chapter 4 lec 2

Stoichiometry of Chemical Reactions

Acid-Base Reactions Leading to Gas Formation

Some acid-base reactions result in the formation of gases. This typically occurs when acids react with certain anions such as carbonates, bicarbonates, sulfites, and sulfides. The gas produced depends on the anion present in the reactant.

• Carbonates (CO32−) and Bicarbonates (HCO3−): React with acids to form carbonic acid (H2CO3), which decomposes into carbon dioxide gas and water.

• Sulfites (SO32−): React with acids to form sulfur dioxide gas and water.

• Sulfides (S2−): React with acids to form hydrogen sulfide gas.

Examples:

• $2 \ \mathrm{HNO}_3\ (aq) + \mathrm{Na}_2\mathrm{CO}_3\ (aq) \rightarrow 2 \ \mathrm{NaNO}_3\ (aq) + \mathrm{H}_2\mathrm{CO}_3\ (aq)$ $\mathrm{H}_2\mathrm{CO}_3\ (aq) \rightarrow \mathrm{CO}_2\ (g) + \mathrm{H}_2\mathrm{O}\ (l)$

• Net ionic: $2 \ \mathrm{H}^+\ (aq) + \mathrm{CO}_3^{2-}\ (aq) \rightarrow \mathrm{CO}_2\ (g) + \mathrm{H}_2\mathrm{O}\ (l)$

• $\mathrm{Li}_2\mathrm{SO}_3\ (aq) + 2 \ \mathrm{HCl}\ (aq) \rightarrow 2 \ \mathrm{LiCl}\ (aq) + \mathrm{SO}_2\ (g) + \mathrm{H}_2\mathrm{O}\ (l)$

• Net ionic: $\mathrm{SO}_3^{2-}\ (aq) + 2 \ \mathrm{H}^+\ (aq) \rightarrow \mathrm{SO}_2\ (g) + \mathrm{H}_2\mathrm{O}\ (l)$

• $\mathrm{Rb}_2\mathrm{S}\ (aq) + 2 \ \mathrm{HCl}\ (aq) \rightarrow 2 \ \mathrm{RbCl}\ (aq) + \mathrm{H}_2\mathrm{S}\ (g)$

• Net ionic: $\mathrm{S}^{2-}\ (aq) + 2 \ \mathrm{H}^+\ (aq) \rightarrow \mathrm{H}_2\mathrm{S}\ (g)$

Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between substances, resulting in changes in oxidation states. These reactions are distinct from acid-base reactions, which involve proton transfer. Redox processes are fundamental in both natural and industrial chemical transformations.

• Oxidation: Loss of electrons by a substance. The oxidation number increases.

• Reduction: Gain of electrons by a substance. The oxidation number decreases.

• Oxidizing Agent: The substance that is reduced (gains electrons).

• Reducing Agent: The substance that is oxidized (loses electrons).

Examples:

• Rusting of Iron: $4 \ \mathrm{Fe}\ (s) + 3 \ \mathrm{O}_2\ (g) \rightarrow 2 \ \mathrm{Fe}_2\mathrm{O}_3\ (s)$ Iron is oxidized, oxygen is reduced.

• Combustion of Hydrogen: $2 \ \mathrm{H}_2\ (g) + \mathrm{O}_2\ (g) \rightarrow 2 \ \mathrm{H}_2\mathrm{O}\ (g)$

• Combustion of Octane: $2 \ \mathrm{C}_8\mathrm{H}_{18}\ (l) + 25 \ \mathrm{O}_2\ (g) \rightarrow 16 \ \mathrm{CO}_2\ (g) + 18 \ \mathrm{H}_2\mathrm{O}\ (g)$

• Formation of Sodium Chloride: $2 \ \mathrm{Na}\ (s) + \mathrm{Cl}_2\ (g) \rightarrow 2 \ \mathrm{NaCl}\ (s)$ Half-reactions: $\mathrm{Na}\ (s) \rightarrow \mathrm{Na}^+ + e^-$ (oxidation) $\mathrm{Cl}_2 + 2e^- \rightarrow 2 \ \mathrm{Cl}^-$ (reduction)

Assigning Oxidation Numbers

Oxidation numbers (states) are assigned to atoms in molecules and ions to keep track of electron transfer in redox reactions. The following rules are used:

1. Free elements: Oxidation number is 0. (e.g., F2, K, O2, P4)

2. Monatomic ions: Oxidation number equals the ion's charge. (e.g., Na+ = +1, S2− = -2)

3. Oxygen: Usually -2; in peroxides, -1. (e.g., H2O2)

4. Hydrogen: Usually +1; in metal hydrides, -1. (e.g., NaH)

5. Fluorine: Always -1. Other halogens are -1 with metals, but can be positive with F or O.

6. Neutral molecules: Sum of oxidation numbers is 0.

7. Polyatomic ions: Sum of oxidation numbers equals the ion's charge.

8. Non-integer values: Some compounds have non-integer oxidation numbers. (e.g., NaO2: O = -1/2)

9. Group 1A metals: Always +1 in compounds.

10. Group 2A metals: Always +2 in compounds.

Additional info: Metals only have positive oxidation numbers; nonmetals can have positive or negative values. Transition metals often have multiple possible oxidation states (e.g., Cu+ and Cu2+).

