Summary of Water Chemistry: Acid-Base and Equilibrium

Reactions in Water

  • Water molecules continually undergo self-reaction; bonds are broken and new bonds form.
  • Reactions are written with reactants on the left and products on the right; atoms must balance so the number of atoms of each element is the same on both sides.
  • Arrows indicate reaction direction:
    • Single arrow ((\rightarrow)) = irreversible reaction
    • Double arrows ((\rightleftharpoons)) = reversible reaction
  • Balance principle: Number of reactant atoms=Number of product atoms\text{Number of reactant atoms} = \text{Number of product atoms}

Writing Chemical Equations

  • Arrows represent the direction of the reaction; reversible vs irreversible.
  • Example balance concept: \ce{2 H2 + O2 -> 2 H2O} (illustrates atoms conserved across sides)
  • For reversible: \ce{N2 + 3 H2

Dissociation of Water and Hydronium

  • The self-ionization (dissociation) of water: 2 H<em>2O⇌H</em>3O++OH−2\,\mathrm{H<em>2O} \rightleftharpoons \mathrm{H</em>3O^+} + \mathrm{OH^-}
  • Hydronium representation: protons associate with water to form (\mathrm{H_3O^+}).
  • In pure water, the concentration of hydrogen ions equals the concentration of hydroxide ions: [H+]=[OH−][\mathrm{H^+}] = [\mathrm{OH^-}] (in reality, protons exist as hydronium: (\mathrm{H^+}\ in solution is (\mathrm{H_3O^+})).
  • A proton donor fixes: acids; a proton acceptor fixes: bases.

Acids, Bases, and Proton Transfer

  • Acids donate protons (increase (\mathrm{H^+}) or (\mathrm{H_3O^+}) in solution).
  • Bases accept protons (increase (\mathrm{OH^-}) or reduce (\mathrm{H^+}) concentration).
  • Water is amphoteric: can act as both an acid and a base.
  • A chemical reaction involving proton transfer is an acid–base reaction: requires a proton donor (acid) and a proton acceptor (base).
  • Definition restated: substances that give up protons are acids; substances that acquire protons are bases.

Water as an Acid and Base; Weakness/Strength

  • Water is an extremely weak acid and a weak base (only a small fraction dissociates at any moment).
  • Strong acids dissociate readily in water to form hydronium ions: HCl+H<em>2O→H</em>3O++Cl−\mathrm{HCl + H<em>2O \rightarrow H</em>3O^+ + Cl^-}
  • Strong bases dissociate completely in water to form hydroxide and cation: NaOH<em>(aq)→Na+</em>(aq)+OH(aq)−\mathrm{NaOH<em>{(aq)} \rightarrow Na^+</em>{(aq)} + OH^-_{(aq)}}
  • Note: in water, many acids/bases mostly act in a single role, but water itself can act as both acid and base.

Acid–Base Equilibria: Key Takeaways

  • Acid–base reactions involve proton donors and acceptors.
  • Self-ionization of water establishes a defined but small concentration of (\mathrm{H_3O^+}) and (\mathrm{OH^-}) in pure water.
  • Equilibrium implies forward and backward reactions occur at equal rates, yielding constant concentrations.
  • Equilibrium dynamics are essential for understanding pH and buffer behavior in biological systems.