Acid–base homeostasis

Acid–base homeostasis is the homeostatic regulation of the pH of the body's extracellular fluid (ECF).[1] The proper balance between the acids and bases (i.e. the pH) in the ECF is crucial for the normal physiology of the body, and cellular metabolism.[1] The pH of the intracellular fluid and the extracellular fluid need to be maintained at a constant level.[2]

Many extracellular proteins such as the plasma proteins and membrane proteins of the body's cells are very sensitive for their three dimensional structures to the extracellular pH.[3][4] Stringent mechanisms therefore exist to maintain the pH within very narrow limits. Outside the acceptable range of pH, proteins are denatured (i.e. their 3-D structure is disrupted), causing enzymes and ion channels (among others) to malfunction.

In humans and many other animals, acid–base homeostasis is maintained by multiple mechanisms involved in three lines of defence[5][6]:

  • The first line of defence are the various chemical buffers which minimize pH changes that would otherwise occur in their absence. They do not correct pH deviations, but only serve to reduce the extent of the change that would otherwise occur. These buffers include the bicarbonate buffer system, the phosphate buffer system, and the protein buffer system.
  • The second line of defence of the pH of the ECF consists of controlling of the carbonic acid concentration in the ECF. This is achieved by changes in the rate and depth of breathing (i.e. by hyperventilation or hypoventilation), which blows off or retains carbon dioxide (and thus carbonic acid) in the blood plasma.[5][7]
  • The third line of defence is the renal system, which can add or remove bicarbonate ions to or from the ECF.[5] The bicarbonate is derived from metabolic carbon dioxide which is enzymatically converted to carbonic acid in the renal tubular cells.[5][8][9] The carbonic acid spontaneously dissociates into hydrogen ions and bicarbonate ions.[5] When the pH in the ECF tends to fall (i.e. become more acidic) the hydrogen ions are excreted into the urine, while the bicarbonate ions are secreted into the blood plasma, causing the plasma pH to rise (correcting the initial fall).[10] The converse happens if the pH in the ECF tends to rise: the bicarbonate ions are then excreted into the urine and the hydrogen ions into the blood plasma.

Physiological corrective measures make up the second and third lines of defence. This is because they operate by making changes to the buffers, each of which consists of two components: a weak acid and its conjugate base.[5][11] It is the ratio concentration of the weak acid to its conjugate base that determines the pH of the solution.[12] Thus, by manipulating firstly the concentration of the weak acid, and secondly that of its conjugate base, the pH of the extracellular fluid (ECF) can be adjusted very accurately to the correct value. The bicarbonate buffer, consisting of a mixture of carbonic acid (H2CO3) and a bicarbonate (HCO
3
) salt in solution, is the most abundant buffer in the extracellular fluid, and it is also the buffer whose acid to base ratio can be changed very easily and rapidly.[13]

An acid–base imbalance is known as acidaemia when the acidity is high, or alkalaemia when the acidity is low.

Acid–base balance

The pH of the extracellular fluid, including the blood plasma, is normally tightly regulated between 7.32 and 7.42,[14] by the chemical buffers, the respiratory system, and the renal system.[11][15][16][17]

Aqueous buffer solutions will react with strong acids or strong bases by absorbing excess hydrogen H+
ions, or hydroxide OH
ions, replacing the strong acids and bases with weak acids and weak bases.[11] This has the effect of damping the effect of pH changes, or reducing the pH change that would otherwise have occurred. But buffers cannot correct abnormal pH levels in a solution, be that solution in a test tube or in the extracellular fluid. Buffers typically consist of a pair of compounds in solution, one of which is a weak acid and the other a weak base.[11] The most abundant buffer in the ECF consists of a solution of carbonic acid (H2CO3), and the bicarbonate (HCO
3
) salt of, usually, sodium (Na+).[5] Thus, when there is an excess of OH
ions in the solution carbonic acid partially neutralizes them by forming H2O and bicarbonate (HCO
3
) ions.[5][13] Similarly an excess of H+ ions is partially neutralized by the bicarbonate component of the buffer solution to form carbonic acid (H2CO3), which, because it is a weak acid, remains largely in the undissociated form, releasing far fewer H+ ions into the solution than the original strong acid would have done.[5]

The pH of a buffer solution depends solely on the ratio of the molar concentrations of the weak acid to the weak base. The higher the concentration of the weak acid in the solution (compared to the weak base) the lower the resulting pH of the solution. Similarly, if the weak base predominates the higher the resulting pH.