Types of Redox Reactions

Redox reactions can be classified into several types based on the nature of the reactants and products:

• Combination Reactions: Two or more reactants form a single product. Example: $2 \ \mathrm{S}\ (s) + 3 \ \mathrm{O}_2\ (g) \rightarrow 2 \ \mathrm{SO}_3\ (g)$

• Decomposition Reactions: A single reactant breaks down into multiple products. Example: $2 \ \mathrm{KBrO}_3\ (s) \rightarrow 2 \ \mathrm{KBr}\ (s) + 3 \ \mathrm{O}_2\ (g)$

• Combustion Reactions: A substance reacts with oxygen, releasing heat and light. Example: $\text{hydrocarbon} + \mathrm{O}_2 \rightarrow \mathrm{CO}_2\ (g) + \mathrm{H}_2\mathrm{O}\ (g)$

• Displacement Reactions: An atom or ion in a compound is replaced by another atom or ion.

Displacement Reactions

• Hydrogen Displacement: Reactive metals (Group 1A and some Group 2A) displace hydrogen from water or acids. Examples:

  • $2 \ \mathrm{K}\ (s) + 2 \ \mathrm{H}_2\mathrm{O}\ (l) \rightarrow 2 \ \mathrm{KOH}\ (aq) + \mathrm{H}_2\ (g)$

  • $\mathrm{Zn}\ (s) + 2 \ \mathrm{HCl}\ (aq) \rightarrow \mathrm{ZnCl}_2\ (aq) + \mathrm{H}_2\ (g)$

• Metal Displacement: A more active metal displaces a less active metal from its compound. Example:

  • $\mathrm{CuCl}_2\ (aq) + \mathrm{Zn}\ (s) \rightarrow \mathrm{ZnCl}_2\ (aq) + \mathrm{Cu}\ (s)$

  • Net ionic: $\mathrm{Zn}\ (s) + \mathrm{Cu}^{2+}\ (aq) \rightarrow \mathrm{Cu}\ (s) + \mathrm{Zn}^{2+}\ (aq)$

  Activity Series: A ranking of metals by their tendency to be oxidized. Metals above hydrogen in the series can displace hydrogen from acids or water.

• Halogen Displacement: A more reactive halogen displaces a less reactive halide ion from solution. Reactivity order: F2 > Cl2 > Br2 > I2 Examples:

  • $\mathrm{Cl}_2\ (g) + 2 \ \mathrm{NaBr}\ (aq) \rightarrow 2 \ \mathrm{NaCl}\ (aq) + \mathrm{Br}_2\ (l)$

  • $\mathrm{F}_2\ (g) + 2 \ \mathrm{NaCl}\ (aq) \rightarrow 2 \ \mathrm{NaF}\ (aq) + \mathrm{Cl}_2\ (g)$

  • $\mathrm{Br}_2\ (l) + 2 \ \mathrm{NaI}\ (aq) \rightarrow 2 \ \mathrm{NaBr}\ (aq) + \mathrm{I}_2\ (s)$

• Disproportionation Reactions: An element is simultaneously oxidized and reduced in the same reaction. Example: $\mathrm{Br}_2 \rightarrow \mathrm{BrO}_3^- + \mathrm{Br}^-$

Balancing Redox Reactions

Redox reactions are balanced by ensuring that the number of electrons lost in oxidation equals the number gained in reduction. This is often done by writing half-reactions for oxidation and reduction, balancing them for mass and charge, and then combining them.

Example 1:

• $\mathrm{Cu}^{2+}\ (aq) + \mathrm{Fe}\ (s) \rightarrow \mathrm{Fe}^{3+}\ (aq) + \mathrm{Cu}\ (s)$

• Half-reactions:

  • Reduction: $3 \ \mathrm{Cu}^{2+}\ (aq) + 6e^- \rightarrow 3 \ \mathrm{Cu}\ (s)$

  • Oxidation: $2 \ \mathrm{Fe}\ (s) \rightarrow 2 \ \mathrm{Fe}^{3+}\ (aq) + 6e^-$

• Combined: $3 \ \mathrm{Cu}^{2+}\ (aq) + 2 \ \mathrm{Fe}\ (s) \rightarrow 2 \ \mathrm{Fe}^{3+}\ (aq) + 3 \ \mathrm{Cu}\ (s)$

Example 2:

• $\mathrm{Mg}\ (s) + \mathrm{N}_2\ (g) \rightarrow \mathrm{Mg}_3\mathrm{N}_2\ (s)$

• Half-reactions:

  • Oxidation: $3 \ \mathrm{Mg}\ (s) \rightarrow 3 \ \mathrm{Mg}^{2+}\ (aq) + 6e^-$

  • Reduction: $\mathrm{N}_2\ (g) + 6e^- \rightarrow 2 \ \mathrm{N}^{3-}\ (aq)$

• Combined: $3 \ \mathrm{Mg}\ (s) + \mathrm{N}_2\ (g) \rightarrow 3 \ \mathrm{Mg}^{2+}\ (aq) + 2 \ \mathrm{N}^{3-}\ (aq)$ (which forms $\mathrm{Mg}_3\mathrm{N}_2$)

• Identifying Agents: Mg is oxidized (reducing agent); N2 is reduced (oxidizing agent).

Summary Table: Activity Series of Metals and Halogens

Metal Activity Series (Partial)	Halogen Activity Series
Li > K > Ba > Ca > Na > Mg > Al > Zn > Fe > Pb > (H) > Cu > Ag > Au Metals above H: Displace H from acids/water Metals below H: Do not displace H	F2 > Cl2 > Br2 > I2 More reactive halogen displaces less reactive halide ion

Additional info: The activity series is a practical tool for predicting the outcomes of single displacement reactions. The higher a metal or halogen is in its series, the more readily it is oxidized (for metals) or reduced (for halogens).