Acids, Bases, and Buffers: Foundations and Applications in Chemistry

Acids, Bases, and Buffers

Introduction

This study guide covers the fundamental concepts of acids, bases, and buffers, including their definitions, properties, and roles in chemical and biological systems. It is designed for students in an Introduction to Chemistry course.

Definitions and Historical Context

Early Classification of Acids and Bases

• Acids: Substances that taste sour, are corrosive to metals, change litmus paper red, and become less acidic when mixed with bases.

• Bases: Substances that taste bitter, feel slippery, change litmus paper blue, and become less basic when mixed with acids.

• Safety Note: Never touch or taste chemical solutions.

Theories of Acids and Bases

Brønsted-Lowry Definition

• Acid: A substance that can donate a proton (H+).

• Base: A substance that can accept a proton (H+).

• General reaction: $HA + B \rightarrow A^- + BH^+$ Where HA is the acid, B is the base, A- is the conjugate base, and BH+ is the conjugate acid.

• Each acid has a corresponding conjugate base, and each base has a corresponding conjugate acid.

• Examples:

  • $HCl + Cl^- \rightarrow H^+ + Cl^-$ (HCl is acid, Cl- is conjugate base)

  • $NH_3 + H^+ \rightarrow NH_4^+$ (NH3 is base, NH4+ is conjugate acid)

Arrhenius Definition

• Acids: Substances that produce hydrogen ions (H+) when dissolved in water.

• Bases: Substances that produce hydroxide ions (OH-) when dissolved in water.

• Examples:

  • $HCl_{(aq)} \rightarrow H^+ + Cl^-$

  • $NaOH_{(aq)} \rightarrow Na^+ + OH^-$

Acid-Base Properties of Water

Hydrogen Bonds and Polarity

• Water molecules are polar, with a partial negative charge on oxygen and partial positive charges on hydrogen.

• Hydrogen bonding makes water a versatile solvent.

Auto-Ionization of Water

• Water can ionize itself: $H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-$

• Dynamic equilibrium exists between hydronium ions ($H_3O^+$) and hydroxide ions ($OH^-$).

Amphoteric Nature of Water

• Water can act as both an acid and a base, depending on the dissolved substance.

• A substance that can act as either an acid or base is called amphoteric.

Acids and Bases in Water

Acids in Water

• An acid donates a proton to water: $AH + H_2O \rightarrow A^- + H_3O^+$

• A- is the conjugate base of the acid; $H_3O^+$ is the conjugate acid of water.

Bases in Water

• A base accepts a proton from water: $B + H_2O \rightarrow BH^+ + OH^-$

• BH+ is the conjugate acid of the base; $OH^-$ is the conjugate base of water.

Neutralization Reactions

Definition and Example

• Neutralization occurs when H+ ions from an acid combine with OH- ions from a base to form water: $H^+ + OH^- \rightarrow H_2O$

• Example: $HCl + NaOH \rightarrow H_2O + NaCl$

• The resulting solution is neutral (neither acidic nor basic).

pH and the pH Scale

Water Ionization and pH Calculation

• In neutral water at 25°C: $[H^+] = [OH^-] = 1.0 \times 10^{-7}$ M

• The equilibrium expression for water ionization: $K_w = [H^+][OH^-]$ $K_w = 1.0 \times 10^{-14}$ at 25°C

• pH is defined as: $pH = -\log[H^+]$

• pH 7 is neutral at 25°C.

The pH Scale

• Ranges from 0 (most acidic) to 14 (most basic).

• Each unit change in pH represents a tenfold change in [H+].

• Acidic: $[H^+] > [OH^-]$; Basic: $[H^+] < [OH^-]$; Neutral: $[H^+] = [OH^-]$

Table: pH, [H+], and [OH-] Relationships

pH	[H+] (M)	[OH-] (M)
0	1 × 100	1 × 10-14
1	1 × 10-1	1 × 10-13
2	1 × 10-2	1 × 10-12
3	1 × 10-3	1 × 10-11
4	1 × 10-4	1 × 10-10
5	1 × 10-5	1 × 10-9
6	1 × 10-6	1 × 10-8
7	1 × 10-7	1 × 10-7
8	1 × 10-8	1 × 10-6
9	1 × 10-9	1 × 10-5
10	1 × 10-10	1 × 10-4
11	1 × 10-11	1 × 10-3
12	1 × 10-12	1 × 10-2
13	1 × 10-13	1 × 10-1
14	1 × 10-14	1 × 100

Table: pH of Common Solutions

pH Value	Aqueous Solution
1	Battery Acid
2	Gastric Acid (stomach), Lemon
3	Vinegar, Beer, Wine, Cola
5	Rainwater
6	Urine
7	Pure Water
7.4	Human Blood
8.5	Seawater
11.5	Household Ammonia
12.5	Household Bleach
13.5	Oven Cleaner

Table: pH Indicator Colour Chart

pH	Methyl Orange	Methyl Red	Phenol Red	Phenolphthalein
1	red	red	yellow	milky white
3.2	orange	red	yellow	milky white
4.2	red	yellow	yellow	milky white
6.8	yellow	yellow	yellow	milky white
8.2	yellow	yellow	pink	milky pink
10	yellow	yellow	red	red
12	yellow	yellow	red	red
14	yellow	yellow	red	red

Strengths of Acids and Bases

Strong Acids

• Completely ionize in water to donate all protons.

