Acid-Base and Oxidation-Reduction Reactions: Study Notes for General Chemistry

Acid-Base Reactions

Introduction to Acid-Base Reactions

Acid-base reactions are fundamental chemical processes in which acids and bases interact, often resulting in the formation of water and a salt. Understanding these reactions is essential for predicting product formation and for quantitative analysis in titrations.

• Acid: Produces H+ ions in water.

• Base: Produces OH- ions in water.

• Be able to write net ionic equations for acid-base reactions.

• Strong acids and strong bases ionize (or dissociate) completely in water.

• Weak acids and weak bases do not completely ionize (will be specified in exams).

Net Ionic Equations for Acid-Base Reactions

Net ionic equations show only the species that actually participate in the reaction, omitting spectator ions.

• Example: The reaction between hydrochloric acid and potassium hydroxide.

• Complete ionic equation: $\mathrm{HCl}(aq) + \mathrm{KOH}(aq) \rightarrow \mathrm{KCl}(aq) + \mathrm{H_2O}(l)$

• Net ionic equation: $\mathrm{H^+}(aq) + \mathrm{OH^-}(aq) \rightarrow \mathrm{H_2O}(l)$

Additional info: Spectator ions (K+ and Cl-) are omitted in the net ionic equation.

Acid-Base Titrations

Introduction to Titrations

Titration is a laboratory technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

• A standard solution of known concentration is added to a solution of unknown concentration until the reaction is complete (equivalence point).

• Indicators are often used to signal the end point of the titration.

Example: Aspirin Tablet Titration

• 25.0 mL of 0.500 M NaOH is required to react with an aspirin tablet weighing 2.50 g.

• Calculate the acetylsalicylic acid (active component) in the tablet.

• Reaction: $\mathrm{NaOH}(aq) + \mathrm{HC_9H_7O_4}(aq) \rightarrow \mathrm{NaC_9H_7O_4}(aq) + \mathrm{H_2O}(l)$

Application: Vinegar Analysis

• To be called "vinegar," a product must have at least 5% acetic acid (HC2H3O2).

• Example calculation: 100.0 g sample of "raspberry vinegar" titrated with 0.1250 M NaOH, requiring 25.0 mL for complete neutralization.

• Determine if the product meets the standard for "vinegar."

Oxidation-Reduction Reactions (REDOX)

Introduction to Redox Reactions

Oxidation-reduction (redox) reactions involve the transfer of electrons between chemical species. These reactions are essential in energy production, corrosion, and many industrial processes.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Redox reactions can be identified by changes in oxidation states of elements.

Oxidation State/Number

Oxidation state (or number) is a system to keep track of electrons involved in the transfer during redox reactions.

• Recall: Covalent sharing is not always equal; some atoms attract electrons more strongly.

• Order of electronegativity: F > O > N > Cl

Rules for Assigning Oxidation Numbers

• Elemental form: Oxidation number is 0 (e.g., O2, N2).

• Monatomic ion: Oxidation number equals the charge (e.g., Na+ is +1).

• Oxygen: Usually -2 (except in peroxides, where it is -1).

• Hydrogen: +1 when bonded to nonmetals, -1 when bonded to metals.

• Fluorine: Always -1.

• The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, it equals the ion's charge.

Example: Assigning Oxidation Numbers

• Na: 0 (elemental form)

• NaCl: Na is +1, Cl is -1

• Ag2O: Ag is +1, O is -2

• NaNO3: Na is +1, N is +5, O is -2

Balancing Redox Reactions

Methods for Balancing Redox Reactions

Redox reactions cannot always be balanced by simple inspection. The half-reaction method is commonly used, especially for reactions in acidic or basic medium.

• Half-reaction method involves separating the oxidation and reduction processes, balancing each for mass and charge, and then combining them.

• Acidic and basic media require different approaches for balancing.

Half-Reaction Method (Acidic Medium)

1. Write the oxidation and reduction half-reactions.

2. Balance all elements except H and O.

3. Balance O by adding H2O.

4. Balance H by adding H+.

5. Balance charge by adding electrons.

6. Combine the half-reactions and ensure electrons cancel.

Half-Reaction Method (Basic Medium)

1. Follow the steps for acidic medium.

2. After balancing, add OH- to both sides to neutralize H+ and form H2O.

Example: Balancing Redox Reaction in Acidic Medium

• Balance: $\mathrm{Cl^-} + \mathrm{MnO_4^-} \rightarrow \mathrm{Cl_2} + \mathrm{Mn^{2+}}$

• Identify which species are oxidized and reduced, and the corresponding agents.

Example: Balancing Redox Reaction in Basic Medium

• Balance: $\mathrm{Cl^-} + \mathrm{MnO_4^-} \rightarrow \mathrm{Cl_2} + \mathrm{MnO_2}$

• Balance the following in both acidic and basic medium: $\mathrm{NO_2^-} (aq) + \mathrm{Al}(s) + \mathrm{NH_4^+}(aq) \rightarrow \mathrm{N_2}(g) + \mathrm{AlO_2^-}(aq)$

Redox Titrations

Example: Permanganate-Iron(II) Titration

• Balanced equation: $\mathrm{MnO_4^-} (aq) + 5\mathrm{Fe^{2+}} (aq) + 8\mathrm{H^+} (aq) \rightarrow \mathrm{Mn^{2+}} (aq) + 5\mathrm{Fe^{3+}} (aq) + 4\mathrm{H_2O} (l)$

• Calculate the volume of 0.64 M KMnO4 solution required to react completely with 27.50 mL of 0.250 M Fe(NO3)2.

Practice Problems

• Assign oxidation numbers to elements in various compounds (e.g., Sc2O3, H2O2, Al, NH4CrO4).

• Balance redox reactions in acidic and basic media.

• Calculate volumes and concentrations in titration problems.

Summary Table: Common Oxidation States

Element	Common Oxidation State(s)	Notes
Na	+1	Alkali metal, always +1 in compounds
Cl	-1	Halide, usually -1 except in compounds with O or F
O	-2	Except in peroxides (-1) and OF2 (+2)
H	+1	With nonmetals; -1 with metals
F	-1	Most electronegative, always -1

Additional info: This table summarizes the most common oxidation states for selected elements, useful for assigning oxidation numbers in redox reactions.