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Acids, bases and pH Knowledge-Based Study

Biology · Chemistry of Life · Biology · Atoms, Water & Chemical Bonds

Biology topic explanations on Acids, bases and pH.

Study notes

  1. In pure liquid water, a small fraction of water molecules naturally dissociates into ions in a reversible process. When two water molecules interact, one can transfer a proton to another, forming a hydronium ion (H3O+H_3O^+, commonly referred to simply as a hydrogen ion, H+H^+) and a hydroxide ion (OH−OH^-):

    H2OightleftharpoonsH++OH−H_2O ightleftharpoons H^+ + OH^-

    In pure water at 25 °C, this dissociation produces an equal concentration of both ions:

    • [H+]=10−7 M[H^+] = 10^{-7}\text{ M}
    • [OH−]=10−7 M[OH^-] = 10^{-7}\text{ M}

    Because the concentrations of H+H^+ and OH−OH^- are exactly equal, pure water is considered chemically neutral.

    Solutes dissolved in water can alter the balance of hydrogen and hydroxide ions:

    • Acids: Substances that increase the relative concentration of hydrogen ions (H+H^+) in an aqueous solution, typically by donating protons. For example, hydrochloric acid (HClHCl) dissociates completely in water:
      HClightarrowH++Cl−HCl ightarrow H^+ + Cl^-, raising [H+][H^+] significantly above 10−7 M10^{-7}\text{ M}.
    • Bases: Substances that decrease the relative concentration of hydrogen ions, either by directly accepting H+H^+ ions or by releasing hydroxide ions (OH−OH^-) that combine with H+H^+ to form water. For instance, sodium hydroxide (NaOHNaOH) dissociates into Na+Na^+ and OH−OH^-, reducing the free H+H^+ concentration.

    The pH scale is a logarithmic measure of the concentration of free hydrogen ions in an aqueous solution. It is mathematically defined as:

    pH=−log⁡10[H+]\text{pH} = -\log_{10}[H^+]

    Key characteristics of the pH scale include:

    • Neutral solution: pH=7\text{pH} = 7 (where [H+]=10−7 M[H^+] = 10^{-7}\text{ M}).
    • Acidic solution: pH<7\text{pH} < 7 (higher [H+][H^+] than neutral water, e.g., gastric juice ≈1.5−2.0\approx 1.5 - 2.0).
    • Basic (alkaline) solution: pH>7\text{pH} > 7 (lower [H+][H^+] than neutral water, e.g., household bleach ≈12−13\approx 12 - 13).

    Because the scale is logarithmic (base 10), each whole pH unit represents a tenfold difference in hydrogen ion concentration. For example, a solution with a pH of 5 has ten times more H+H^+ ions than a solution with a pH of 6, and 100 times more H+H^+ ions than a solution with a pH of 7.

    Most biological fluids and cellular processes require a strictly regulated pH range. For example, human arterial blood must remain between 7.35 and 7.45. Significant shifts in internal pH can denature proteins, disrupt enzyme function, and compromise cellular membrane stability.

    Organisms regulate internal pH using buffers—substances that minimize changes in the concentrations of H+H^+ and OH−OH^- when an acid or base is added. A buffer typically consists of a weak acid and its corresponding conjugate base.

    In human blood, the carbonic acid–bicarbonate buffering system plays a primary role:

    H2CO3ightleftharpoonsHCO3−+H+H_2CO_3 ightleftharpoons HCO_3^- + H^+

    • If blood pH drops (excess H+H^+), bicarbonate ions (HCO3−HCO_3^-) bind free H+H^+ to form carbonic acid (H2CO3H_2CO_3).
    • If blood pH rises (depleted H+H^+), carbonic acid dissociates to release H+H^+ into the blood.