This principle is exploited to regulate the pH of the extracellular fluids (rather than just buffering the pH). For the carbonic acid-bicarbonate buffer, a molar ratio of weak acid to weak base of 1:20 produces a pH of 7.4; and vice versa - when the pH of the extracellular fluids is 7.4 then the ratio of carbonic acid to bicarbonate ions in that fluid is 1:20.[12]

This relationship is described mathematically by the Henderson–Hasselbalch equation, which, when applied to the carbonic acid-bicarbonate buffer system in the extracellular fluids, states that:[12]

where:
  • pH is the negative logarithm (or cologarithm) of molar concentration of hydrogen ions in the ECF. It indicates the acidity in the ECF in an inverse manner: the lower the pH the greater the acidity of the solution.
  • pKa H2CO3 is the cologarithm of the acid dissociation constant of carbonic acid. It is equal to 6.1.
  • [HCO
    3
    ]
    is the molar concentration of bicarbonate in the blood plasma
  • [H2CO3] is the molar concentration of carbonic acid in the ECF.
However, since the carbonic acid concentration is directly proportional to the partial pressure of carbon dioxide () in the extracellular fluid, the equation can be rewritten as follows:[5][12]
where:
  • pH is the negative logarithm of molar concentration of hydrogen ions in the ECF, as before.
  • [HCO
    3
    ]
    is the molar concentration of bicarbonate in the plasma
  • PCO2 is the partial pressure of carbon dioxide in the blood plasma.

The pH of the extracellular fluids can thus be controlled by separately regulating the partial pressure of carbon dioxide (which determines the carbonic acid concentration), and the bicarbonate ion concentration in the extracellular fluids.

There are therefore at least two homeostatic negative feedback systems responsible for the regulation of the plasma pH. The first is the homeostatic control of the blood partial pressure of carbon dioxide, which determines the carbonic acid concentration in the plasma, and can change the pH of the arterial plasma within a few seconds.[5] The partial pressure of carbon dioxide in the arterial blood is monitored by the central chemoreceptors of the medulla oblongata, and so are part of the central nervous system.[5][18] These chemoreceptors are sensitive to the pH and levels of carbon dioxide in the cerebrospinal fluid.[12][10][18] (The peripheral chemoreceptors are located in the aortic bodies and carotid bodies adjacent to the arch of the aorta and to the bifurcation of the carotid arteries, respectively.[18] These chemoreceptors are sensitive primarily to changes in the partial pressure of oxygen in the arterial blood and are therefore not directly involved with pH homeostasis.[18])

The central chemoreceptors send their information to the respiratory centres in the medulla oblongata and pons of the brainstem.[10] The respiratory centres then determine the average rate of ventilation of the alveoli of the lungs, to keep the partial pressure carbon dioxide in the arterial blood constant. The respiratory center does so via motor neurons which activate the muscles of respiration (in particular the diaphragm).[5][19] A rise in the partial pressure of carbon dioxide in the arterial blood plasma above 5.3 kPa (40 mmHg) reflexly causes an increase in the rate and depth of breathing. Normal breathing is resumed when the partial pressure of carbon dioxide has returned to 5.3 kPa.[7] The converse happens if the partial pressure of carbon dioxide falls below the normal range. Breathing may be temporally halted, or slowed down to allow carbon dioxide to accumulate once more in the lungs and arterial blood.

