Problem 9 - Chemistry for Allied Health

Question

For each equation, 1 ) identify the oxidation numbers of each element, 2) determine if it is a redox reaction or not, and for redox reactions, 3) identify the species being oxidized and the species being reduced, and 4) identify the oxidizing and reducing agents. a. \(2 \mathrm{KClO}_{3}(s) \rightarrow 2 \mathrm{KCl}(s)+3 \mathrm{O}_{2}(g)\) b. \(\mathrm{H}_{2}(g)+\mathrm{CuO}(s) \rightarrow \mathrm{Cu}(s)+\mathrm{H}_{2} \mathrm{O}(l)\) c. \(2 \mathrm{Al}(s)+3 \operatorname{Sn}\left(\mathrm{NO}_{3}\right)_{2}(a q) \rightarrow 2 \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(a q)+3 \mathrm{Sn}(s)\) d. \(2 \mathrm{HNO}_{3}(a q)+6 \mathrm{HI}(a q) \rightarrow 2 \mathrm{NO}(g)+3 \mathrm{I}_{2}(s)+4 \mathrm{H}_{2} \mathrm{O}(l)\) e. \(\mathrm{AgNO}_{3}(a q)+\mathrm{NaCl}(a q) \rightarrow \mathrm{AgCl}(s)+\mathrm{NaNO}_{3}(a q)\) f. \(2 \mathrm{FeCl}_{3}(a q)+\mathrm{H}_{2} \mathrm{~S}(g) \rightarrow 2 \mathrm{FeCl}_{2}(a q)+2 \mathrm{HCl}(a q)+\mathrm{S}(s)\)

Step-by-Step Solution

Answer

Reactions A, B, C, D, and F are redox reactions. A: Cl is reduced, O is oxidized; B: H is oxidized, Cu is reduced; C: Al is oxidized, Sn is reduced; D: I is oxidized, N is reduced; F: S is oxidized, Fe is reduced. Reaction E is not a redox reaction.

Step 1: Determine Oxidation Numbers, Reaction A

For the reaction \(2 \mathrm{KClO}_{3}(s) \rightarrow 2 \mathrm{KCl}(s)+3 \mathrm{O}_{2}(g)\):- **K** in both \(\mathrm{KClO}_{3}\) and \(\mathrm{KCl}\) has an oxidation number of +1.- **Cl** in \(\mathrm{KClO}_{3}\) is +5 and in \(\mathrm{KCl}\) it is -1.- **O** in \(\mathrm{KClO}_{3}\) is -2 and in \(\mathrm{O}_2\), it is 0.

Step 2: Identify if Reaction A is Redox

Since the oxidation state of Cl changes from +5 in \(\mathrm{KClO}_{3}\) to -1 in \(\mathrm{KCl}\) and that of O changes from -2 in \(\mathrm{KClO}_{3}\) to 0 in \(\mathrm{O}_2\), it is a redox reaction.

Step 3: Determine Oxidized and Reduced Species, Reaction A

- **Cl** is reduced because its oxidation number decreases from +5 to -1. - **O** is oxidized because its oxidation number increases from -2 to 0.

Step 4: Identify Reducing and Oxidizing Agents, Reaction A

- **Reducing Agent:** \(\mathrm{KClO}_{3}\) because it provides the oxygen that is oxidized.- **Oxidizing Agent:** \(\mathrm{KClO}_{3}\) because it accepts electrons during the reduction of Cl.

Step 5: Determine Oxidation Numbers, Reaction B

For the reaction \(\mathrm{H}_{2}(g)+\mathrm{CuO}(s) \rightarrow \mathrm{Cu}(s)+\mathrm{H}_{2} \mathrm{O}(l)\):- **H** in \(\mathrm{H}_{2}\) is 0 and in \(\mathrm{H}_{2}\mathrm{O}\) is +1.- **Cu** in \(\mathrm{CuO}\) is +2 and in \(\mathrm{Cu}\) is 0.- **O** in \(\mathrm{CuO}\) and \(\mathrm{H}_{2}\mathrm{O}\) is -2.

Step 6: Identify if Reaction B is Redox

This is a redox reaction because the oxidation number of H increases (is oxidized) and the oxidation number of Cu decreases (is reduced).

Step 7: Determine Oxidized and Reduced Species, Reaction B

- **H** is oxidized from 0 to +1. - **Cu** is reduced from +2 to 0.

