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Exercise causes an imbalance of pH within the body; therefore our body has several mechanisms to counterbalance these fluctuations. They are called Buffers. They work to keep the body at a constant pH of around 7.4, more easily said than done. When exercising our metabolism speeds up and produces more CO2 and H+ ions in the muscles. We breathe more to keep up with the oxygen demand but eventually are unable to compensate and lactic acid is produced in the muscles by anaerobic respiration. Buffers kick in to stabilise the pH. Some organs that help in maintaining a balance is the Heart which at high rates hinders CO2 removal, Kidneys remove HCO3- and lungs remove CO2. As well as these organs, there are other pH buffer systems one prominent example is the Bicarbonate-Buffer System. These all work together to maintain homeostasis in the blood.
The most important buffer in the blood stream is the carbonic-acid-bicarbonate buffer. It is the top acid-base balancer in the blood.
The simultaneous equilibrium reactions of concern are
The pH of the blood depends on the concentration of H+, (HCO3-, H2CO3, and CO2) molecules to form equilibrium constant.
According to the universal Brønstead-Lowry definition, Acids donate protons or H+ molecules and a base accepts these.
The definition of pH:
[H+] is the molar concentration of protons or Hydrogen atoms in a solution. When an acid is placed into water, free protons are produced according to equation 3.
Equation 3 complies with the Brønstead-Lowry definition (The HA (acid) gives away a proton to become A- while the water acts like a base, accepting the proton.
The H+ is actually shared by many water molecules and the equation for this equilibrium is written as:
Where H2O is the base of this equation.
The Law of Mass Action and the so called Equilibrium Constants.
Using this Law of Mass Action; this states that for a balanced (both sides are equal) chemical equation of the nature;
A, B, C, and D are elements and a, b, c, and d are then their stoichiometric coefficients, a equilibrium constant (K) can then be found from the equation:
[A], [B], [C], and [D] specify the concentration of the elements at equilibrium.
The Equilibrium Constant for an Acid-Base Reaction
In equation 4 (HA equation), the acid dissociation equilibrium reaction can formulate an equilibrium constant. This equilibrium constant, known as Ka, is described by equation 7;
The Equilibria Constant for the Dissociation of Water.
A use of the Law of Mass Action can be show in the dissociation of H2O into H+ and OH-;
Kw is equilibrium constant for this dissociation reaction, given by;
Increasing OH- concentrations in aqueous solutions reduces the H+ concentration as the product of the two concentrations at a given temperature has to remain constant. Therefore the equilibrium in equation 8 means the equivalency of the Brønstead-Lowry classification and the Arrhenius description of a base are feasible.
To better illustrate the two equilibrium reactions of the carbonic-acid-bicarbonate buffer (The most important one in the blood), equation 1 is slightly modified to clearly illustrate the contribution of water:
[H30+(aq) + HCO3-(aq) â‡Œ H2CO3(aq) + H2O(l)] Is the acid-base reaction in equation 10. It is simply written in the reverse format of equation 3. Water is the base while H2CO3 (Carbonic acid) is the acid. The conjugate base of H2CO3 is, HCO3- (bicarbonate ion). Carbonic acid quickly dissipates to produce more water and carbon dioxide (as shown on right side of equation 10) This isn't an acid-base reaction however is crucial to the blood's buffering capacity as shown below in equation 11 below:
The source of equation is shown from equations 13-18 below.
Observe that Equation 11 is comparable to the Henderson-Hasselbach equation but does not meet the strict description. This is because equation 11 incorporates a non-acid-base reaction; the dissociation of carbonic acid to CO2 and H2O as well as the ratio in parentheses is not actually the concentration ratio of the acid to its conjugate base. Rather, the Henderson-Hasselbach equation for all the buffers in physiological applications is commonly referred to the association shown in equation 11 above. For the buffer, pK is equivalent to the negative log of the equilibrium constant K, as seen in equation 12 below;
Where K=Ka/K2 is from equation 10.
This quantity presents a suggestion of the extent of which HCO3- reacts with H+ (originally H3O+ as shown in equation 10) to form H2CO3, and consequently forms CO2 and H2O. In this situation of the carbonic-acid-bicarbonate buffer, pK=6.1 at the average body temperature.
