Acids, Bases, Buffers, and Equilibrium: General Chemistry Study Guide

Acids and Bases

Definitions and Roles

Acids and bases are fundamental concepts in chemistry, describing substances that can donate or accept protons (H+). Their behavior is explained by the Brønsted-Lowry and Lewis definitions.

• Brønsted-Lowry Acid: A substance that donates a proton (H+).

• Brønsted-Lowry Base: A substance that accepts a proton (H+).

• Lewis Acid: An electron pair acceptor.

• Lewis Base: An electron pair donor.

• Example: In the reaction H2O + NH3 → OH- + NH4+, water acts as a Brønsted-Lowry acid by donating a proton.

Strength of Acids and Bases

The strength of an acid or base is determined by its ability to donate or accept protons. Strong acids/bases dissociate completely in water, while weak acids/bases only partially dissociate.

• pKa and pKb: Lower pKa values indicate stronger acids; lower pKb values indicate stronger bases.

• Relationship: $pK_a + pK_b = 14$

• Example: If pKa of acid HX is 7.52 and pKa of HY is 10.64, HX is the stronger acid.

pH, pOH, and Calculations

Definitions

pH and pOH are measures of the acidity and basicity of a solution, respectively. They are related to the concentration of hydrogen and hydroxide ions in solution.

• pH: $pH = -\log[H^+]$

• pOH: $pOH = -\log[OH^-]$

• Relationship: $pH + pOH = 14$

• Example: A solution with [H+] = 1.0 × 10-7 M has a pH of 7 (neutral).

Calculating pH of Weak Acids and Bases

Weak acids and bases do not fully dissociate, so equilibrium calculations are required to determine pH.

• Equilibrium Expression for Weak Acid: $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• Percent Ionization: $\%\,\text{ionization} = \frac{[H_3O^+]_{eq}}{[HA]_i} \times 100\%$

• Example: For 0.27 M acetic acid ($K_a = 1.3 \times 10^{-5}$), pH can be calculated using the equilibrium expression.

Buffer Solutions

Definition and Function

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

• Buffer Capacity: The ability of a buffer to maintain pH depends on the concentrations of the acid/base pair.

• Henderson-Hasselbalch Equation: $pH = pK_a + \log \frac{[A^-]}{[HA]}$

• Example: Mixing 0.10 M HC2H3O2 and 0.25 M NH4Cl forms a buffer with high capacity.

Buffer Action and Neutralization

When strong acids or bases are added to a buffer, the weak acid/base pair reacts to minimize changes in pH.

• Example: Adding NaOH to a buffer containing H2PO4- will consume the conjugate acid and shift equilibrium.

Equilibrium Constants

Acid and Base Dissociation Constants

The equilibrium constant for acids ($K_a$) and bases ($K_b$) quantifies their strength in water.

• Acid Dissociation Constant: $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• Base Dissociation Constant: $K_b = \frac{[BH^+][OH^-]}{[B]}$

• Water Dissociation Constant: $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$

• Relationship: $K_w = K_a \times K_b$

Titrations and pH Curves

Titration of Acids and Bases

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

• Equivalence Point: The point at which the amount of titrant added is stoichiometrically equivalent to the amount of analyte in solution.

• pH at Equivalence: For strong acid-strong base titrations, pH = 7. For weak acid-strong base, pH > 7 at equivalence.

• Example: Titration curves can be used to compare the strength of unknown acids by their equivalence point pH values.

Periodic Table and Ion Charges

Periodic Table Overview

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties. Ion charges are indicated for common ions.

• Group 1 (Alkali Metals): +1 charge

• Group 2 (Alkaline Earth Metals): +2 charge

• Halogens: -1 charge

• Transition Metals: Variable charges

• Example: Na+, Ca2+, Cl-

Key Equations and Constants

Essential Equations

• $pH = pK_a + \log \frac{[A^-]}{[HA]}$

• $\%\,\text{ionization} = \frac{[H_3O^+]_{eq}}{[HA]_i} \times 100\%$

• $pK = -\log K$

• $10^{-pK} = K$

• $K_w = [H_3O^+][OH^-] = K_a \times K_b$

• $pK_w = pH + pOH = pK_a + pK_b$

• $K_w = 1.0 \times 10^{-14}$

Sample Table: Common Ion Charges

The following table summarizes common ion charges for selected elements (inferred from the periodic table image):

Element	Symbol	Common Ion Charge
Sodium	Na	+1
Calcium	Ca	+2
Chlorine	Cl	-1
Iron	Fe	+2, +3
Potassium	K	+1
Magnesium	Mg	+2
Aluminum	Al	+3
Oxygen	O	-2

Additional info: Table entries inferred from standard periodic table conventions.

Summary Table: Acid/Base Properties and Calculations

Property	Equation	Description
pH	$pH = -\log[H^+]$	Measures acidity
pOH	$pOH = -\log[OH^-]$	Measures basicity
Ka	$K_a = \frac{[H_3O^+][A^-]}{[HA]}$	Acid dissociation constant
Kb	$K_b = \frac{[BH^+][OH^-]}{[B]}$	Base dissociation constant
Kw	$K_w = [H_3O^+][OH^-]$	Water dissociation constant
Henderson-Hasselbalch	$pH = pK_a + \log \frac{[A^-]}{[HA]}$	Buffer pH calculation

Additional Info

• All equations and constants are standard for General Chemistry and are essential for solving acid-base equilibrium and buffer problems.

• Periodic table ion charges are crucial for predicting compound formation and solution properties.