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An acidity function is a measure of the acidity of a medium or solvent system, [1] [2] usually expressed in terms of its ability to donate protons to (or accept protons from) a solute (Brønsted acidity). The pH scale is by far the most commonly used acidity function, and is ideal for dilute aqueous solutions.
A smaller H + concentration means a greater OH − concentration and, therefore, a greater K b and a greater pH. NaOH (s) (sodium hydroxide) is a stronger base than (CH 3 CH 2) 2 NH (l) (diethylamine) which is a stronger base than NH 3 (g) (ammonia). As the bases get weaker, the smaller the K b values become. [1]
Lawrence Joseph Henderson was a biological chemist and Karl Albert Hasselbalch was a physiologist who studied pH. [2] [3] In 1908, Lawrence Joseph Henderson [4] derived an equation to calculate the hydrogen ion concentration of a bicarbonate buffer solution, which rearranged looks like this:
The relative concentration of undissociated acid is shown in blue, and of its conjugate base in red. The pH changes relatively slowly in the buffer region, pH = pK a ± 1, centered at pH = 4.7, where [HA] = [A −]. The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in ...
The pH of a solution is defined as the negative logarithm of the concentration of H +, and the pOH is defined as the negative logarithm of the concentration of OH −. For example, the pH of a 0.01 in moles per litreM solution of hydrochloric acid (HCl) is equal to 2 (pH = −log 10 (0.01)), while the pOH of a 0.01 M solution of sodium ...
Conversely, when pH = pK a, the concentration of HA is equal to the concentration of A −. The buffer region extends over the approximate range pK a ± 2. Buffering is weak outside the range pK a ± 1. At pH ≤ pK a − 2 the substance is said to be fully protonated and at pH ≥ pK a + 2 it is fully dissociated (deprotonated).
The Charlot equation, named after Gaston Charlot, is used in analytical chemistry to relate the hydrogen ion concentration, and therefore the pH, with the formal analytical concentration of an acid and its conjugate base. It can be used for computing the pH of buffer solutions when the approximations of the Henderson–Hasselbalch equation ...
3] + [CO 2− 3]. K 1, K 2 and DIC each have units of a concentration, e.g. mol/L. A Bjerrum plot is obtained by using these three equations to plot these three species against pH = −log 10 [H +] eq, for given K 1, K 2 and DIC. The fractions in these equations give the three species' relative proportions, and so if DIC is unknown, or the ...