Understanding acid-base equilibria is crucial in chemistry, especially in biological systems. This topic covers definitions, pH scales, strengths, and the role of buffers, all of which help maintain homeostasis in living organisms.
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Definition of acids and bases (Arrhenius, Brønsted-Lowry, Lewis)
- Arrhenius: Acids produce H⁺ ions in solution, while bases produce OH⁻ ions.
- Brønsted-Lowry: Acids are proton donors, and bases are proton acceptors.
- Lewis: Acids are electron pair acceptors, and bases are electron pair donors.
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pH scale and pOH
- pH measures the concentration of H⁺ ions in a solution; lower pH indicates higher acidity.
- pOH measures the concentration of OH⁻ ions; lower pOH indicates higher basicity.
- The relationship: pH + pOH = 14 at 25°C.
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Acid and base strength
- Strong acids/bases completely dissociate in water, while weak acids/bases only partially dissociate.
- Strength is determined by the degree of ionization and the stability of the resulting ions.
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Ka and Kb constants
- Ka (acid dissociation constant) quantifies the strength of an acid in solution.
- Kb (base dissociation constant) quantifies the strength of a base in solution.
- Larger Ka or Kb values indicate stronger acids or bases, respectively.
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Autoionization of water and Kw
- Water undergoes autoionization to produce H⁺ and OH⁻ ions.
- Kw is the ion product constant of water (1.0 x 10⁻¹⁴ at 25°C), representing the product of [H⁺] and [OH⁻].
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Conjugate acid-base pairs
- A conjugate acid is formed when a base gains a proton; a conjugate base is formed when an acid loses a proton.
- The strength of a conjugate acid-base pair is inversely related; strong acids have weak conjugate bases and vice versa.
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Buffer solutions and their mechanisms
- Buffers resist changes in pH upon the addition of small amounts of acids or bases.
- They typically consist of a weak acid and its conjugate base or a weak base and its conjugate acid.
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Henderson-Hasselbalch equation
- pH = pKa + log([A⁻]/[HA]), where [A⁻] is the concentration of the conjugate base and [HA] is the concentration of the weak acid.
- Useful for calculating the pH of buffer solutions.
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Acid-base titrations
- A technique to determine the concentration of an acid or base by neutralizing it with a base or acid of known concentration.
- The endpoint is often indicated by a color change from an acid-base indicator.
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Polyprotic acids and bases
- Polyprotic acids can donate more than one proton (e.g., H₂SO₄, H₃PO₄).
- Each dissociation step has its own Ka value, with the first dissociation being the strongest.
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Amino acids and their acid-base properties
- Amino acids can act as both acids and bases due to the presence of both amino (NH₂) and carboxyl (COOH) functional groups.
- The isoelectric point (pI) is the pH at which the amino acid has no net charge.
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Biological buffer systems (e.g., bicarbonate buffer)
- Biological systems use buffers to maintain pH homeostasis; bicarbonate (HCO₃⁻) is a key buffer in blood.
- Buffers work by neutralizing excess acids or bases through reversible reactions.
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Le Châtelier's principle in acid-base equilibria
- States that if a system at equilibrium is disturbed, the system will shift to counteract the disturbance.
- In acid-base equilibria, adding or removing H⁺ or OH⁻ will shift the equilibrium position.
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Common ion effect
- The presence of a common ion in solution decreases the solubility of a salt due to Le Châtelier's principle.
- This effect can also influence the dissociation of weak acids and bases.
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Acid-base indicators
- Indicators are substances that change color at a specific pH range, signaling the endpoint of a titration.
- Common indicators include phenolphthalein (colorless in acid, pink in base) and bromothymol blue (yellow in acid, blue in base).