The sensor for the plasma HCO
3
concentration is not known for certain. It is very probable that the renal tubular cells of the distal convoluted tubules are themselves sensitive to the pH of the plasma. The metabolism of these cells produces CO2, which is rapidly converted to H+ and HCO
3
through the action of carbonic anhydrase.[5][8][9] When the extracellular fluids tend towards acidity, the renal tubular cells secrete the H+ ions into the tubular fluid from where they exit the body via the urine. The HCO
3
ions are simultaneously secreted into the blood plasma, thus raising the bicarbonate ion concentration in the plasma, lowering the carbonic acid/bicarbonate ion ratio, and consequently raising the pH of the plasma.[5][10] The converse happens when the plasma pH rises above normal: bicarbonate ions are excreted into the urine, and hydrogen ions into the plasma. These combine with the bicarbonate ions in the plasma to form carbonic acid (H+ + HCO
3
= H2CO3), thus raising the carbonic acid:bicarbonate ratio in the extracellular fluids, and returning its pH to normal.[5]

In general, metabolism produces more waste acids than bases.[5] The urine is therefore generally acid. This urinary acidity is, to a certain extent, neutralized by the ammonia (NH3) which is excreted into the urine when glutamate and glutamine (carriers of excess, no longer needed, amino groups) are deaminated by the distal renal tubular epithelial cells.[5][9] Thus some of the "acid content" of the urine resides in the resulting ammonium ion (NH4+) content of the urine, though this has no effect on pH homeostasis of the extracellular fluids.[5][20]

Imbalance

Acid-base nomogram
An acid base nomogram for human plasma, showing the effects on the plasma pH when carbonic acid (partial pressure of carbondioxide) or bicarbonate occur in excess or are deficient in the plasma

Acid–base imbalance occurs when a significant insult causes the blood pH to shift out of the normal range (7.32 to 7.42[14]). An abnormally low pH in the ECF is called an acidaemia and an abnormally high pH is called an alkalaemia.

A second pair of terms is used in acid-base pathophysiology: "acidosis" and "alkalosis". They are often used as synonyms for "acidaemia" and "alkalaemia",[21] though this can cause confusion. "Acidaemia" refers unambiguously to the actual change in the pH of the ECF, whereas "acidosis", strictly speaking, refers to either a rise in the amount of carbonic acid in the ECF or to a decrease in the amount of HCO
3
in the ECF. Either change would on its own (i.e. if left "uncompensated" by an alkalosis) cause an acidaemia.[21] Similarly an alkalosis refers to a rise in the concentration of bicarbonate in the ECF, or to a fall on the partial pressure of carbon dioxide, either of which would on their own raise the pH of the ECF above the normal value.[21] The terms acidosis and alkalosis should always be qualified by an adjective to indicate the cause of the disturbance: "respiratory" (indicating a change in the partial pressure of carbon dioxide),[22] or "metabolic" (indicating a change in the bicarbonate concentration of the ECF).[5][23] There are therefore four different acid-base problems: metabolic acidosis, respiratory acidosis, metabolic alkalosis, and respiratory alkalosis.[5] One or a combination these conditions may occur simultaneously. For instance, a metabolic acidosis (as in uncontrolled diabetes mellitus) is almost always partially compensated by a respiratory alkalosis (hyperventilation), or a respiratory acidosis can be completely or partially corrected by a metabolic alkalosis.

Whether an acidosis causes an acidaemia or not depends on the magnitude of the accompanying alkalosis. If the one cancels the other out (i.e. the ratio of carbonic acid to bicarbonate is returned to 1:20) then there is neither an acidaemia or an alkalaemia.[5] If the accompanying alkalosis overwhelms the acidosis then an alkalaemia results; whereas if the acidosis is greater than the alkalosis then an acidaemia is the inevitable result. The same considerations determine whether an alkalosis results in an alkalaemia or not.

