Acids, Bases, Buffers, and Electrochemistry: General Chemistry II Study Notes

Acids and Bases

Bronsted-Lowry Acid-Base Theory

The Bronsted-Lowry theory defines acids as proton (H+) donors and bases as proton acceptors. Each acid-base reaction involves the transfer of a proton from the acid to the base, forming conjugate acid-base pairs.

• Acid: Species that donates a proton.

• Base: Species that accepts a proton.

• Conjugate Acid: Formed when a base gains a proton.

• Conjugate Base: Formed when an acid loses a proton.

• Example: $\mathrm{HCl} + \mathrm{H_2O} \rightarrow \mathrm{Cl^-} + \mathrm{H_3O^+}$ - HCl is the acid, H2O is the base, Cl- is the conjugate base, H3O+ is the conjugate acid.

Acid and Base Strength

Acid and base strength is determined by the degree of ionization in water. Strong acids/bases ionize completely, while weak acids/bases ionize partially.

• Strong Acids: HCl, HNO3, H2SO4

• Weak Acids: CH3COOH, HF

• Strong Bases: NaOH, KOH

• Weak Bases: NH3

• Relationship: The stronger the acid, the weaker its conjugate base.

pH and pOH Calculations

pH is a measure of the hydrogen ion concentration in a solution. It is calculated using the following formula:

• $\mathrm{pH} = -\log[\mathrm{H^+}]$

• $\mathrm{pOH} = -\log[\mathrm{OH^-}]$

• $\mathrm{pH} + \mathrm{pOH} = 14$

• Example: For $[\mathrm{H^+}] = 1.0 \times 10^{-3}$ M, $\mathrm{pH} = 3$

Acid Dissociation Constant (Ka)

The acid dissociation constant quantifies the strength of a weak acid in solution.

• $K_a = \frac{[\mathrm{H^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$

• Relationship to pKa: $\mathrm{p}K_a = -\log K_a$

• Example: If pH and concentration are known, $K_a$ can be calculated.

Buffers and Buffer Calculations

Buffer Solutions

Buffers are solutions that resist changes in pH upon addition of small amounts of acid or base. They typically consist of a weak acid and its conjugate base.

• Common Buffer System: Acetic acid (CH3COOH) and acetate (CH3COO-)

• Buffer Equation (Henderson-Hasselbalch):

$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[\mathrm{A^-}]}{[\mathrm{HA}]} \right)$

• Example: Calculate the pH after adding acid or base to a buffer using the above equation.

Buffer Capacity and Preparation

Buffer capacity refers to the amount of acid or base a buffer can neutralize before the pH changes significantly.

• Preparation: Mix a weak acid with its conjugate base in appropriate ratios.

• Example: Adding NaOH to a buffer solution and recalculating pH.

Titrations and Acid-Base Curves

Titration of Acids and Bases

Titration is a technique used to determine the concentration of an unknown acid or base by reacting it with a standard solution.

• Strong Acid with Strong Base: Sharp pH change at equivalence point.

• Weak Acid with Strong Base: Gradual pH change, buffer region before equivalence point.

• Equivalence Point: The point at which moles of acid equal moles of base.

Titration Curves

Titration curves plot pH versus volume of titrant added. The shape of the curve depends on the strength of the acid and base.

• Strong Acid/Strong Base: Steep rise at equivalence.

• Weak Acid/Strong Base: Buffer region, less steep at equivalence.

• Weak Base/Strong Acid: Similar buffer region, but pH decreases.

Electrochemistry

Redox Reactions

Electrochemistry involves chemical reactions where electrons are transferred between species (oxidation-reduction reactions).

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation, is itself reduced.

• Reducing Agent: Causes reduction, is itself oxidized.

• Example: $\mathrm{Cu(s)} + \mathrm{2Ag^+(aq)} \rightarrow \mathrm{Cu^{2+}(aq)} + \mathrm{2Ag(s)}$

Balancing Redox Reactions

Redox reactions are balanced by separating into half-reactions and ensuring both mass and charge are balanced.

• Steps:

  1. Write oxidation and reduction half-reactions.

  2. Balance atoms other than O and H.

  3. Balance O by adding H2O.

  4. Balance H by adding H+.

  5. Balance charge by adding electrons.

  6. Combine half-reactions and cancel electrons.

Galvanic Cells

Galvanic (voltaic) cells convert chemical energy into electrical energy via spontaneous redox reactions.

• Anode: Electrode where oxidation occurs (loss of electrons).

• Cathode: Electrode where reduction occurs (gain of electrons).

• Cell Notation: Anode | Anode solution || Cathode solution | Cathode

• Example: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Cell Potentials

The cell potential (Ecell) is the voltage produced by a galvanic cell, calculated from standard reduction potentials.

• $E_{cell} = E_{cathode} - E_{anode}$

• Standard Reduction Potentials: Found in tables; higher E0 means stronger oxidizing agent.

Tables

Acid and Conjugate Base Strengths

The following table compares the strength of acids and their conjugate bases:

Acid	Conjugate Base	Acid Strength	Conjugate Base Strength
HClO4	ClO4-	Strong	Very weak
HCl	Cl-	Strong	Very weak
HNO3	NO3-	Strong	Very weak
CH3COOH	CH3COO-	Weak	Moderate
H2O	OH-	Very weak	Strong

Solution Classification Table

Classify solutions as acidic, basic, or neutral:

Solution	Acidic	Basic	Neutral
NaCl			X
KNO3			X
NH4Cl	X
NaCH3COO		X
NaF		X
NH4NO3	X

Additional info:

• Some calculations and explanations were expanded for clarity and completeness.

• Examples and equations were added to illustrate key concepts.

• Tables were reconstructed and summarized based on the provided content.