Step 8: Identify Reducing and Oxidizing Agents, Reaction B

- **Reducing Agent:** \(\mathrm{H}_{2}\) because it gives electrons to \(\mathrm{CuO}\).- **Oxidizing Agent:** \(\mathrm{CuO}\) because it accepts electrons from \(\mathrm{H}_{2}\).

Step 9: Determine Oxidation Numbers, Reaction C

For the reaction \(2 \mathrm{Al}(s)+3 \operatorname{Sn}\left(\mathrm{NO}_{3}\right)_{2}(a q) \rightarrow 2 \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(a q)+3 \mathrm{Sn}(s)\):- **Al** is 0 in \(\mathrm{Al}\) and +3 in \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}\).- **Sn** is +2 in \(\operatorname{Sn}\left(\mathrm{NO}_{3}\right)_{2}\) and 0 in \(\mathrm{Sn}\).- **N** and **O** are unchanged.

Step 10: Identify if Reaction C is Redox

This is a redox reaction because the oxidation number of Al increases (is oxidized) and the oxidation number of Sn decreases (is reduced).

Step 11: Determine Oxidized and Reduced Species, Reaction C

- **Al** is oxidized from 0 to +3. - **Sn** is reduced from +2 to 0.

Step 12: Identify Reducing and Oxidizing Agents, Reaction C

- **Reducing Agent:** \(\mathrm{Al}\) because it provides electrons to \(\operatorname{Sn}\left(\mathrm{NO}_{3}\right)_{2}\).- **Oxidizing Agent:** \(\operatorname{Sn}\left(\mathrm{NO}_{3}\right)_{2}\) because it accepts electrons from \(\mathrm{Al}\).

Step 13: Determine Oxidation Numbers, Reaction D

For the reaction \(2 \mathrm{HNO}_{3}(a q)+6 \mathrm{HI}(a q) \rightarrow 2 \mathrm{NO}(g)+3 \mathrm{I}_{2}(s)+4 \mathrm{H}_{2} \mathrm{O}(l)\):- **N** in \(\mathrm{HNO}_{3}\) is +5 and in \(\mathrm{NO}\) is +2.- **I** in \(\mathrm{HI}\) is -1 and in \(\mathrm{I}_{2}\) is 0.- **H** remains +1 and **O** is -2.

Step 14: Identify if Reaction D is Redox

This is a redox reaction because the oxidation number of N decreases (is reduced) and the oxidation number of I increases (is oxidized).

Step 15: Determine Oxidized and Reduced Species, Reaction D

- **I** is oxidized from -1 to 0. - **N** is reduced from +5 to +2.

Step 16: Identify Reducing and Oxidizing Agents, Reaction D

- **Reducing Agent:** \(\mathrm{HI}\) because it donates electrons to \(\mathrm{HNO}_{3}\).- **Oxidizing Agent:** \(\mathrm{HNO}_{3}\) because it accepts electrons from \(\mathrm{HI}\).

Step 17: Determine Oxidation Numbers, Reaction E

For the reaction \(\mathrm{AgNO}_{3}(a q)+\mathrm{NaCl}(a q) \rightarrow \mathrm{AgCl}(s)+\mathrm{NaNO}_{3}(a q)\):- **Ag, Na, N, O, Cl** remain unchanged in oxidation numbers (Ag +1, Na +1, N +5, Cl -1, O -2).

Step 18: Identify if Reaction E is Redox

This is not a redox reaction because none of the elements change their oxidation states.

Step 19: Determine Oxidation Numbers, Reaction F

For the reaction \(2 \mathrm{FeCl}_{3}(a q)+\mathrm{H}_{2} \mathrm{~S}(g) \rightarrow 2 \mathrm{FeCl}_{2}(a q)+2 \mathrm{HCl}(a q)+\mathrm{S}(s)\):- **Fe** is +3 in \(\mathrm{FeCl}_{3}\) and +2 in \(\mathrm{FeCl}_{2}\).- **S** is -2 in \(\mathrm{H}_{2} \mathrm{~S}\) and 0 in \(\mathrm{S}\).- **Cl** and **H** remain unchanged.

Step 20: Identify if Reaction F is Redox

This is a redox reaction because the oxidation number of Fe decreases (is reduced) and the oxidation number of S increases (is oxidized).

Step 21: Determine Oxidized and Reduced Species, Reaction F

- **S** is oxidized from -2 to 0. - **Fe** is reduced from +3 to +2.

