PH Buffer Solution

Buffer solution is as below:
a. Mixed of weak acid with salt from that weak acid substance

Example:
- CH₃COOH mixed CH₄COONa
- H₃PO₄  mixed NaH₂PO₄
b. Mixed of weak base with salt from that weak base substance.

Example
NH₄OH mixed NH₄ Cl
The properties of buffer solution:
- pH solution will not change if dissolved.
- pH solution will not changed if added a little of acid or base solution..

The way to calculate buffer solution PH:
To calculate pH of buffer solution of mixed weak acid with their salt (the solution will have pH < 7) using the formula as follows:
 [H+] = Ka. Ca/Cg
pH = pKa - log Ca/Cg = - log Ka + log Cg/Ca

where::
Ca = concentration of weak acid
Cg = concentration of the salt Ka = weak acid ionization constant

Example:
1. Calculate the pH solution for mixed of acetic acid 0.01 mol with solution sodium acetate 0.1 mol in 1 liter of water. Known Ka for sodium acetate = 10^-5 answer:

Ca = 0.01 mol/liter = 10^-2 M
Cg = 0.10 mol/liter = 10^-1 M
pH= pKa + log Cg/Ca = -log 10^-5 + log10^-1/log10^-2 = 5 + 1 = 6

 2. To calculate of buffer solution contain of mixed weak base with thier base salt (the solution will have pH  > 7), use the formula as follows:

[OH^-] = Kb . Cb/Cg
pOH = pKb + log Cg/Cb

where:
Cb = weak base concentration
Cg = salt concentration
Kb = constant of weak base ionization

Example:
Calculate mixed solution of 1 liter contain of NH₄OH 0.2 mol with HCl 0.1 mol (Kb= 10^-5)
answer:
NH₄OH(aq) + HCl(aq) <====> NH ₄Cl(aq) + H ₂ O(l)
mol NH₄OH reacted = mol HCl available = 0.1 mol
mol NH₄OH rest = 0.2 - 0.1 = 0.1 mol
mol NH₄Cl formed = mol NH ₄OH reacted = 0.1 mol

Because weak acid rest and mixed with salt (NH₄Cl), the mixture will formed buffer solution:
Cb (rest) = 0.1 mol/liter = 10^-1 M
Cg (formed) = 0.1 mol/liter = 10^-1 M
pOH = pKb + log Cg/Cb = -log 10 ^-5  + log (10 ^-1 /10 ^-1 ) = 5 + log 1 = 5
pH = 14 - pOH = 14 - 5 = 9

The Basicities of Metallic Elements and Their Compounds

Although the terms basicity, as applied to the metallic elements and their compounds, has been employed to cover a variety of phenomena, from the ease with which the free elements lose electrons under oxidizing conditions through the extent to which metal salts hydrolyze in aqueous solution to the ease with which oxygen-containing salts decompose when heated, all such phenomona, are reducible, directly or indirectly, to relative attractions (or lack of attraction) for aniona or electrons. In this respect, they are, therefore, all manifestations of acid base behavior in terms of the broad electronic concept offered by Usanovich.

Since basicity involves the loss of aniona or electrons, any property which measures the tendency of an element to low electrons or which measures the lack of attraction which an ion has for electrons or anions in turn measures the basicity of that elements or ion. There is, therefore, no ambiguity in referring to such a variety of phenomena as mentioned above as measure of basicities. A broad generalization of this sort would be impossible in term of older approaches to acid-base character, and in this respect the Usanovich interpretation does much to clarify an otherwise confused situation.

Solubility in Liquid Ammonia

Inasmuch as the solubilities of materials in liquid ammonia are often markedly different from the corresponding solubilities in water and inasmuch as the reaction solute undergo are often functions of their solubilities, a general summary of solubilities is desirable. Perhaps the outstanding difference between ammonia and water is the ability of ammonia to dissolve, without chemical reaction, five metals which are strongly reducing in character. Thus the alkali metals dissolve readily to yield characteristic blue solution from which the free metals can be recovered by evaporation of the solvent. The alkaline earth metal (calcium, strontium, and barium) shows similar behaviors, although their solubilities are not as large and evaporation of ammonia yields, as first solid phases, their unstable metal ammonates of composition M(NH3)6. Magnesium exhibits a slight, though measurable, solubility, as does aluminum, and the same is apparently true to lesser degree of beryllium, zinc, gallium, lanthanum, cerium and manganese.

Non metals such as iodine, sulfur, selenium and phosphorus are somewhat soluble in liquid ammonia. With sulfur, and perhaps with the other as well, this solubility is due, at least in part, to reaction with the solvent. The solubility of inorganic salts show trends markedly different from those noted in water. As might be expected, those salts which are most readily and extensively solvated are the most soluble. With ammonia the nature of the anion appear to play an even more important role than with water, the nature of the natural of the cation being comparatively unimportant except that a number of ammonium salts are soluble irrespective of anion. Among the halides solubility increses markedly from fluoride to iodide, essentially al fluorides being insoluble and even such iodides as that of silver being very soluble. The only chlorides which are really soluble are ammonium and beryllium chloride. Soluble inorganic salt ordinarily contain anions as iodide, perchlorate, nitrate, thiocyanate, cyanide, or nitrite, whereas salt containing carbonate, oxalate, sulfate, sulfite, sulfide, arsenate, phosphate, hydroxide, or oxide ions uniformly insoluble.