Acids and bases are key players in chemistry. Brønsted-Lowry theory focuses on proton transfer, while Lewis theory looks at electron pair sharing. These concepts help us understand how molecules interact and react in various settings.

pH measures acidity, with 7 being neutral. Strong acids and bases fully break apart in water, while weak ones only partly do. This affects how they behave in reactions and their impact on living things.

Brønsted-Lowry Acids and Bases

Defining Acids and Bases in the Brønsted-Lowry Theory

• Brønsted-Lowry acid functions as a proton donor in chemical reactions
• Brønsted-Lowry base acts as a proton acceptor during chemical processes
• Proton donor transfers a hydrogen ion (H+) to another species in solution
• Proton acceptor receives a hydrogen ion (H+) from another species in solution
• Conjugate acid-base pairs form when an acid donates a proton to a base
  • Consists of the acid and its conjugate base, or the base and its conjugate acid
  • Differ by a single proton in their chemical formulas
• Amphoteric substances can act as both acids and bases depending on the reaction conditions
  • Includes water (H2O), which can donate or accept protons
  • Bicarbonate ion (HCO3-) serves as another common amphoteric species

Examples and Applications of Brønsted-Lowry Theory

• HCl (hydrochloric acid) donates a proton to H2O (water) in aqueous solution
  • HCl + H2O → H3O+ + Cl-
  • HCl acts as the acid, H2O as the base
  • H3O+ becomes the conjugate acid of H2O, Cl- the conjugate base of HCl
• NH3 (ammonia) accepts a proton from H2O in aqueous solution
  • NH3 + H2O ⇌ NH4+ + OH-
  • NH3 functions as the base, H2O as the acid
  • NH4+ forms the conjugate acid of NH3, OH- the conjugate base of H2O
• Amino acids demonstrate amphoteric behavior in biological systems
  • Can donate protons through their carboxyl groups (-COOH)
  • Can accept protons via their amino groups (-NH2)

Lewis Acids and Bases

Fundamentals of Lewis Acid-Base Theory

• Lewis acid behaves as an electron pair acceptor in chemical reactions
• Lewis base serves as an electron pair donor during chemical processes
• Electron pair acceptor receives a pair of electrons to form a covalent bond
• Electron pair donor provides a pair of electrons to form a covalent bond
• Hard and soft acids and bases (HSAB) principle categorizes Lewis acids and bases
  • Hard acids and bases tend to be small, highly charged, and weakly polarizable
  • Soft acids and bases are typically large, have low charge states, and are highly polarizable

Examples and Applications of Lewis Acid-Base Theory

• BF3 (boron trifluoride) acts as a Lewis acid by accepting an electron pair from NH3 (ammonia)
  • BF3 + NH3 → BF3NH3
  • BF3 functions as the Lewis acid, NH3 as the Lewis base
• Cu2+ ion serves as a Lewis acid in complex formation with water molecules
  • Cu2+ + 6H2O → [Cu(H2O)6]2+
  • Cu2+ accepts electron pairs from water molecules to form coordinate covalent bonds
• Carbocation (R3C+) behaves as a Lewis acid in organic reactions
  • Accepts an electron pair from nucleophiles in substitution or elimination reactions
• HSAB principle applied in coordination chemistry
  • Hard acids (Fe3+) prefer hard bases (F-)
  • Soft acids (Pt2+) favor soft bases (I-)

Acid-Base Strength and pH

Understanding Acid-Base Strength

• Acid-base strength measures the extent of dissociation or proton transfer in solution
• Strong acids and bases dissociate completely in aqueous solutions
  • Includes HCl (hydrochloric acid) and NaOH (sodium hydroxide)
• Weak acids and bases partially dissociate in aqueous solutions
  • Encompasses CH3COOH (acetic acid) and NH3 (ammonia)
• Acid strength inversely relates to the strength of its conjugate base
  • Strong acids have weak conjugate bases
  • Weak acids possess strong conjugate bases

pH Scale and Its Applications

• pH scale quantifies the acidity or basicity of a solution
• Ranges from 0 to 14 in aqueous solutions at 25°C
  • pH below 7 indicates acidic solutions
  • pH above 7 signifies basic solutions
  • pH of 7 represents neutral solutions
• Calculated using the negative logarithm of hydrogen ion concentration
  • pH = -log[H+]
• pOH scale measures hydroxide ion concentration
  • pOH = -log[OH-]
  • Relates to pH through the equation: pH + pOH = 14 (in water at 25°C)
• Buffer solutions resist pH changes when small amounts of acid or base are added
  • Consist of a weak acid and its conjugate base or a weak base and its conjugate acid
  • Maintain relatively constant pH in biological systems (blood pH ≈ 7.4)