Step 22: Identify Reducing and Oxidizing Agents, Reaction F

- **Reducing Agent:** \(\mathrm{H}_{2} \mathrm{~S}\) because it donates electrons to \(\mathrm{FeCl}_{3}\).- **Oxidizing Agent:** \(\mathrm{FeCl}_{3}\) because it accepts electrons from \(\mathrm{H}_{2} \mathrm{~S}\).

Key Concepts

Oxidation Numbers

Understanding oxidation numbers is a crucial step in analyzing redox reactions. An oxidation number is a charge an element would have if electrons were transferred completely, rather than shared as they are in covalent bonds. Oxidation numbers help us track the electron transfer in a chemical reaction, crucial for identifying what is being oxidized and what is being reduced.

Here are some common rules to determine oxidation numbers:

• Pure elements have an oxidation number of 0. For instance, in \( ext{O}_2\) or \( ext{H}_2\), each oxygen and hydrogen atom is 0.
• Fluorine is always -1 in compounds, while alkali metals (like \( ext{K}\)) are typically +1.
• In most of its compounds, oxygen is -2, except in peroxides where it is -1.
• Hydrogen is usually +1 when bonded with non-metals, but -1 when bonded with metals.
• The sum of oxidation numbers in a neutral compound must be zero. For polyatomic ions, it equals the ion's charge.

Oxidation numbers change during redox reactions:

• An increase in oxidation number indicates oxidation occurred.
• A decrease in oxidation number indicates reduction.

Oxidizing Agent

In a redox reaction, the oxidizing agent plays a vital role. This is the species that accepts electrons. By accepting these electrons, the oxidizing agent itself gets reduced.

Here are some key points about oxidizing agents:

• Oxidizing agents cause another substance to lose electrons (or get oxidized).
• These agents often include nonmetals and oxygen-containing compounds, like \( ext{KClO}_3\) or \( ext{HNO}_3\).
• They are crucial in applications like combustion, bleaching, and disinfecting.

In chemical reactions, identifying the oxidizing agent is essential because it tells us which component of the reaction is undergoing a reduction. For example, in the reaction of hydrogen gas \( ext{H}_2\) with copper(II) oxide \( ext{CuO}\), the copper(II) oxide \( ext{CuO}\) acts as the oxidizing agent because it accepts electrons from hydrogen, allowing the hydrogen to be oxidized.

Reducing Agent

Contrary to an oxidizing agent, a reducing agent donates electrons in a redox reaction. This donation facilitates the reduction of another species. Hence, the reducing agent itself undergoes oxidation.

Essential facts about reducing agents include:

• Reducing agents donate electrons and reduce other substances.
• They tend to be metals or hydrides, such as hydrogen (\( ext{H}_2\)) or \( ext{Al}\) in reactions.
• These agents are used extensively in processes such as metal extraction and organic chemistry reductions.

For students, it's key to recognize that the reducing agent loses electrons, leading to an increase in its oxidation number. In the example with aluminum and tin nitrate (\( ext{Sn(NO}_3)_2\)), aluminum acts as a reducing agent as it loses electrons to tin, indicating aluminum's oxidation.

Oxidized Species

The oxidized species in a redox reaction is characterized by an increase in oxidation number. This species loses electrons, hence undergoes oxidation.

Important points to consider about oxidized species:

• They undergo an increase in oxidation state through electron loss.
• Often involve metals or elements transitioning from a lower to a higher oxidation state.
• Identifying the oxidized species helps in understanding the electron flow in reactions.

For instance, in the equation \(2 ext{FeCl}_3 + ext{H}_2 ext{S} ightarrow 2 ext{FeCl}_2 + 2 ext{HCl} + ext{S}\), sulfur from \( ext{H}_2 ext{S}\) experiences oxidation as its oxidation state changes from -2 to 0, meaning it's the oxidized species.

Reduced Species

In a redox reaction, the reduced species is the one that gains electrons and experiences a decrease in oxidation number, signifying reduction.

Essential insights into reduced species include:

• The reduced species gains electrons, thus its oxidation number decreases.
• It is often associated with transition metals or compounds shifting from higher to lower oxidation states.
• Correctly identifying the reduced species aids in visualizing how electrons move within a reaction.

For example, in the reaction \(2 ext{HNO}_3 + 6 ext{HI} ightarrow 2 ext{NO} + 3 ext{I}_2 + 4 ext{H}_2 ext{O}\), the nitrogen in \( ext{HNO}_3\) is the reduced species, decreasing in oxidation state from +5 to +2.