The normal pH in the fetus differs from that in the adult. In the fetus, the pH in the umbilical vein pH is normally 7.25 to 7.45 and that in the umbilical artery is normally 7.18 to 7.38.[24]

See also

References

  1. ^ a b Hamm, LL; Nakhoul, N; Hering-Smith, KS (7 December 2015). "Acid-Base Homeostasis". Clinical Journal of the American Society of Nephrology. 10 (12): 2232–42. doi:10.2215/CJN.07400715. PMC 4670772. PMID 26597304.
  2. ^ J., Tortora, Gerard (2012). Principles of anatomy & physiology. Derrickson, Bryan. (13th ed.). Hoboken, NJ: Wiley. pp. 42–43. ISBN 9780470646083. OCLC 698163931.
  3. ^ Macefield, Gary; Burke, David (1991). "Paraesthesiae and tetany induced by voluntary hyperventilation: increased excitability of cutaneous and motor axons". Brain. 114 (1): 527–540. doi:10.1093/brain/114.1.527.
  4. ^ Stryer, Lubert (1995). Biochemistry (Fourth ed.). New York: W.H. Freeman and Company. pp. 347, 348. ISBN 0 7167 2009 4.
  5. ^ a b c d e f g h i j k l m n o p q r s t u v Silverthorn, Dee Unglaub (2016). Human physiology. An integrated approach (Seventh, Global ed.). Harlow, England: Pearson. pp. 607–608, 666–673. ISBN 1-292-09493-1.
  6. ^ Adrogué, H. E.; Adrogué, H. J. (April 2001). "Acid-base physiology". Respiratory Care. 46 (4): 328–341. ISSN 0020-1324. PMID 11345941.
  7. ^ a b MedlinePlus Encyclopedia Metabolic acidosis
  8. ^ a b Tortora, Gerard J.; Anagnostakos, Nicholas P. (1987). Principles of anatomy and physiology (Fifth ed.). New York: Harper & Row, Publishers. pp. 581–582, 675–676. ISBN 0-06-350729-3.
  9. ^ a b c Stryer, Lubert (1995). Biochemistry (Fourth ed.). New York: W.H. Freeman and Company. pp. 39, 164, 630–631, 716–717. ISBN 0 7167 2009 4.
  10. ^ a b c d Tortora, Gerard J.; Anagnostakos, Nicholas P. (1987). Principles of anatomy and physiology (Fifth ed.). New York: Harper & Row, Publishers. pp. 494, 556–582. ISBN 0-06-350729-3.
  11. ^ a b c d Tortora, Gerard J.; Anagnostakos, Nicholas P. (1987). Principles of anatomy and physiology (Fifth ed.). New York: Harper & Row, Publishers. pp. 698–700. ISBN 0-06-350729-3.
  12. ^ a b c d e Bray, John J. (1999). Lecture notes on human physiology. Malden, Mass.: Blackwell Science. p. 556. ISBN 978-0-86542-775-4.
  13. ^ a b Garrett, Reginald H.; Grisham, Charles M (2010). Biochemistry. Cengage Learning. p. 43. ISBN 978-0-495-10935-8.
  14. ^ a b Diem, K.; Lentner, C. (1970). "Blood – Inorganic substances". in: Scientific Tables (Seventh ed.). Basle, Switzerland: CIBA-GEIGY Ltd. p. 527.
  15. ^ MedlinePlus Encyclopedia Blood gases
  16. ^ Caroline, Nancy (2013). Nancy Caroline's Emergency care in the streets (7th ed.). Buffer systems: Jones & Bartlett Learning. pp. 347–349. ISBN 978-1449645861.
  17. ^ Hamm, L. Lee; Nakhoul, Nazih; Hering-Smith, Kathleen S. (2015-12-07). "Acid-Base Homeostasis". Clinical Journal of the American Society of Nephrology. 10 (12): 2232–2242. doi:10.2215/CJN.07400715. ISSN 1555-905X. PMC 4670772. PMID 26597304.
  18. ^ a b c d J., Tortora, Gerard (2010). Principles of anatomy and physiology. Derrickson, Bryan. (12th ed.). Hoboken, NJ: John Wiley & Sons. p. 907. ISBN 9780470233474. OCLC 192027371.
  19. ^ Levitzky, Michael G. (2013). Pulmonary physiology (Eighth ed.). New York: McGraw-Hill Medical. p. Chapter 9. Control of Breathing. ISBN 978-0-07-179313-1.
  20. ^ Rose, Burton; Helmut Rennke (1994). Renal Pathophysiology. Baltimore: Williams & Wilkins. ISBN 0-683-07354-0.
  21. ^ a b c Andertson, Douglas M. (2003). Dorland's illustrated medical dictionary (30th ed.). Philadelphia PA: Saunders. pp. 17, 49. ISBN 0-7216-0146-4.
  22. ^ Brandis, Kerry. Acid-base physiology Respiratory acidosis: definition. http://www.anaesthesiamcq.com/AcidBaseBook/ab4_1.php
  23. ^ Brandis, Kerry. Acid-base physiology Metabolic acidosis: definition. http://www.anaesthesiamcq.com/AcidBaseBook/ab5_1.php
  24. ^ Yeomans, ER; Hauth, JC; Gilstrap, LC III; Strickland DM (1985). "Umbilical cord pH, PCO2, and bicarbonate following uncomplicated term vaginal deliveries (146 infants)". Am J Obstet Gynecol. 151: 798–800. doi:10.1016/0002-9378(85)90523-x. PMID 3919587.

