In-depth Notes on Acids, Bases, and Buffers

Introducing Acids and Bases

• Definition & Properties:

  • Acids:

  • Turn litmus indicator red.

  • Tend to be corrosive.

  • Taste sour.

  • React with bases and produce H⁺ ions in solution (low pH).

  • Conduct electricity.

  • Bases:

  • Turn litmus indicator blue.

  • Are caustic and slippery.

  • Taste bitter.

  • React with acids and produce OH^- ions in solution (high pH).

  • Conduct electricity.

Historical Perspectives on Acids and Bases

• Initially categorized by observed properties (taste, reaction).

• Antoine Lavoisier proposed acidity was due to oxygen.

• Humphrey Davy proposed it was about hydrogen.

• Svante Arrhenius refined the ideas:

  • Acids ionize to form H⁺ ions.

  • Bases ionize to form OH^- ions.

• Brønsted-Lowry Theory:

  • Acids are proton donors.

  • Bases are proton acceptors.

Acid-Base Reactions and Hydrolysis

• Reaction Example:

  • HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl^-(aq)

  • HCl donates a proton to water (acting as acid), forming hydronium ion.

• Hydrolysis Example:

  • HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl^-(aq)

• Dissociation Example:

  • HCl(g) → H⁺(aq) + Cl^-(aq)

• Amphiprotic Substances:

  • Can act as either acid or base.

  • Example: Water can act as both donor and acceptor depending on the solute.

Conjugate Acid-Base Pairs

• A conjugate pair differs by one proton.

• For every acid reaction, there are two conjugate acid-base pairs:

  • HCl and Cl^- (acid/base variant).

  • NH₃ (base) and NH₄⁺ (conjugate acid).

Classifying Acids

• Monoprotic Acids: donate one proton (e.g., HCl, CH₃COOH).

• Polyprotic Acids: donate multiple protons in stages:

  • Diprotic Acids: release two protons (H₂SO₄).

  • Triprotic Acids: release three protons (H₃PO₄).

• Ionization occurs stepwise, where each stage may show varying strengths (e.g., weak vs strong acids).

pH Scale and Acid-Base Chemistry

• pH is a logarithmic scale related to H⁺ concentration:

  • pH = -log[H⁺]

  • Pure water: [H⁺] = 1 × 10^-7 mol/L (neutral pH = 7).

  • Acidic: [H⁺] > 1 × 10^-7 mol/L.

  • Basic: [H⁺] < 1 × 10^-7 mol/L.

Self-Ionization of Water

• 2H₂O(l) ⇌ H₃O⁺(aq) + OH^-(aq)

• Ionization constant (K_w) of water at 25°C:

  • K_w = [H⁺][OH^-] = 1 × 10^-14 mol²/L².

Acid-Base Properties in Aqueous Solutions

• Acidic solutions have excess H⁺:

  • pH < 7.

• Basic solutions need less H⁺:

  • pH > 7.

• Neutral solutions correspond to pH 7:

  • [H⁺] = [OH^-].

Acidity of Salt Solutions

• Salt hydrolysis affects the acidity of the solution:

  • Basic salts (derived from weak acids) release OH^-.

  • Acid salts (derived from weak bases) release H⁺.

Practical Applications of pH

• Indicator use in titrations allows tracking of pH changes visually when adding an acid or base.

• In biological systems, buffers are critical for maintaining pH within acceptable ranges to avoid acidosis or alkalosis in organisms.

Buffers in Biological Systems

• Buffers consist of a weak acid and its conjugate base (or vice-versa) that help maintain stable pH.

• E.g., bicarbonate buffer system in blood:

  • H₂CO₃(aq) ⇌ HCO₃^-(aq) + H⁺(aq)

  • Responds to changes in pH from metabolic activity and maintains homeostasis.

• Definition & Properties:- Acids:

  • Turn litmus indicator red due to their ability to donate protons (H⁺ ions).

  • Tend to be corrosive, meaning they can cause damage to living tissue and materials.

  • Taste sour, a characteristic that can be observed in common acids such as vinegar (acetic acid) and citric acid.

  • React with bases in neutralization reactions, resulting in the formation of water and salt. In solution, they produce H⁺ ions, which contribute to a low pH (typically below 7).

  • Conduct electricity due to the presence of ions in solution.

• Bases:

  • Turn litmus indicator blue, indicating their basic nature.

  • Are caustic and slippery, which can be observed with strong bases like sodium hydroxide (NaOH).

  • Taste bitter, often contrasted with the sour taste of acids.

  • React with acids, producing OH^- ions in solution, contributing to a high pH (typically above 7).

  • Like acids, they also conduct electricity as a result of their ionization in water.

