Question

Indicate whether the following balanced equations involve oxidation-reduction. If they do, identify the elements that undergo changes in oxidation number. $$ \begin{array}{l}{\text { (a) } \mathrm{PBr}_{3}(l)+3 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{3} \mathrm{PO}_{3}(a q)+3 \mathrm{HBr}(a q)} \\ {\text { (b) } \mathrm{NaI}(a q)+3 \mathrm{HNOl}(a q) \longrightarrow \mathrm{NaIO}_{3}(a q)+3 \mathrm{HCl}(a q)} \\ {\text { (c) } 3 \mathrm{SO}_{2}(g)+2 \mathrm{HNO}_{3}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow} \\ {\quad 3 \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{NO}(g)}\end{array} $$

Short answer

(a) No oxidation-reduction occurs as there is no change in oxidation state for any element. (b) Oxidation-reduction occurs. Iodine (I) is oxidized from -1 to +5, and Nitrogen (N) is reduced from +5 to +3. (c) Oxidation-reduction occurs. Sulfur (S) is oxidized from +4 to +6, and Nitrogen (N) is reduced from +5 to +2.

Step-by-step solution

(a) PBr3(l) + 3 H2O(l) -> H3PO3(aq) + 3 HBr(aq)

First, assign oxidation numbers to all elements. PBr3: P: +3 (as it is combined with three bromine atoms, each with -1 oxidation state) Br: -1 H2O: H: +1 O: -2 H3PO3: H: +1 P: +3 (as it is combined with three oxygen atoms, each with -2 oxidation state, and three hydrogen atoms, each with +1 oxidation state) O: -2 HBr: H: +1 Br: -1 Now, let's compare the oxidation states: P: +3 in PBr3 and +3 in H3PO3 (No change) Br: -1 in PBr3 and -1 in HBr (No change) H: +1 in H2O and +1 in HBr and H3PO3 (No change) O: -2 in H2O and -2 in H3PO3 (No change) As there is no change in oxidation state for any element, this reaction does not involve oxidation-reduction.

(b) NaI(aq) + 3 HNO3(aq) -> NaIO3(aq) + 3 HNO2(aq)

First, assign oxidation numbers to all elements. NaI: Na: +1 I: -1 HNO3: H: +1 N: +5 (as it is combined with three oxygen atoms, each with -2 oxidation state, and one hydrogen atom with +1 oxidation state) O: -2 NaIO3: Na: +1 I: +5 (as it is combined with three oxygen atoms, each with -2 oxidation state) O: -2 HNO2: H: +1 N: +3 (as it is combined with two oxygen atoms, each with -2 oxidation state, and one hydrogen atom with +1 oxidation state) O: -2 Now, let's compare the oxidation states: Na: +1 in NaI and +1 in NaIO3 (No change) I: -1 in NaI and +5 in NaIO3 (Change) H: +1 in HNO3 and +1 in HNO2 (No change) N: +5 in HNO3 and +3 in HNO2 (Change) O: -2 in HNO3 and -2 in NaIO3 and HNO2 (No change) As there is a change in the oxidation state for I and N, this reaction involves oxidation-reduction. Iodine (I) is oxidized from -1 to +5, and Nitrogen (N) is reduced from +5 to +3 in the process.

(c) 3 SO2(g) + 2 HNO3(aq) + 2 H2O(l) -> 3 H2SO4(aq) + 2 NO(g)

First, assign oxidation numbers to all elements. SO2: S: +4 (as it is combined with two oxygen atoms, each with -2 oxidation state) O: -2 HNO3: H: +1 N: +5 (as it is combined with three oxygen atoms, each with -2 oxidation state, and one hydrogen atom with +1 oxidation state) O: -2 H2O: H: +1 O: -2 H2SO4: H: +1 S: +6 (as it is combined with four oxygen atoms, each with -2 oxidation state, and two hydrogen atoms, each with +1 oxidation state) O: -2 NO: N: +2 (as it is combined with one oxygen atom, with -2 oxidation state) O: -2 Now, let's compare the oxidation states: S: +4 in SO2 and +6 in H2SO4 (Change) O: -2 in SO2, HNO3, H2O, H2SO4, and NO (No change) H: +1 in HNO3, H2O, and H2SO4 (No change) N: +5 in HNO3 and +2 in NO (Change) As there is a change in the oxidation state for S and N, this reaction involves oxidation-reduction. Sulfur (S) is oxidized from +4 to +6 in the process, and Nitrogen (N) is reduced from +5 to +2.

Key concepts

Oxidation Number

The oxidation number, also known as oxidation state, is a concept used to keep track of how many electrons an atom uses or has available during chemical reactions. It often relates to the balance of electron sharing in molecules. In atoms, electrons are either lost, gained, or shared during reactions. This transfer or sharing allows us to assign each atom an oxidation number. Here's a simple way to determine oxidation numbers for common elements:

• For a pure element like O₂, the oxidation number is 0.
• For monoatomic ions, the oxidation number is the charge of the ion. For example, Na⁺ is +1.
• In compounds, hydrogen is generally +1, and oxygen is usually -2.

By examining the changes in these numbers, we can identify if oxidation or reduction occurs, which are key processes in redox reactions.

Oxidation State

The oxidation state is a useful indicator in determining an atom's behavior in a chemical reaction, particularly in redox reactions. Essentially, it represents the effective charge of an atom when electrons are counted according to certain rules. Assigning oxidation states can help identify which elements undergo changes in a chemical reaction. To figure out oxidation states in a reaction:

• For each atom in a molecule, consider how electrons are distributed. Elements in their most stable form, like diatomic oxygen, have oxidation states of 0.
• Typically, more electronegative elements have negative oxidation states.
• The sum of oxidation states for all atoms in a molecule or compound is equal to the overall charge of that molecule or compound.

Understanding oxidation states ensures we can pinpoint the electron transfer activity in redox reactions. This makes it easier to determine whether a substance is oxidized or reduced.

Chemical Equations

Chemical equations represent the substances involved in a reaction, displaying reactants transformed into products. These equations must always be balanced to respect the conservation of mass and charge. In redox reactions, balanced equations are crucial for identifying changes in oxidation states. Here's how to approach balancing these equations:

• Assign oxidation numbers to each element in the equation to identify what is oxidized and what is reduced.
• Ensure that the number of atoms for each element is the same on both sides of the equation.
• Check that the algebraic sum of oxidation numbers equals zero for neutral substances or matches the net charge for ionic compounds.

Balancing chemical equations helps us accurately represent chemical reactions and the conservation laws that govern them. This is essential to reliably analyze and understand redox processes.

Redox Reactions

Redox reactions, or oxidation-reduction reactions, are chemical processes where electrons are transferred between species. This involves two complementary processes: oxidation (loss of electrons) and reduction (gain of electrons). Understanding redox reactions revolves around identifying changes in oxidation states of elements within a reaction. Here's an approach to identifying redox processes:

• Look for changes in oxidation numbers to spot which elements have undergone oxidation and which have undergone reduction.
• Typically, an increase in oxidation number indicates oxidation, while a decrease suggests reduction.
• Ensure that the overall number of electrons lost equals the number of electrons gained, keeping the reaction balanced.

Redox reactions are part of many everyday processes, including combustion, respiration, and more. Mastery of these reactions opens up insight into a vast array of chemical, biological, and industrial applications.