The Source of the pH Equation for the Carbonic-Acid-Bicarbonate Buffer.
Use the Law of Mass Action Rule to describe the equilibrium constant, K1 is used, for the left-hand reaction in the equation 10;
The switched version of the left-hand reaction in the equation 10 is Ka, the equilibrium constant for this acid-base reaction observed in equation 7 above. The formula for Ka is;
The Law of Mass Action is also described by the right-hand reaction in equation 10, the equilibrium constant, K2, as shown below;
Equations 14 and 15 can be written as two simultaneous equations. This is because of the two equilibrium reactions in equation 10 above occur concurrently. When you solve for the specific equilibrium concentration of carbonic acid, you get the formula;
To solve for the equilibrium proton concentration, equation 16 has to be rearranged. In terms of the two equilibrium constants (Ka and K2) and the concentrations of the other two species in the equation;
Using negative log on both sides of equation 17 will give us the pH of the blood;
To give the relation revealed in equation 11, equation 18 can be modified using more conventional notation, which is replicated below;
Above is the titration curve of the carbonic-acid-bicarbonate buffer.
The formula for percentage buffer in the form of HCO3- is known as;
1) How does exercise have an effect on the body?
A) Exercise is promoted physical activity. It comprises of chemical and physical factors that work together in order to produce physical movement. Exercise has many beneficial factors which is why many people today make it a priory to get at least 30 minutes of exercise a day. In order to move, muscles contract. To do this, muscles need energy (ATP - Adenosine Triphosphate) and oxygen (If anaerobic respiration is activated). ATP liberates energy for work by a reaction which releases one of the phosphate-oxygen groups, leaving ADP (Adenosine Diphosphate). These molecules are then transferred to the mitochondria which 'refills' the ADP molecules to return as ATP molecules ready for use once again. A by product of respiration is Carbon Dioxide and Lactic Acid. These substances decrease the pH of the blood; therefore it is important that these substances are removed as swiftly as possible. In order to do this buffers are used.
2) How are chemicals exchanged in the body?
A) Chemicals are constantly being exchanged with the external fluid throughout the body by all cells. These chemicals include; nutrients, waste products and ions. The external fluid then exchanges with the capillaries. The foremost mode used is diffusion via the concentration gradient. This means that the external fluid needs to maintain a stable pH otherwise H+ ions will enter the cell and will eventually destroy it. Thus the external fluid must retain a suitable chemical composition appropriate for the surrounding cells. This consistency is revered to as Homeostasis. The body has several mechanisms in order to maintain homeostasis; by far the most important are the buffers. As well as the buffers the body has organs to aid in the maintenance of the body's pH. One of these organs is the kidneys, which remove H+ ions and other chemicals that affect the pH negatively. Another organ which aids in the maintenance of the pH is the lungs. The lungs remove CO2 which greatly affects the pH in its large quantities.
3) How do buffers work?
A) Buffers are a 'resistance' to change in pH when either hydrogen or hydroxide ions are added to the mix. The buffers typically contain weak acid and its conjugate base. When H+ ions are added to the equation, the base reacts with it to form weak acid. When OH- is added, it reacts with the weak acid to form a conjugate base and water, only raising the pH slightly. These buffers work because they are in very large concentration therefore when H+ or OH- ions are added there is very little change in concentration in the solution hence the ratio changes only slightly.
4) How is Le Chatelier's Principle related to maintaining equilibrium in the blood?
A) Le Chatelier's Principle states that "if a change in conditions [(an external) 'stress'] is imposed on a system at equilibrium, the equilibrium position will shift in a direction that tends to reduce that change in conditions." Thus when a product or reactant of an equilibrium reaction is then added to another solution at equilibrium, the concentration of the species will change so a new equilibrium is formed. This is recognized as a shift of the equilibrium. An example of this relating to the pH of the blood is with the kidneys. When the pH of the blood is rising, the kidneys remove HCO3-, thus the equilibrium shifts to the left, producing more H+ ion and HCO3- ions thus the pH declines.