Strong and Weak Acids and Bases, and Buffers

[Bernard NebelKeymaster]

As we have seen, any chemical formula (non carbon, inorganic) with an H at the front or an organic (carbon containing) compound with an COOH at the  end is an acid; it will tend to release hydrogen atoms into solution as hydrogen ions (H⁺). 

However, this release of H⁺ into solution is not an all or none process. Think of it as the compound having a certain internal electric pressure to pop off their H⁺. At the same time however, the concentration of H⁺ already in solution, the pH if you will, creates counter pressure. There is a pressure among these to pop onto whatever is available. 

The result is that there are only few acids that release 100% of their hydrogens ions to water solution as H⁺ ions. These are known as strong acids;  hydrochloric, sulfuric, and nitric acids (HCl, H₂SO₄, and HNO₃). are most important. 

Most acids, especially those found in food stuffs, vinegar, and other common materials, are weak acids. They will only release H⁺ so long as its concentration in the surrounding solution is less than a certain amount. If the H⁺ concentration is greater than that amount, H⁺ will be pushed back to its place of bonding. 

Putting this in terms of pH, adding a weak acid to a solution will only take pH down to a certain level; at lower pH (greater H⁺ concentration) the weak acid will not ionize (will not release H⁺). 

Exactly the same reasoning follows for strong and weak bases.

The following table gives the pH value where some common substances  (weak acids and bases) reach their “balance point” of releasing/taking back H⁺. For bases it is a matter of releasing/taking back OH^– ions. (Exactly the same reasoning applies to strong and weak bases.)

pH values for some common substances TABLE

http://coolperiodictable.com/resources/acids-and-bases/pH-of-some-common-substances.php

 

Buffers

 Neutralizing an acid does not always require a base. There are non base (non OH-) compounds that will react with and thus absorb H⁺ ions. Hence, they prevent the water from becoming more acidic. Such compounds are called buffers. 

The most important example, both agriculturally and environmentally is the naturally occurring mineral, lime stone, which is chemically calcium carbonate, CaCO₃. Acid reacts with limestone as shown below.

 CaCO₃          +       2H⁺ Cl-       ——->         Ca Cl₂      +      CO₂   +   H₂O 

calcium            hydrochloric                     calcium          Carbon       water           carbonate                acid                             chloride          dioxide                           limestone                                                                   no acid

Limestone (ground to fine granules) is periodically spread on agricultural fields to keep pH close to neutral. A near neutral pH of ponds and lakes is also maintained by adding limestone. (See video below)

 

Where limestone occurs naturally, it may provide natural buffering as shown in the video below.

 

 

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