• Examples: HCl, H2SO4, HNO3

• For 0.1 M HCl: $[H^+] = 0.1$ M

Weak Acids

• Only partially ionize in water.

• Equilibrium: $HA \rightleftharpoons A^- + H^+$

• Acid dissociation constant: $K_a = \frac{[A^-][H^+]}{[HA]}$

pKa

• Negative logarithm of $K_a$:

$pK_a = -\log_{10}K_a$

• Formic acid: $pK_a = 3.75$

• Acetic acid: $pK_a = 4.76$

• The smaller the pKa, the stronger the acid.

Strong and Weak Bases

• Strong bases (e.g., NaOH) dissociate completely in water.

• pH of strong base solution can be calculated using $K_w$:

$[H^+][OH^-] = 10^{-14}$

$[H^+] = \frac{10^{-14}}{[OH^-]}$

• Weak bases only partially ionize; their strength is often measured by the pKa of their conjugate acid.

Table: Base Strength and Conjugate Acid pKa

Base	Conjugate Acid	pKa	Strength
OH-	H2O	14	Very weak acid
NH3	NH4+	9.2	Weak acid

Buffer Solutions

Definition and Characteristics

• Buffers are solutions that resist large changes in pH when acids or bases are added.

• Composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Keep the pH constant within a limited range and capacity.

Acidic Buffer Solutions

• pH less than 7.

• Made from a weak acid and one of its salts (e.g., sodium acetate).

• Example: $CH_3COOH + H_2O \rightleftharpoons CH_3COO^- + H_3O^+$

Basic Buffer Solutions

• pH greater than 7.

• Made from a weak base and one of its salts (e.g., ammonia buffer).

• Example: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$

Common Buffers Table

Buffer	pKa	Neutral Form (g/L)	Acid Form (g/L)	Base Form (g/L)
Acetate	4.76	---	---	---
Phosphate	7.2	---	---	---
Tris	8.1	---	---	---
HEPES	7.5	---	---	---
Carbonate	10.3	---	---	---
Additional info: Table entries inferred for illustration; refer to full buffer tables for laboratory use.

The Henderson-Hasselbalch Equation

Relationship Between pH and Buffer Components

• Describes how pH relates to the concentrations of acid and conjugate base in a buffer.

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

• Can be used to determine the pH of a buffer solution.

Example Calculations

• For a solution with 0.5 M acetic acid and 0.5 M sodium acetate ($pK_a = 4.75$): $pH = 4.75 + \log\left(\frac{0.5}{0.5}\right) = 4.75$

• For a buffer with 0.10 M CH3COOH and 0.15 M CH3COONa ($pK_a = 4.76$): $pH = 4.76 + \log\left(\frac{0.15}{0.10}\right) = 4.76 + 0.18 = 4.94$

• Adding 0.01 M HCl to 1 L of the above buffer: New concentrations: $[CH_3COOH] = 0.11$ M, $[CH_3COO^-] = 0.14$ M $pH = 4.76 + \log\left(\frac{0.14}{0.11}\right) = 4.86$ pH change is very small (buffer resists change).

Biological Buffer Systems

Blood Buffer System

• Blood maintains a pH of 7.4, optimal for oxygen transport.

• Carbonic acid-bicarbonate buffer system: $H_2CO_3 + H_2O \rightleftharpoons HCO_3^- + H_3O^+$

• CO2 in the blood participates in this system: $CO_2 + 2H_2O \rightleftharpoons HCO_3^- + H_3O^+$

• Higher CO2 shifts equilibrium right, increasing H+ and lowering pH (acidosis).

• Lower CO2 shifts equilibrium left, decreasing H+ and making pH more basic (alkalosis).

Importance of pH and pH Control

• Physiology: Blood pH is critical; even small shifts can be fatal.

• Agriculture: Crop growth depends on soil pH.

• Drug Delivery: Drug formulation, administration, and potency are affected by pH.

• Certain drugs lose potency or become toxic depending on pH and route of administration.

Summary of Key Learning Outcomes

• Brønsted-Lowry and Arrhenius definitions of acids and bases

• Water ionization, pH calculation, and the pH scale

• Calculation of pKa for conjugate acids and bases

• Henderson-Hasselbalch equation and buffer systems

• Importance of pH and pH control in chemical and biological contexts

• Blood as a biological buffer system

Additional info: Some buffer table entries and indicator chart details inferred for completeness. For laboratory use, consult full reference tables.