During exercise, the muscles use up oxygen as they convert chemical energy in glucose to mechanical energy. This O 2 comes from hemoglobin in the blood. These chemical changes, unless offset by other physiological functions, cause the pH of the blood to drop. If the pH of the body gets too low below 7. This can be very serious, because many of the chemical reactions that occur in the body, especially those involving proteins, are pH-dependent.
Ideally, the pH of the blood should be maintained at 7. If the pH drops below 6. Fortunately, we have buffers in the blood to protect against large changes in pH. All cells in the body continually exchange chemicals e. This external fluid, in turn, exchanges chemicals with the blood being pumped throughout the body. A dominant mode of exchange between these fluids cellular fluid, external fluid, and blood is diffusion through membrane channels, due to a concentration gradient associated with the contents of the fluids.
Recall your experience with concentration gradients in the "Membranes, Proteins, and Dialysis" experiment. Hence, the chemical composition of the blood and therefore of the external fluid is extremely important for the cell.
As mentioned above, maintaining the proper pH is critical for the chemical reactions that occur in the body. In order to maintain the proper chemical composition inside the cells, the chemical composition of the fluids outside the cells must be kept relatively constant.
This constancy is known in biology as homeostasis. This is a schematic diagram showing the flow of species across membranes between the cells, the extracellular fluid, and the blood in the capillaries.
The body has a wide array of mechanisms to maintain homeostasis in the blood and extracellular fluid. The most important way that the pH of the blood is kept relatively constant is by buffers dissolved in the blood.
Other organs help enhance the homeostatic function of the buffers. The kidneys help remove excess chemicals from the blood, as discussed in the Kidney Dialysis tutorial. Acidosis that results from failure of the kidneys to perform this excretory function is known as metabolic acidosis. However, excretion by the kidneys is a relatively slow process, and may take too long to prevent acute acidosis resulting from a sudden decrease in pH e.
The lungs provide a faster way to help control the pH of the blood. The increased-breathing response to exercise helps to counteract the pH-lowering effects of exercise by removing CO 2 , a component of the principal pH buffer in the blood. Acidosis that results from failure of the lungs to eliminate CO 2 as fast as it is produced is known as respiratory acidosis.
The kidneys and the lungs work together to help maintain a blood pH of 7. Therefore, to understand how these organs help control the pH of the blood, we must first discuss how buffers work in solution. Acid-base buffers confer resistance to a change in the pH of a solution when hydrogen ions protons or hydroxide ions are added or removed.
An acid-base buffer typically consists of a weak acid , and its conjugate base salt see Equations in the blue box, below. Buffers work because the concentrations of the weak acid and its salt are large compared to the amount of protons or hydroxide ions added or removed.
When protons are added to the solution from an external source, some of the base component of the buffer is converted to the weak-acid component thus using up most of the protons added ; when hydroxide ions are added to the solution or, equivalently, protons are removed from the solution; see Equations in the blue box, below , protons are dissociated from some of the weak-acid molecules of the buffer, converting them to the base of the buffer and thus replenishing most of the protons removed.
However, the change in acid and base concentrations is small relative to the amounts of these species present in solution. Hence, the ratio of acid to base changes only slightly. 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.
Hence, the conjugate base of an acid is the species formed after the acid loses a proton; the base can then gain another proton to return to the acid. In solution, these two species the acid and its conjugate base exist in equilibrium.
When an acid is placed in water, free protons are generated according to the general reaction shown in Equation 3. Note : HA and A - are generic symbols for an acid and its deprotonated form, the conjugate base. Hence, the equilibrium is often written as Equation 4, where H 2 O is the base :. Using the Law of Mass Action, which says that for a balanced chemical equation of the type.
Using the Law of Mass Action, we can also define an equilibrium constant for the acid dissociation equilibrium reaction in Equation 4. This equilibrium constant, known as K a , is defined by Equation The equilibrium constant for this dissociation reaction, known as K w , is given by.
H 2 O is not included in the equilibrium-constant expression because it is a pure liquid. 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:. Methods figures The ISE pH electrode and associated materials. Remove the ISE pH electrode from its bottle of buffer by unscrewing the cap a little so that you can easily pull out the electrode.
Make sure top portion of the bottle is removed. Replace screw cap of bottle if totally separated and place back in dry storage beaker where you found the electrode. Rinse the electrode with distilled water while holding the electrode over the large beaker used for the collection of waste liquids.
Place the tip of the ISE pH electrode in a small beaker containing enough room temperature distilled water to submerge the tip. Keep the electrode in distilled water for a few seconds. Obtain the calibration buffers: One beaker will contain pH 7 buffer and the other will contain pH 10 buffer.
Remove the electrode from the distilled water and gently blot any extra drops of water using Kimwipes. Submerge the end of the electrode in the beaker of pH 7 buffer. Click Record the red button on the upper toolbar on the LabScribe Main window to begin recording. After a few seconds, the trace will reach a stable baseline.