Historical Perspectives on Acids and Bases

• Initially categorized based on their observable properties, such as taste and reactions with various substances.

• Antoine Lavoisier, often regarded as the "father of modern chemistry," proposed in the late 1700s that all acids contained oxygen, a theory later found to be incomplete.

• Humphrey Davy, in the early 1800s, brought forward the idea that acidity was related to hydrogen.

• Svante Arrhenius, in the late 19th century, refined earlier theories by defining:

  • Acids as substances that ionize to produce H⁺ ions in solution.

  • Bases as substances that ionize to produce OH^- ions in solution.

• The Brønsted-Lowry Theory further expanded the definitions:

  • Accurately describes acids as proton donors and bases as proton acceptors, enhancing our understanding of acid-base reactions.

Acid-Base Reactions and Hydrolysis

• Reaction Example: HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl^-(aq)

  • In this reaction, hydrochloric acid (HCl) donates a proton to water, forming the hydronium ion (H₃O⁺), which is responsible for the acidic properties of the solution.

• Hydrolysis Example: Similar to the reaction, it involves the breakdown of a compound in the presence of water, further demonstrating the role water plays in acid-base reactions.

• Dissociation Example: HCl(g) → H⁺(aq) + Cl^-(aq)

  • This illustrates the dissociation of hydrochloric acid in solution, providing clarity on how acids release protons into the solution.

• Amphiprotic Substances:

  • Water is a prime example of an amphiprotic solvent that can both donate and accept protons depending on the solute, playing a crucial role in many biological and chemical processes.

Conjugate Acid-Base Pairs

• A conjugate pair consists of two species that differ by a single proton (H⁺).

• In any acid-base reaction, you can find two conjugate pairs:

  • For example, HCl and Cl^- act as an acid and its conjugate base, respectively.

  • In another instance, ammonia (NH₃) serves as a base and can accept a proton to form its conjugate acid, ammonium (NH₄⁺).

Classifying Acids

• Monoprotic Acids: These acids donate one proton (e.g., HCl, acetic acid, CH₃COOH).

• Polyprotic Acids: These donate multiple protons in steps:

  • Diprotic Acids: Acids like sulfuric acid (H₂SO₄) release two protons in two distinct steps.

  • Triprotic Acids: Acids like phosphoric acid (H₃PO₄) release three protons, highlighting the complexity of their ionization processes.

  • Each ionization step can exhibit varying degrees of acidity, emphasizing the differences between strong and weak acids in their ionization behavior.

pH Scale and Acid-Base Chemistry

• The pH scale is a logarithmic scale relating to the concentration of H⁺ in solution:

  • Formula: pH = -log[H⁺].

  • Pure water at 25°C has [H⁺] = 1 × 10^-7 mol/L, indicating a neutral pH of 7.

  • Acidic solutions have [H⁺] concentrations greater than 1 × 10^-7 mol/L (pH < 7).

  • Basic solutions have [H⁺] concentrations less than 1 × 10^-7 mol/L (pH > 7).

Self-Ionization of Water

• The reaction of water molecules with themselves can be represented as:

  • 2H₂O(l) ⇌ H₃O⁺(aq) + OH^-(aq).

• The ionization constant (K_w) of pure water at 25°C is:

  • K_w = [H⁺][OH^-] = 1 × 10^-14 mol²/L².

Acid-Base Properties in Aqueous Solutions

• Acidic solutions contain an excess of H⁺; thus, pH is less than 7.

• Basic solutions, conversely, contain an excess of OH^-; hence, pH exceeds 7.

• Neutral solutions are characterized by [H⁺] = [OH^-], corresponding to a pH of 7. This balance is critical for many biological processes.

Acidity of Salt Solutions

• The hydrolysis of salts significantly influences the acidity of a solution:

  • Basic salts, derived from weak acids, will produce OH^- in solution, making the solution basic.

  • Acid salts, derived from weak bases, release H⁺, resulting in an acidic solution.

Practical Applications of pH

• The use of indicators in titrations allows for visual tracking of pH changes as acids and bases are added, crucial for analytical chemistry.

• In biological systems, buffers stabilize pH levels to prevent acidosis or alkalosis, essential for maintaining physiological functions and overall homeostasis.

Buffers in Biological Systems

• Buffers consist of a weak acid and its conjugate base—or vice versa—to resist pH changes in solutions.

• A notable example is the bicarbonate buffer system in blood:

  • H₂CO₃(aq) ⇌ HCO₃^-(aq) + H⁺(aq).

  • This buffer system adjusts to changes in pH resulting from metabolic activities, illustrating how crucial buffers are for sustaining life.