External links

Acidosis

Acidosis is a process causing increased acidity in the blood and other body tissues (i.e., an increased hydrogen ion concentration). If not further qualified, it usually refers to acidity of the blood plasma.

The term acidemia describes the state of low blood pH, while acidosis is used to describe the processes leading to these states. Nevertheless, the terms are sometimes used interchangeably. The distinction may be relevant where a patient has factors causing both acidosis and alkalosis, wherein the relative severity of both determines whether the result is a high, low, or normal pH.

Acidemia is said to occur when arterial pH falls below 7.35 (except in the fetus – see below), while its counterpart (alkalemia) occurs at a pH over 7.45. Arterial blood gas analysis and other tests are required to separate the main causes.

The rate of cellular metabolic activity affects and, at the same time, is affected by the pH of the body fluids. In mammals, the normal pH of arterial blood lies between 7.35 and 7.50 depending on the species (e.g., healthy human-arterial blood pH varies between 7.35 and 7.45). Blood pH values compatible with life in mammals are limited to a pH range between 6.8 and 7.8. Changes in the pH of arterial blood (and therefore the extracellular fluid) outside this range result in irreversible cell damage.

Ailsa A. Welch

Professor Ailsa A. Welch is a professor of nutritional epidemiology at Norwich Medical School (part of the University of East Anglia) in the UK. Her research focuses on the impact of human nutrition on health, disease and aging. She is listed as a notable scientist in Thomson Reuters' Highly Cited Researchers 2014, ranking her among the top 1% most cited scientists.

Albert Baird Hastings

Albert Baird Hastings (November 20, 1895 – September 24, 1987) was an American biochemist and physiologist. He spent 28 years as the department chair and Hamilton Kuhn Professor of Biological Chemistry at Harvard University. After retiring from Harvard, Hastings moved to the Scripps Clinic and Research Foundation (now the Scripps Research Institute), where he became the director of the division of biochemistry and helped to establish the institution's emerging program in basic research. In 1966, he became one of the first faculty members at the University of California, San Diego's new medical school. His research focused on the biochemical underpinnings of physiology and included characterizing acid-base homeostasis in blood and pioneering the use of radioactive tracers for studying metabolism. Hastings received a number of honors and awards for his work, including election to the National Academy of Sciences in 1937 and the President's Medal for Merit in 1948 following his wartime service on the Committee for Medical Research. Hastings died of heart failure in 1987 at age 91.

