ASPIRIN OVERDOSE

Una sobredosis es la ingestion excesiva de cualquier sustancia fuera de la dosis predeterminada.

Como todos sabemos, la aspirina es un analgesico debil, aunque efectivo, pero se lo caracteriza de esta forma por su mecanismo de acción en comparasión con los mecanismos de acción de los analgesicos fuertes como la morfina.

Aunque sabemos que la aspirina es muy beneficiosa de algun modo, por ejemplo el efecto analgésico, anticoagulante y tambien bloquean los receptores que activan la inflamación luego de una lesión.

The Carbonic-Acid-Bicarbonate Buffer in the Blood

By far the most important buffer for maintaining acid-base balance in the blood is the carbonic-acid-bicarbonate buffer. The simultaneous equilibrium reactions of interest are
.
(1)
We are interested in the change in the pH of the blood; therefore, we want an expression for the concentration of H+ in terms of an equilibrium constant (see blue box, below) and the concentrations of the other species in the reaction (HCO3-, H2CO3, and CO2).
Review of acid-base concepts
To more clearly show the two equilibrium reactions in the carbonic-acid-bicarbonate buffer, Equation 1 is rewritten to show the direct involvement of water:
(10)
The equilibrium on the left is an acid-base reaction that is written in the reverse format from Equation 3. Carbonic acid (H2CO3) is the acid and water is the base. The conjugate base for H2CO3 is HCO3- (bicarbonate ion). (Note: To view the three-dimensional structure of HCO3-, consult the Table of Common Ions in the Periodic Properties tutorial from Chem 151.) Carbonic acid also dissociates rapidly to produce water and carbon dioxide, as shown in the equilibrium on the right of Equation 10. This second process is not an acid-base reaction, but it is important to the blood's buffering capacity, as we can see from Equation 11, below.
.
(11)
The derivation for this equation is shown in the yellow box, below. Notice that Equation 11 is in a similar form to the Henderson-Hasselbach equation presented in the introduction to the Experiment (Equation 16 in the lab manual). Equation 11 does not meet the strict definition of a Henderson-Hasselbach equation, because this equation takes into account a non-acid-base reaction (i.e., the dissociation of carbonic acid to carbon dioxide and water), and the ratio in parentheses is not the concentration ratio of the acid to the conjugate base. However, the relationship shown in Equation 11 is frequently referred to as the Henderson-Hasselbach equation for the buffer in physiological applications.
In Equation 11, pK is equal to the negative log of the equilibrium constant, K, for the buffer (Equation 12).
where K=Ka/K2 (from Equation 10).
(12)
This quantity provides an indication of the degree to which HCO3- reacts with H+ (or with H3O+ as written in Equation 10) to form H2CO3, and subsequently to form CO2 and H2O. In the case of the carbonic-acid-bicarbonate buffer, pK=6.1 at normal body temperature.
Derivation of the pH Equation for the Carbonic-Acid-Bicarbonate Buffer
We may begin by defining the equilibrium constant, K1, for the left-hand reaction in Equation 10, using the Law of Mass Action:
.
(13)
Ka (see Equation 9, above) is the equilibrium constant for the acid-base reaction that is the reverse of the left-hand reaction in Equation 10. It follows that the formula for Ka is
.
(14)
The equilibrium constant, K2, for the right-hand reaction in Equation 10 is also defined by the Law of Mass Action:
.
(15)
Because the two equilibrium reactions in Equation 10 occur simultaneously, Equations 14 and 15 can be treated as two simultaneous equations. Solving for the equilibrium concentration of carbonic acid gives
.
(16)
Rearranging Equation 16 allows us to solve for the equilibrium proton concentration in terms of the two equilibrium constants and the concentrations of the other species:
.
(17)
Because we are interested in the pH of the blood, we take the negative log of both sides of Equation 17:
,
(18)
Recalling the definitions of pH and pK (Equations 2 and 12, above), Equation 18 can be rewritten using more conventional notation, to give the relation shown in Equation 11, which is reproduced below:
As shown in Equation 11, the pH of the buffered solution (i.e., the blood) is dependent only on the ratio of the amount of CO2 present in the blood to the amount of HCO3-(bicarbonate ion) present in the blood (at a given temperature, so that pK remains constant). This ratio remains relatively constant, because the concentrations of both buffer components (HCO3- and CO2) are very large, compared to the amount of H+ added to the blood during normal activities and moderate exercise. When H+ is added to the blood as a result of metabolic processes, the amount of HCO3- (relative to the amount of CO2) decreases; however, the amount of the change is tiny compared to the amount of HCO3- present in the blood. This optimal buffering occurs when the pH is within approximately 1 pH unit from the pK value for the buffering system, i.e., when the pH is between 5.1 and 7.1.
However, the normal blood pH of 7.4 is outside the optimal buffering range; therefore, the addition of protons to the blood due to strenuous exercise may be too great for the buffer alone to effectively control the pH of the blood. When this happens, other organs must help control the amounts of CO2 and HCO3- in the blood. The lungs remove excess CO2 from the blood (helping to raise the pH via shifts in the equilibria in Equation 10), and the kidneys remove excess HCO3- from the body (helping to lower the pH). The lungs' removal of CO2 from the blood is somewhat impeded during exercise when the heart rate is very rapid; the blood is pumped through the capillaries very quickly, and so there is little time in the lungs for carbon dioxide to be exchanged for oxygen. The ways in which these three organs help to control the blood pH through the bicarbonate buffer system are highlighted in Figure 3, below.