Acid-Base Theories and Reactions in Chemistry

• Acid-Base Theories and Acid-Base Reactions are a consideration in Analytical Chemistry.

Arrhenius Acid-Base Definition

• Classifies acids and bases based on formulas and behavior in water.
• Arrhenius considers acids and bases as ionic compounds.

Arrhenius Acid

• An Arrhenius Acid is a substance with H in its formula that yields H3O+ or H+ when dissociated.
• Examples of Arrhenius Acids include HCl, HNO3, and HCN.

Arrhenius Base

• An Arrhenius Base is a substance with OH in its formula that yields OH- when dissociated.
• Examples of Arrhenius Bases include NaOH, KOH, and Ba(OH)2.

Neutralization

• Neutralization is a reaction where H+ ions from an acid and OH- ions from a base react to form water (H₂O).
• The ions "cancel each other out" in simple terms.
• Example of Neutralization: HCl (acid) + NaOH (base) yields NaCl (salt) + H₂O (water).

Brønsted-Lowry Acid-Base Definition

• The Brønsted-Lowry definition expands on the Arrhenius definition.
• It removes the restriction of water as a solvent, focusing on the exchange of H+ ions.

Brønsted-Lowry Acid

• A Brønsted-Lowry Acid is any substance that donates H+ ions (proton donor).
• Examples of Brønsted-Lowry acids include all Arrhenius acids like HCl, HNO3
• HCN (Hydrogen cyanide) also donates H+.

Brønsted-Lowry Base

• A Brønsted-Lowry Base is any substance that accepts a H+ ion (proton acceptor).
• Examples of Brønsted-Lowry bases include all Arrhenius bases like NaOH, KOH.
• NH3 (Ammonia) is a base because it accepts H+ ions.

Arrhenius vs Brønsted-Lowry

• Arrhenius definition states acids release H+ ions in water, while bases release OH- ions in water.
• Brønsted-Lowry definition states acids donate H+ ions (proton donor) and bases accept H+ ions (proton acceptor).
• Examples of Arrhenius acids: HCl, HNO3, HCN.
• Examples of Arrhenius bases: NaOH, KOH, Ba(OH)2.
• Examples of Brønsted-Lowry acids: HCl, HNO3, HCN, H2SO4.
• Examples of Brønsted-Lowry bases: NaOH, KOH, NH3, NaF.
• Arrhenius limitations: only in water-based solutions; some substances don't fit
• Brønsted-Lowry definition works in aqueous, non-aqueous, and gas phases.

Maintenance of H+

• An acid is a substance that can donate hydrogen ions (H+) when dissolved in water.
• A base is a substance that can accept hydrogen ions.
• pH is defined as the negative log of the hydrogen ion concentration.
• The logarithmic pH scale expresses H+ concentration (c is concentration) as pH = log(1/cH+) = -log(cH+).
• Through metabolism, the body produces approximately 150 g of H+ each day.
• H+ concentration in extracellular body fluids is maintained in a narrow range: 36 to 44 nmol/L (pH 7.35 to 7.45).
• The relative strengths of acids and bases are described by their dissociation constant (Ka).
• pKa is defined as the negative log of the dissociation constant.

Buffer Systems

• A buffer consists of a weak acid and a salt of its conjugate base.
• A buffer resists changes in pH upon adding acid or base.
• An example is H₂CO₃ ↔ HCO₃⁻ + H⁺ (Carbonic acid ↔ Bicarbonate).
• The Henderson-Hasselbalch equation expresses acid-base relationships in a mathematical formula: pH = pKa + log(cA/cHA).

Buffer System in Extracellular Fluids

• H2CO3 dissociates into CO2 and H2O, allowing CO2 to be eliminated by the lungs and H+ as water.
• Changes in CO2 modify the ventilation (respiratory) rate.
• HCO3-concentration can be changed by the kidneys.