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:. The equilibrium on the left is an acid-base reaction that is written in the reverse format from Equation 3. Carbonic acid H 2 CO 3 is the acid and water is the base.
Carbonic acid also dissociates rapidly to produce water and carbon dioxide, as shown in the equilibrium on the right of Equation 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. 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.
However, the relationship shown in Equation 11 is frequently referred to as the Henderson-Hasselbach equation for the buffer in physiological applications. In such cases, bicarbonate ions are not conserved from the filtrate to the blood, which will also contribute to a pH imbalance and acidosis. The hydrogen ions also compete with potassium to exchange with sodium in the renal tubules.
If more potassium is present than normal, potassium, rather than the hydrogen ions, will be exchanged, and increased potassium enters the filtrate. When this occurs, fewer hydrogen ions in the filtrate participate in the conversion of bicarbonate into CO 2 and less bicarbonate is conserved. If there is less potassium, more hydrogen ions enter the filtrate to be exchanged with sodium and more bicarbonate is conserved. Chloride ions are important in neutralizing positive ion charges in the body.
If chloride is lost, the body uses bicarbonate ions in place of the lost chloride ions. Thus, lost chloride results in an increased reabsorption of bicarbonate by the renal system. Acid-Base Balance: KetoacidosisDiabetic acidosis, or ketoacidosis, occurs most frequently in people with poorly controlled diabetes mellitus.
When certain tissues in the body cannot get adequate amounts of glucose, they depend on the breakdown of fatty acids for energy. When acetyl groups break off the fatty acid chains, the acetyl groups then non-enzymatically combine to form ketone bodies, acetoacetic acid, beta-hydroxybutyric acid, and acetone, all of which increase the acidity of the blood.
Ketoacidosis can be severe and, if not detected and treated properly, can lead to diabetic coma, which can be fatal. A common early symptom of ketoacidosis is deep, rapid breathing as the body attempts to drive off CO 2 and compensate for the acidosis. Another common symptom is fruity-smelling breath, due to the exhalation of acetone. Other symptoms include dry skin and mouth, a flushed face, nausea, vomiting, and stomach pain. Treatment for diabetic coma is ingestion or injection of sugar; its prevention is the proper daily administration of insulin.
A person who is diabetic and uses insulin can initiate ketoacidosis if a dose of insulin is missed. Among people with type 2 diabetes, those of Hispanic and African-American descent are more likely to go into ketoacidosis than those of other ethnic backgrounds, although the reason for this is unknown. A variety of buffering systems exist in the body that helps maintain the pH of the blood and other fluids within a narrow range—between pH 7. A buffer is a substance that prevents a radical change in fluid pH by absorbing excess hydrogen or hydroxyl ions.
Several substances serve as buffers in the body, including cell and plasma proteins, hemoglobin, phosphates, bicarbonate ions, and carbonic acid. The bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body. The respiratory and renal systems also play major roles in acid-base homeostasis by removing CO 2 and hydrogen ions, respectively, from the body.
Skip to content Learning Objectives By the end of this section, you will be able to: Identify the most powerful buffer system in the body Identify the most rapid buffer system in the body Describe the protein buffer systems.
Explain the way in which the respiratory system affects blood pH Describe how the kidney affects acid-base balance. Chapter Review A variety of buffering systems exist in the body that helps maintain the pH of the blood and other fluids within a narrow range—between pH 7. Review Questions. Critical Thinking Questions 1. Describe the conservation of bicarbonate ions in the renal system. Describe the control of blood carbonic acid levels through the respiratory system.
Glossary hypercapnia abnormally elevated blood levels of CO 2 hypocapnia abnormally low blood levels of CO 2. Solutions Answers for Critical Thinking Questions Bicarbonate ions are freely filtered through the glomerulus. They cannot pass freely into the renal tubular cells and must be converted into CO 2 in the filtrate, which can pass through the cell membrane. Sodium ions are reabsorbed at the membrane, and hydrogen ions are expelled into the filtrate. Article Navigation.
Close mobile search navigation Article Navigation. Volume 4. Buffer Capacities of Human Blood and Plasma. George Ellison , George Ellison. Department of Biochemistry. Oxford Academic. Google Scholar. Jon V Straumfjord, Jr. J P Hummel. Select Format Select format. Permissions Icon Permissions. Abstract Normal values for buffer capacity of blood and plasma of a number of normal subjects over various intervals between pH 7.
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