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Acids & Bases - pH - Buffers

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Svante Arrhenius Definition of Acids/Bases

acid - a compound that increases concentration of H+/protons when in aqueous solution (or H3O/hydronium when H+ bonds with H2O)
base - a compound that increases concentration of OH-/hydroxide when in aqueous solution

example strong acid - HCl

HCl(aq) + H2O(l) -> H+(aq) + CL-(aq)

example strong base - NaOH

NaOH(aq) + H2O(l) -> Na+(aq) + OH-(aq)

Brønsted-Lowry Definition of Acids/Bases

acid - H+/proton donor

base - H+/proton acceptor

example strong acid - HCl

HCl(aq) + H2O(l) -> H+(aq) + CL-(aq)

example strong base - NaOH

NaOH(aq) + H2O(l) -> Na+(aq) + OH-(aq)

Lewis Definition of Acids/Bases

TODO - article

pH Definition

pH = -log(hydrogen protons concentration)
pH = -log(hydronium concentration)

Buffers

Most organisms, including humans, need to maintain pH within a fairly narrow range in order to survive. For instance, human blood needs to keep its pH right around 7.4

Buffers - solutions that can resist changes in pH, are key to maintaining stable H+ ion concentrations in biological systems.

  • when there are too many H+ ions, a buffer will absorb some of them, bringing pH back up
  • when there are too few H+ ions, a buffer will donate some of its own H+ions to reduce the pH

Buffers typically consist of an acid-base pair, with the acid and base differing by the presence or absence of a proton (a conjugate acid-base pair). For instance, one of the buffers that maintain the pH of human blood involves carbonic acid (H2CO3) and its conjugate base, the bicarbonate ion (HCO3-).

  • if too many H+ ions build up, the equation above will be pushed to the right, and bicarbonate ions will absorb the H+ to form carbonic acid
  • if H+ concentrations drop too low, the equation will be pulled the left and carbonic acid will turn into bicarbonate, donating H+ ions to the solution