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1. Enumerate the normal Ph of important body fluids (eg. Saliva, tears, blood, etc.).

The normal pH of blood running through arteries (large elastic-walled blood vessels that carry blood from theheart to other parts of the body) is 7.4; the pH of blood in the veins (vessels that transports blood to the heart) is

about 7.35. Normal urine pH averages about 6.0. Saliva has a pH between 6.0 and 7.4. Semen = 7.5

Blood = 7.4

Gastric Juice = 0.7

Urine = 6.0

Pancreatic Juice = 8.1

Cerebrospinal Fluid = 7.3

Saliva = 7.0

2. Give the importance of knowing pH of these fluids in the body.

Essentially every function of the body is dependent on our bodies maintaining a precisely balanced pH in the

 blood, and other critical bodily systems. We must remember that the changes in pH do not have to be major.

Most elements of the body have a very definite pH range and slight changes impact chemical reactions both

within and outside the cell. This is why maintenance of pH levels are critical to overall body balance

3. Give 5 examples each of foods that make the body fluids acidic and basic.

Acid forming foods include asparagus, barley, beans (dried), beechnuts, and beef.

Alkaline forming foods include agar, alfalfa (sprouts), almonds, APPLES(apple cider), andapricots.

4. Explain the carbonic acid buffer system. Why is it considered the most important buffer in our 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:

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(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 thePeriodic 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 labmanual). 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. Howeverthe 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=K a/K 2 (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, K 1, for the left-hand

reaction in Equation 10, using the Law of Mass Action:

. (13)

K a (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 K a is

.(14

)

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The equilibrium constant, K 2, 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. Solvingfor the equilibrium concentration of carbonic acid gives

.(16

)

Rearranging Equation 16 allows us to solve for the equilibrium protonconcentration 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 therelation 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

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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), andthe 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 thecapillaries very quickly, and so there is little time in the lungs for carbon dioxide to be exchanged for oxygen.

5. Define the following conditions and give the probable causes of each as well as their possible treatments or 

management.

a. Acidosis is a condition in which there is excessive acid in the body fluids. The kidneysand lungs maintain the balance (proper pH level) of chemicals called acids and bases inthe body. Acidosis occurs when acid builds up or when bicarbonate (a base) is lost.

Acidosis is classified as either respiratory acidosis or metabolic acidosis. Respiratoryacidosis develops when there is too much carbon dioxide (an acid) in the body. This typeof acidosis is usually caused by a decreased ability to remove carbon dioxide from thebody through effective breathing. Other names for respiratory acidosis are hypercapnicacidosis and carbon dioxide acidosis. Causes of respiratory acidosis include Chest

deformities, such as kyphosis, Chest injuries, Chest muscle weakness, Chronic lung disease, Overuse of sedative drugs. Metabolic acidosis develops when too much acid is produced or when the kidneys cannot

remove enough acid from the body.

 b. Alkalosis is a condition in which the body fluids have excess base (alkali). The kidneys and lungs maintainthe proper balance of chemicals, called acids and bases, in the body. Decreased carbon dioxide (an acid) or 

increased bicarbonate (a base) levels make the body too alkaline, a condition called alkalosis. Respiratory

alkalosis is caused by low carbon dioxide levels in the blood. This can be due to Fever, Being at a high altitude,Lack of oxygen, Liver disease, Lung disease, which causes you to breathe faster (hyperventilate), Salicylate

 poisoning. Metabolic alkalosis is caused by too much bicarbonate in the blood. Hypochloremic alkalosis is

caused by an extreme lack or loss of chloride, which can occur with prolonged vomiting. Hypokalemic alkalosis

is caused by the kidneys' response to an extreme lack or loss of potassium, which can occur when people takecertain diuretic medications. Compensated alkalosis occurs when the body returns the acid - base balance to

normal in cases of alkalosis, but bicarbonate and carbon dioxide levels remain abnormal. Symptoms are

Confusion (can progress to stupor or coma), Hand tremor, Lightheadedness, Muscle twitching, Nausea,

vomiting, Numbness or tingling in the face or extremities, Prolonged muscle spasms (tetany). Treatment of alkalosis depends on finding the specific cause. For alkalosis caused by hyperventilation, breathing into a paper 

 bag causes you to retain more carbon dioxide and improves the alkalosis. If your oxygen level is low, you mayreceive oxygen to help the alkalosis. Some people need medications to correct chemical loss (such as chloride

and potassium). Your health care provider will monitor your vital signs (temperature, pulse, rate of breathing,

 blood pressure).

.

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