chm 12901 general chemistry with a biological focus unit 6 study guides

Key Concepts

• Acids donate protons (H⁺) in aqueous solutions while bases accept protons
• Brønsted-Lowry theory defines acids as proton donors and bases as proton acceptors
• Arrhenius theory states acids dissociate in water to form H⁺ ions and bases dissociate to form OH⁻ ions
• pH is a logarithmic scale that measures the concentration of H⁺ ions in a solution
  • Ranges from 0 to 14 with 7 being neutral, below 7 acidic, and above 7 basic
• pOH measures the concentration of OH⁻ ions in a solution and is related to pH by the equation $pH + pOH = 14$
• Acid-base reactions involve the transfer of protons between reactants
• Buffers are solutions that resist changes in pH when small amounts of acid or base are added

Acid-Base Theories

• Arrhenius theory was the first to describe acids and bases in terms of their behavior in aqueous solutions
  • Limited to substances that contain H⁺ or OH⁻ ions
• Brønsted-Lowry theory expands the definition of acids and bases to include any species that can donate or accept protons
  • Allows for the classification of substances like ammonia (NH₃) as a base
• Lewis theory further broadens the definition of acids and bases based on electron pair interactions
  • Acids are electron pair acceptors and bases are electron pair donors
• Conjugate acid-base pairs consist of a species and its corresponding proton-transferred form (HA and A⁻)
• Amphoteric substances can act as both acids and bases depending on the environment (water, amino acids)

Properties of Acids and Bases

• Acids have a sour taste (citric acid in lemons, acetic acid in vinegar)
• Bases have a bitter taste and feel slippery (soap, baking soda)
• Acids react with metals to produce hydrogen gas (H₂)
  • Hydrochloric acid (HCl) reacting with zinc (Zn) to form zinc chloride (ZnCl₂) and H₂
• Bases react with oils and fats to form soaps in a process called saponification
• Acids change the color of pH indicators like litmus paper (red) and phenolphthalein (colorless)
• Bases change the color of pH indicators like litmus paper (blue) and phenolphthalein (pink)

pH and pOH

• pH is calculated using the negative logarithm of the hydrogen ion concentration: $pH = -log[H⁺]$
• pOH is calculated using the negative logarithm of the hydroxide ion concentration: $pOH = -log[OH⁻]$
• The relationship between pH and pOH is given by $pH + pOH = 14$ at 25°C
• A neutral solution has a pH of 7 and equal concentrations of H⁺ and OH⁻ ions ($[H⁺] = [OH⁻] = 10^{-7} M$)
• Acidic solutions have a pH below 7 and a higher concentration of H⁺ ions compared to OH⁻ ions
• Basic solutions have a pH above 7 and a higher concentration of OH⁻ ions compared to H⁺ ions
• The pH scale is logarithmic, meaning a change of one pH unit represents a tenfold change in H⁺ concentration

Strength of Acids and Bases

• Acid and base strength refers to the extent of dissociation in aqueous solutions
• Strong acids and bases completely dissociate in water (HCl, NaOH)
  • Have a higher degree of ionization and conductivity
• Weak acids and bases only partially dissociate in water (acetic acid, ammonia)
  • Exist in equilibrium with their dissociated ions
• The acid dissociation constant (Ka) and base dissociation constant (Kb) quantify the strength of weak acids and bases
  • Higher values indicate a stronger acid or base
• The relationship between Ka and Kb for a conjugate acid-base pair is given by $Ka × Kb = Kw$, where $Kw$ is the ionization constant of water ($1.0 × 10^{-14}$ at 25°C)

Acid-Base Reactions

• Neutralization reactions occur when an acid and a base react to form water and a salt
  • Hydrochloric acid (HCl) and sodium hydroxide (NaOH) form water and sodium chloride (NaCl)
• Titration is a technique used to determine the concentration of an unknown acid or base solution
  • Involves the gradual addition of a known concentration of acid or base until the endpoint is reached
• Indicators like phenolphthalein are used to visually detect the endpoint of a titration
• Acid-base reactions can also involve the transfer of protons between molecules (NH₃ + H₂O ⇌ NH₄⁺ + OH⁻)
• The net ionic equation for an acid-base reaction shows only the species that participate in the proton transfer

Buffers and Their Biological Importance

• Buffers are solutions that resist changes in pH 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 a relatively constant pH through the equilibrium between the acid and base components
• The Henderson-Hasselbalch equation relates the pH of a buffer to the pKa of the acid and the concentrations of the acid and base components: $pH = pKa + log([base]/[acid])$
• Biological systems rely on buffers to maintain homeostasis and proper cellular function
  • Blood is buffered by the bicarbonate buffer system (H₂CO₃/HCO₃⁻) to maintain a pH of 7.4
• Buffers are important in regulating the pH of enzymatic reactions, as enzymes function optimally within a specific pH range

Applications in Biochemistry

• Amino acids, the building blocks of proteins, contain both acidic (carboxyl) and basic (amino) functional groups
  • Act as buffers and help maintain the pH of biological systems
• The isoelectric point (pI) of an amino acid is the pH at which it has a net neutral charge
  • Determines the solubility and mobility of amino acids in different pH environments
• Enzymes, biological catalysts, are sensitive to changes in pH
  • Optimal pH range is essential for proper enzyme function and cellular processes
• Nucleic acids (DNA and RNA) are acidic due to the presence of phosphate groups in their backbone
• The pH of bodily fluids is tightly regulated to ensure proper physiological function
  • Stomach acid (HCl) has a pH of 1.5-3.5 to aid in digestion and protect against pathogens
  • Blood plasma has a pH of 7.4, maintained by the bicarbonate buffer system and respiratory regulation of CO₂