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The molar ionic strength, I, of a solution is a function of the concentration of all ions present in that solution. [3]= = where one half is because we are including both cations and anions, c i is the molar concentration of ion i (M, mol/L), z i is the charge number of that ion, and the sum is taken over all ions in the solution.
This "salting in" is called the "uncommon-ion effect" (also "salt effect" or the "diverse-ion effect"). It occurs because as the total ion concentration increases, inter-ion attraction within the solution can become an important factor. [3] This alternate equilibrium makes the ions less available for the precipitation reaction.
Molar concentration or molarity is most commonly expressed in units of moles of solute per litre of solution. [1] For use in broader applications, it is defined as amount of substance of solute per unit volume of solution, or per unit volume available to the species, represented by lowercase : [2]
The relative activity of a species i, denoted a i, is defined [4] [5] as: = where μ i is the (molar) chemical potential of the species i under the conditions of interest, μ o i is the (molar) chemical potential of that species under some defined set of standard conditions, R is the gas constant, T is the thermodynamic temperature and e is the exponential constant.
A metal ion in aqueous solution or aqua ion is a cation, dissolved in water, of chemical formula [M(H 2 O) n] z+.The solvation number, n, determined by a variety of experimental methods is 4 for Li + and Be 2+ and 6 for most elements in periods 3 and 4 of the periodic table.
The molar concentration of hydronium or H + ions determines a solution's pH according to pH = -log([H 3 O +]/M) where M = mol/L. The concentration of hydroxide ions analogously determines a solution's pOH. The molecules in pure water auto-dissociate into aqueous protons and hydroxide ions in the following equilibrium: H 2 O ⇌ OH − (aq) + H ...
Ammonia solutions decrease in density as the concentration of dissolved ammonia increases. At 15.6 °C (60.1 °F), the density of a saturated solution is 0.88 g/ml; it contains 35.6% ammonia by mass, 308 grams of ammonia per litre of solution, and has a molarity of approximately 18 mol/L.
The concentration of ions in a solution of a weak electrolyte is less than the concentration of the electrolyte itself. For acids and bases the concentrations can be calculated when the value or values of the acid dissociation constant are known.