14.6 Buffers

A buffer that consists of a weak acid and its salt is an example of a buffer that resists a change in pH.

The unbuffered solution has become acidic due to the change in color of the orange, which is red at a pH of 4.

It is a buffer because it contains both acid and salt.
Adding either a strong acid or a strong base keeps the hydronium ion concentration almost constant.
Adding a base such as sodium hydroxide will cause the hydroxide ion to react with the few hydronium ion present.

To prevent pH changes that might change the biochemical activity of compounds, acetic acid is used.

A solution with a volume of 101 mL is given by adding NaOH to 100 mL of this buffer.

The concentrations of the intermediate mixture are calculated from the complete reaction between the acid in the buffer and the added base.

When the same amount of NaOH was added to the buffered solution, the change was 4.7%.

If we add an acid or a base to a buffer that is a mixture of a weak base and its salt, the calculations of the changes in pH are similar to those for a buffer mixture of a weak acid and its salt.

If we add so much base to a buffer that the weak acid is exhausted, no more buffering action is possible.

If we added more acid, the weak base would be exhausted, and no more buffering action would be possible.
We don't need to exhaust all of the acid or base in a buffer to overwhelm it; its buffering action will diminish as a given component nears depletion.

The indicator color shows that a small amount of acid added to a buffered solution of pH 8 has little effect on the buffered system.

A large amount of acid exhausts the buffering capacity of the solution and the pH changes dramatically.

The buffer capacity is dependent on the amount of the weak acid and its conjugate base.

When one component of the buffer pair is less than 10%, a buffer solution loses its usefulness.

Weak acids and their salts are better as buffers for pHs less than 7 than weak bases and salts.

The concentrations of the weak acid and its salt in a buffered solution are related to this equation.

Lawrence Joseph Henderson was an American physician, biochemist and physiologist.

After obtaining a medical degree from Harvard, he spent 2 years studying in Strasbourg, then a part of Germany, before taking a lecturer position at Harvard.
He worked at Harvard for the rest of his life.
The acid-base balance in human blood is regulated by a buffer system formed by dissolved carbon dioxide.
The carbonic acid-carbonate buffer system in blood was described in an equation written in 1908.
In addition to his important research on the physiology of blood, Henderson wrote on the adaptation of organisms and their fit with their environments, on sociology and on university education.
He founded the Fatigue Laboratory at the Harvard Business School, which examined human physiology with a specific focus on work in industry, exercise, and nutrition.

The ability of the blood to bind with oxygen was shown to be related to the acidity of the water in it.

The pH scale was introduced in 1909 by another Danes, Sorensen, and in 1912 by another Danes, Hasselbalch.
The Henderson-Hasselbalch equation was born in 1916 after Hasselbalch expressed Henderson's equation in logarithmic terms.

The 2CO3 concentration is lower than the HCO3 ion due to the fact that most of the by-products of our metabolism are acidic.

The capacity of the buffer will not be exceeded if there is a larger proportion of base than acid.