Type pH 7 in the Mark box to the right of the Mark button at the top- center of the screen. Press the Enter key on the keyboard to mark the stable baseline of the recording.
You may continue recording while changing the beakers of buffers. Hold the electrode over the beaker used for collecting waste liquid and rinse it with distilled water.
Blot any extra drops of water with Kimwipe. Position the electrode in the beaker of pH 10 buffer. As you continue to record, the trace will reach a stable baseline, allow it to record for several seconds. Type pH 10 in the Mark box. Click Stop on the LabScribe Main toolbar to stop recording.
Remove the electrode from the beaker of pH 10 buffer. Hold the electrode over the beaker used for collecting waste liquid, and rinse it with distilled water from a wash bottle.
Blot any extra drops of water and place the electrode in a clean beaker of room temperature distilled water. To show the data collected at pH 7 and pH 10 on the Main window at the same time, you can use either the Double Time Display icon to adjust the Display Time. Click Double Time several times until your pH 7 mark and the pH 10 mark are on the same screen or the Double Cursor icon will allow two red cursors to appear on the Main window.
Place one red cursor on a flat section of data collected when the ISE pH electrode was in the pH 7 buffer and the second red cursor on a flat section of data collected when the electrode was in the pH 10 buffer. Click on the Auto Scale icon in the tool bar above the graph screen.
Methods figure 4. Note gray triangle in orange circle. To convert the voltages at the positions of the cursors to pH values, use the Simple Units Conversion dialogue window. To access this dialogue window, click on the grey arrow the upper left corner of the screen, see orange circle on figure 4, above to the left of the channel title, pH, to open the channel menu.
Select Units from the channel menu , and select Simple from the units submenu. On the Units Conversion window, see picture below make sure 2 point calibration is selected in the pull-down menu in the upper-left corner of the window.
Put a check mark in the box next to Apply units to all blocks. Notice that the voltages from the positions of the cursors are automatically entered into the value equations. Enter the two buffers used in the calibration recording in the corresponding boxes on the right side of the conversion equations. Enter the name of the units, pH, in the box below the buffer values. Click on the OK button in the lower right corner of the window to activate the Units Conversion.
Methods figure 5. The process that causes the imbalance is classified based on the etiology of the disturbance respiratory or metabolic and the direction of change in pH acidosis or alkalosis. There are four basic processes and one or a combination may occur at any given time. Expiration : When blood pH drops too low, the body compensates by increasing breathing to expel more carbon dioxide. A smaller fraction is transported in the red blood cells that combine with the globin portion of hemoglobin as carbaminohemoglobin.
This is the chemical portion of the red blood cell that aids in the transport of oxygen and nutrients around the body, but, this time, it is carbon dioxide that is transported back to the lung. The basic reaction governed by this principle is as follows:. When the blood pH drops too low acidemia , the body compensates by increasing breathing to expel more CO 2 ; this shifts the above reaction to the left such that less hydrogen ions are free; thus, the pH will rise back to normal.
For alkalemia, the opposite occurs. The kidneys help maintain the acid—base balance by excreting hydrogen ions into the urine and reabsorbing bicarbonate from the urine. Urine testing is important because it can detect acid—base imbalances. For instance, uncontrolled diabetes results in highly acidic urine. If the diabetes remains uncontrolled, the kidneys could become over-stressed and malfunction, which could lead to coma or death.
Within the human body, fluids such as blood must be maintained within the narrow range of 7. Outside that range, pH becomes incompatible with life; proteins are denatured and digested, enzymes lose their ability to function, and the body is unable to sustain itself.
To maintain this narrow range of pH the body has a powerful buffering system. Acid—base imbalances that overcome this system are compensated in the short term by changing the rate of ventilation. The kidneys are slower to compensate than the lungs, but renal physiology has several powerful mechanisms to control pH by the excretion of excess acid or base.
The major, homeostatic control point for maintaining a stable pH balance is renal excretion. In response to acidosis, the tubular cells reabsorb more bicarbonate from the tubular fluid, and the collecting duct cells secrete more hydrogen and generate more bicarbonate, and ammoniagenesis leads to an increase in the formation of the NH 3 buffer.
In response to alkalosis, the kidneys may excrete more bicarbonate by decreasing hydrogen ion secretion from the tubular epithelial cells, and lowering the rates of glutamine metabolism and ammonium excretion. Privacy Policy. Skip to main content. Body Fluids and Acid-Base Balance. Search for:. Acid-Base Balance. Learning Objectives Explain the composition of buffer solutions and how they maintain a steady pH.
Key Terms alkaline : having a pH greater than 7; basic acidic : having a pH less than 7 buffer : a solution composed of a weak acid and its conjugate base that can be used to stabilize the pH of a solution.
0コメント