Buffers in the human body are essential to maintain normal metabolism. The internal pH in most cells is around 7, but small variations can be fatal within minutes and natural ways of buffering cellular pH are vital. Specifically, in human body, the kidneys and the lungs work together to help maintain a blood pH of 7.4 by affecting the components of the buffers in the blood.
Important buffer systems in the body include bicarbonate buffers, phosphate buffer and protein buffers.
In fact, there are many different factors that play a role in choosing a particular buffer system for a particular biological application. These factors include working temperature, desired pH, toxicity to the system, and interactions with other components.
1. Bicarbonate buffers
By far the most important buffer for maintaining acid-base balance in the blood is the carbonic-acid-bicarbonate buffer. In this buffer, carbonic acid (H2CO3) act as a weak acid and hydrogen carbonate ion (HCO3–) act as conjugate base of a weak acid or salt of weak acid.
H2CO3(aq) 
CO2(aq) + H2O(l)
An enzyme called carbonic anhydrase catalyzes the conversion of carbonic acid to dissolved carbon dioxide. In the lungs, excess dissolved carbon dioxide is exhaled as carbon dioxide gas.
CO2(aq) CO2(g)
The concentration of hydrogen carbonate ions is controlled through the kidneys. Excess hydrogen carbonate ions are excreted in the urine.
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The much higher concentration of hydrogen carbonate ion over that of carbonic acid in blood plasma allows the buffer to respond effectively to the most common materials that are released into the blood. Specifically, normal metabolism releases mainly acidic materials: carboxylic acids such as lactic acid (HLac). These acids react with hydrogen carbonate ion and form carbonic acid.
HLac(aq) + HCO3–(aq) Lac–(aq) + H2CO3(aq)
The carbonic acid is converted through the action of the enzyme carbonic anhydrase into aqueous carbon dioxide.
H2CO3(aq) CO2(aq) + H2O(l)
An increase in CO2(aq) concentration stimulates increased breathing, and the excess carbon dioxide is released into the air.
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The condition called respiratory acidosis occurs when blood pH falls as a result of decreased respiration. When respiration is restricted, the concentration of dissolved carbon dioxide in the blood increases, making the blood too acidic. Such a condition can be produced by asthma, pneumonia, emphysema, or inhaling smoke.
Metabolic acidosis is the decrease in blood pH that results when excessive amounts of acidic substances are released into the blood. This can happen through prolonged physical exertion, by diabetes, or restricted food intake. The normal body response to this condition is increases breathing to reduce the amount of dissolved carbon dioxide in the blood. This is why we breathe more heavily after climbing several flights of stairs.
Respiratory alkalosis results from excessive breathing that produces an increase in blood pH. Hyperventilation causes too much dissolved carbon dioxide to be removed from the blood, which decreases the carbonic acid concentration, which raises the blood pH. Often, the body of a hyperventilating person will react by fainting, which slows the breathing.
Metabolic alkalosis is an increase in blood pH resulting from the release of alkaline materials into the blood. This can result from the ingestion of alkaline materials, and through overuse of diuretics. Again, the body usually responds to this condition by slowing breathing, possibly through fainting.
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In conclusion, the carbonic acid-hydrogen carbonate ion buffer works throughout the body to maintain the pH of blood plasma close to 7.40. The body maintains the buffer by eliminating either the acid (carbonic acid) or the base (hydrogen carbonate ions). Changes in carbonic acid concentration can be effected within seconds through increased or decreased respiration. Changes in hydrogen carbonate ion concentration, however, require hours through the relatively slow elimination through the kidneys.
2. Phosphate buffer (Buffering of internal cell fluids)
The phosphate buffer system operates in the internal fluid of all cells. This buffer system consists of dihydrogen phosphate ions (H2PO4–) as hydrogen-ion donor (acid) and hydrogen phosphate ions (HPO42-) or phosphate ions (PO43-), as the hydrogen-ion acceptor (base). These two ions are in equilibrium with each other as indicated by the chemical equation below.
H+(aq) + HPO42-(aq)H+(aq) + PO43-(aq)In the absence of this phosphate buffer, sharp changes in pH of cell fluids may cause cell death or improper working of different proteins and cell organelles present within the cell.
3. Phosphate Protein buffer (Buffering in Cells and Tissues)
Proteins are mainly composed of amino acids. These amino acids contain functional groups – including carboxyl (COOH) and amino groups (NH2) – that can act as weak acid and bases with buffering capacities. This seems to be a particularly important system in cells and tissues, but researchers believe it may also have an impact even in blood.
The figure below shows the molecular structure of an amino acid:
The carboxyl group
The carboxyl group is attached to the amino acid central carbon: C-COOH. In the figure above, you can see the carboxyl group off to the left. You can see that the carboxyl group consists of a double bond to one of the oxygen and a single bond to the hydroxyl group. The important part of the carboxyl group that is to be noted is the hydrogen atom within the hydroxyl group.
At a near neutral pH, like the pH of blood, the carboxyl group is actually COO– instead of COOH. Then, if a protein finds itself in a more acidic solution, the carboxyl group will be able to take on the extra hydrogen ions and return to the COOH configuration.
The amino group
The amino group is attached to the amino acid central carbon: C – NH2. The amino group is shown at the right hand side of the diagram of the amino acid above. However, at a near neutral pH, like in blood, the amino group is actually NH3+ rather than just NH2. It actually tends to carry an extra hydrogen ion on it at a normal pH. Then, if a protein finds itself in a more basic environment, its amino groups on its amino acids can actually release their hydrogen ions and return to NH2.
Diagrammatic Overview of amino acids providing buffering functions:
So, amino acids can accept or donate hydrogen ions, making them excellent buffers. And any given protein typically has hundreds of amino acids. So, proteins make superb buffers. Remember, they are found in very high concentration in intracellular solutions and in blood.
As all cells and tissues are composed of proteins mainly so in the absence of protein buffer, the sharp changes in pH may cause cell death or tissue damage in living organisms.
Other implications to society:
(Click on each of these to dwell into the specific implication of buffers to society)
1. Buffers and Industrial Applications
2. Buffers in Nature
References:
- http://environmental-realm.blogspot.sg/2012/04/importance-of-buffers-in-physiological.html
- http://www.thechemicalblog.co.uk/buffers-in-our-daily-life/
- http://scifun.chem.wisc.edu/chemweek/biobuff/biobuffers.html
- http://faculty.stcc.edu/AandP/AP/AP2pages/Units21to23/ph/buffers.htm
- http://www.thechemicalblog.co.uk/what-is-a-buffer-solution/
Pictures:
- http://www.thechemicalblog.co.uk/what-is-a-buffer-solution/
- https://fmss12uchemd.files.wordpress.com/2013/05/screen-shot-2013-05-07-at-5-04-55-pm.png
- http://faculty.stcc.edu/AandP/AP/AP2pages/Units21to23/ph/buffers.htm