Acids, Bases, and Equilibrium: Comprehensive Study Notes

Acids and Bases: Introduction and Clinical Relevance

Clinical Laboratory Applications

Acids and bases play a crucial role in clinical laboratory settings, where technicians analyze specimens and interpret test results for medical diagnoses, such as cancer detection and blood typing. Understanding acid-base chemistry is essential for accurate analysis and interpretation.

Key Math and Chemistry Skills for Acid-Base Chemistry

• Solving Equations: Essential for quantitative analysis.

• Scientific Notation: Used for expressing concentrations.

• Writing Ionic Formulas: Important for representing acids and bases.

• Balancing Chemical Equations: Fundamental for reaction stoichiometry.

• Mole–Mole Factors: Used in titration and reaction calculations.

• Concentration Conversions: Key for solution chemistry.

Acids and Bases: Definitions and Properties

Acids: Characteristics and Naming

Acids are molecular substances that produce ions in water, including the hydronium ion (H3O+). They are classified as electrolytes and are known for their sour taste, ability to turn blue litmus red, and their corrosive properties.

• Arrhenius Acids: Produce H+ ions in water.

• Brønsted–Lowry Acids: Donate H+ ions.

• Naming Acids:

  • Acids with H+ and a nonmetal: Prefix "hydro" and suffix "ic acid" (e.g., HCl: hydrochloric acid).

  • Acids with H+ and a polyatomic ion: "ate" becomes "ic acid"; "ite" becomes "ous acid" (e.g., HNO3: nitric acid).

Bases: Characteristics and Naming

Bases are substances that produce hydroxide ions (OH−) in water. They are electrolytes, taste bitter or chalky, feel slippery, and turn litmus paper blue.

• Arrhenius Bases: Produce OH− ions in water.

• Brønsted–Lowry Bases: Accept H+ ions.

• Naming Bases: Named as hydroxides (e.g., NaOH: sodium hydroxide).

Acid-Base Theories

Brønsted–Lowry Theory

According to the Brønsted–Lowry theory, acids donate H+ ions, while bases accept H+ ions. This theory expands the definition of acids and bases beyond aqueous solutions.

Characteristics of Acids and Bases

Acids and bases differ in their properties, including taste, feel, and reactions with indicators.

Characteristic	Acids	Bases
Arrhenius	Produce H+	Produce OH−
Brønsted–Lowry	Donate H+	Accept H+
Electrolytes	Yes	Yes
Taste	Sour	Bitter, chalky
Feel	May sting	Soapy, slippery
Litmus	Red	Blue
Phenolphthalein	Colorless	Pink
Neutralization	Neutralize bases	Neutralize acids

Conjugate Acid–Base Pairs

Definition and Identification

In acid–base reactions, conjugate pairs are formed by the loss or gain of H+. Each acid has a conjugate base, and each base has a conjugate acid.

• Conjugate Acid–Base Pair Example: HF and F−; H2O and H3O+

Strengths of Acids and Bases

Strong vs. Weak Acids

Strong acids completely ionize in water, producing high concentrations of H3O+. Weak acids only partially dissociate, resulting in lower ion concentrations.

• Strong Acids: HCl, HBr, HI, HNO3, H2SO4, HClO4

• Weak Acids: HF, H2CO3, CH3COOH

Strong vs. Weak Bases

Strong bases, such as NaOH and KOH, dissociate completely in water. Weak bases, like NH3 and Cu(OH)2, only partially ionize.

Acid–Base Equilibrium

Reversible Reactions and Equilibrium

Acid–base reactions can reach equilibrium, where the rates of the forward and reverse reactions are equal and concentrations remain constant.

Le Châtelier’s Principle

Le Châtelier’s principle states that a system at equilibrium will adjust to relieve any stress caused by changes in concentration, temperature, or pressure.

Stress	Change	Shift in Direction of
Add reactant	Increases forward reaction rate	Products
Remove reactant	Decreases forward reaction rate	Reactants
Add product	Increases reverse reaction rate	Reactants
Remove product	Decreases reverse reaction rate	Products

Dissociation of Water and Ion Product Constant (Kw)

Ionization of Water

Water is amphoteric and can act as both an acid and a base. The dissociation of water produces equal concentrations of H3O+ and OH− ions.

• Ion Product Constant: $Kw = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C

The pH Scale

Definition and Calculation

The pH scale measures the acidity or basicity of a solution, ranging from 0 (very acidic) to 14 (very basic). pH is calculated as the negative logarithm of the hydronium ion concentration:

• pH Equation: $pH = -\log[H_3O^+]$

Reactions of Acids and Bases

Acids Reacting with Metals, Carbonates, and Bases

Acids react with metals to produce hydrogen gas and salts, with carbonates and bicarbonates to produce carbon dioxide, and with bases in neutralization reactions to form water and salts.

Acid–Base Titration

Principles and Calculations

Titration is used to determine the concentration of an acid or base by neutralizing it with a solution of known concentration. The endpoint is identified using an indicator such as phenolphthalein.

Chemistry Link to Health: Antacids

Antacids neutralize excess stomach acid and are composed of aluminum hydroxide and magnesium hydroxide. These compounds are only slightly soluble, minimizing adverse effects on the digestive tract.

Buffers

Definition and Function

A buffer solution maintains pH by neutralizing small amounts of added acid or base. It contains a weak acid and its conjugate base, or a weak base and its conjugate acid.

Acids and Bases—Concept Map

Additional info: These notes cover all major aspects of Chapter 10: Acids and Bases and Equilibrium, including definitions, properties, reactions, calculations, and clinical relevance. All images included are directly relevant to the adjacent content and reinforce key concepts for college-level chemistry students.