Acids and Bases: Chemistry Fundamentals

• Acids and bases are fundamental concepts in chemistry with numerous applications in everyday life and various industries.
• Acids and bases are defined in multiple ways: Arrhenius, Bronsted-Lowry, and Lewis definitions offer different perspectives.

Arrhenius Definition

• Acids are substances that produce hydrogen ions (H+) when dissolved in water.
• Bases are substances that produce hydroxide ions (OH−) when dissolved in water.
• Neutralization involves the reaction of H+ and OH− ions to form water and a salt.
• Limited scope as it only applies to aqueous solutions and substances that directly produce H+ or OH− ions.

Bronsted-Lowry Definition

• Acids are proton (H+) donors.
• Bases are proton acceptors.
• Broader definition; acids and bases don't necessarily need to be in water.
• Includes conjugate acid-base pairs.
• A conjugate acid is formed when a base accepts a proton.
• A conjugate base is formed when an acid donates a proton.
• Example: In the reaction HCl + H2O → H3O+ + Cl−, HCl is the acid, H2O is the base, H3O+ is the conjugate acid, and Cl− is the conjugate base.

Lewis Definition

• Acids are electron-pair acceptors (electrophiles).
• Bases are electron-pair donors (nucleophiles).
• Most inclusive definition; covers reactions that do not involve protons.
• Involves the formation of coordinate covalent bonds.
• Example: BF3 (acid) accepts an electron pair from NH3 (base) to form a complex.

Acid Strength and pH

• Strong acids completely dissociate into ions in water.
• Examples of strong acids: hydrochloric acid (HCl), sulfuric acid (H2SO4), nitric acid (HNO3).
• Weak acids only partially dissociate into ions in water.
• Examples of weak acids: acetic acid (CH3COOH), carbonic acid (H2CO3).
• Acid dissociation constant (Ka) measures the strength of an acid.
• Larger Ka indicates stronger acid.
• Bases can also be strong or weak.
• Strong bases completely dissociate into ions in water.
• Examples of strong bases: sodium hydroxide (NaOH), potassium hydroxide (KOH).
• Weak bases only partially dissociate into ions in water.
• Examples of weak bases: ammonia (NH3), pyridine (C5H5N).
• Base dissociation constant (Kb) measures the strength of a base.
• Larger Kb indicates a stronger base.
• pH is a measure of the acidity or basicity of a solution.
• pH = -log[H+], where [H+] is the concentration of hydrogen ions in moles per liter (M).
• pH scale ranges from 0 to 14.
• pH < 7 indicates an acidic solution.
• pH = 7 indicates a neutral solution.
• pH > 7 indicates a basic (alkaline) solution.
• pOH is a measure of the hydroxide ion concentration.
• pOH = -log[OH−], where [OH−] is the concentration of hydroxide ions in moles per liter (M).
• pH + pOH = 14 in aqueous solutions at 25°C.

Neutralization Reactions

• Reaction between an acid and a base.
• Products usually include a salt and water.
• Example: HCl + NaOH → NaCl + H2O.
• Titration is a technique used to determine the concentration of an acid or a base in a solution.
• Involves the gradual addition of a known concentration of an acid (or base) to a known volume of a base (or acid) until the reaction is complete (equivalence point).
• Equivalence point is the point at which the acid and base have completely reacted.
• Indicators are substances that change color depending on the pH of the solution.
• Used to visually determine the endpoint of a titration, which is an approximation of the equivalence point.

Buffers

• Solutions that resist changes in pH upon addition of small amounts of acid or base.
• Typically composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.
• Example: Acetic acid (CH3COOH) and sodium acetate (CH3COONa) form a buffer.
• Buffering capacity is the amount of acid or base a buffer can neutralize before significant pH change occurs.
• Henderson-Hasselbalch equation relates pH, pKa, and the concentrations of the weak acid and its conjugate base.
• pH = pKa + log([A−]/[HA]), where [A−] is the concentration of the conjugate base and [HA] is the concentration of the weak acid.
• Buffers are crucial in biological systems to maintain stable pH levels necessary for biochemical reactions.
• Example: Blood contains buffers to maintain a pH around 7.4.

Acid-Base Properties of Salts

• Salts are ionic compounds formed from the neutralization reaction of an acid and a base.
• Salts can be neutral, acidic, or basic when dissolved in water, depending on the properties of the ions they produce.
• Salts of strong acids and strong bases form neutral solutions (e.g., NaCl).
• Salts of strong acids and weak bases form acidic solutions (e.g., NH4Cl).
• Salts of weak acids and strong bases form basic solutions (e.g., CH3COONa).
• Salts of weak acids and weak bases can be acidic, basic, or neutral, depending on the relative strengths of the acid and base (e.g., NH4CN).

Applications

• Acids and bases are used widely in industrial processes.
• Sulfuric acid (H2SO4) is used in fertilizer production, chemical synthesis, and petroleum refining.
• Hydrochloric acid (HCl) is used in metal cleaning, food processing, and laboratory analysis.
• Sodium hydroxide (NaOH) is used in the manufacturing of paper, soap, and detergents.
• Acids and bases are in pharmaceuticals.
• Aspirin (acetylsalicylic acid) is a common pain reliever.
• Antacids contain bases like magnesium hydroxide (Mg(OH)2) or calcium carbonate (CaCO3) to neutralize stomach acid.
• Acids and bases are in household products.
• Acetic acid (vinegar) is used in cooking and cleaning.
• Ammonia (NH3) is used in cleaning solutions.
• Soaps and detergents are typically alkaline.
• Acids and bases are present in environmental processes.
• Acid rain, caused by pollutants like sulfur dioxide (SO2) and nitrogen oxides (NOx), can damage ecosystems and structures.
• Buffering systems in natural waters help to maintain stable pH levels, which are important for aquatic life.
• Acid-base chemistry is vital in biological systems.
• Enzymes, which catalyze biochemical reactions, are sensitive to pH changes.
• Proper pH balance is essential for the transport of oxygen in the blood and the function of cells.

Indicators

• Indicators are substances that change color depending on the pH of the solution.
• Used to visually determine the endpoint of a titration, which is an approximation of the equivalence point.
• Examples include litmus paper, phenolphthalein, and methyl orange.
• Different indicators change color at different pH ranges.
• Litmus paper turns red in acidic conditions and blue in basic conditions.
• Phenolphthalein is colorless in acidic solutions and pink in basic solutions.
• Methyl orange is red in acidic solutions and yellow in basic solutions.