reactions, one for oxidation and one for reduction.
• Balance each half-reaction separately, both in terms of atoms and charges. This may involve adding water molecules, hydrogen ions ( H^+ or hydroxide ions ( OH^- ), and electrons as required to achieve balance.
• Ensure that the electrons lost in the oxidation half-reaction equal the electrons gained in the reduction half-reaction. It may be necessary to multiply each half-reaction by suitable coefficients to accomplish this.
• Finally, add up the balanced half-reactions to get the balanced overall reaction, confirming that masses and charges are balanced.

Understanding the role of electrons and making sure both mass and charge balance are crucial for correctly interpreting and predicting the outcomes of redox reactions. Balancing equations fundamentally provides insight into the stoichiometric relationships within a reaction, illuminating how matter rearranges during chemical transformations.