Alkaline diet

Alkaline diet (also known as the alkaline ash diet, alkaline acid diet, acid ash diet, and acid alkaline diet) describes a group of loosely related diets based on the misconception that different types of food can have an effect on the pH balance of the body. It originated from the acid ash hypothesis, which primarily related to osteoporosis research. Proponents of the diet believe that certain foods can affect the acidity (pH) of the body and that the change in pH can therefore be used to treat or prevent disease. Due to the lack of credible evidence supporting the claimed mechanism of this diet, it is not recommended by dietitians or other health professionals, though some have noted that eating unprocessed foods as this diet recommends may have incidental health benefits unrelated to bodily pH.These diets have been promoted by alternative medicine practitioners, who propose that such diets treat or prevent cancer, heart disease, low energy levels, and other illnesses. Human blood is maintained between pH 7.35 and 7.45 by acid–base homeostasis mechanisms. Levels above 7.45 are referred to as alkalosis and levels below 7.35 as acidosis. Both are potentially serious. The idea that these diets can materially affect blood pH for the purpose of treating a range of diseases is not supported by scientific research and makes incorrect assumptions about how alkaline diets function that are contrary to human physiology.While diets avoiding meat, poultry, cheese, and grains can be used in order to make the urine more alkaline (higher pH), difficulties in effectively predicting the effects of these diets have led to medications, rather than diet modification, as the preferred method of changing urine pH. The "acid-ash" hypothesis was once considered a risk factor for osteoporosis, though the current weight of scientific evidence does not support this hypothesis.

Alkalosis

Alkalosis is the result of a process reducing hydrogen ion concentration of arterial blood plasma (alkalemia). In contrast to acidemia (serum pH 7.35 or lower), alkalemia occurs when the serum pH is higher than normal (7.45 or higher). Alkalosis is usually divided into the categories of respiratory alkalosis and metabolic alkalosis or a combined respiratory/metabolic alkalosis.

Base excess

In physiology, base excess and base deficit refer to an excess or deficit, respectively, in the amount of base present in the blood. The value is usually reported as a concentration in units of mEq/L, with positive numbers indicating an excess of base and negative a deficit. A typical reference range for base excess is −2 to +2 mEq/L.Comparison of the base excess with the reference range assists in determining whether an acid/base disturbance is caused by a respiratory, metabolic, or mixed metabolic/respiratory problem. While carbon dioxide defines the respiratory component of acid-base balance, base excess defines the metabolic component. Accordingly, measurement of base excess is defined, under a standardized pressure of carbon dioxide, by titrating back to a standardized blood pH of 7.40.

The predominant base contributing to base excess is bicarbonate. Thus, a deviation of serum bicarbonate from the reference range is ordinarily mirrored by a deviation in base excess. However, base excess is a more comprehensive measurement, encompassing all metabolic contributions.

Bicarbonate

In inorganic chemistry, bicarbonate (IUPAC-recommended nomenclature: hydrogencarbonate) is an intermediate form in the deprotonation of carbonic acid. It is a polyatomic anion with the chemical formula HCO−3.

Bicarbonate serves a crucial biochemical role in the physiological pH buffering system.The term "bicarbonate" was coined in 1814 by the English chemist William Hyde Wollaston. The prefix "bi" in "bicarbonate" comes from an outdated naming system and is based on the observation that there is twice as much carbonate (CO2−3) per sodium ion in sodium bicarbonate (NaHCO3) and other bicarbonates than in sodium carbonate (Na2CO3) and other carbonates. The name lives on as a trivial name.

Bone health

The human skeletal system is a complex organ in constant equilibrium with the rest of the body. In addition to support and structure of the body, bone is the major reservoir for many minerals and compounds essential for maintaining a healthy pH balance. The deterioration of the body with age renders the elderly particularly susceptible to and affected by poor bone health. Illnesses like osteoporosis, characterized by weakening of the bone’s structural matrix, increases the risk of hip-fractures and other life-changing secondary symptoms. In 2010, over 258,000 people aged 65 and older were admitted to the hospital for hip fractures. Incidence of hip fractures is expected to rise by 12% in America, with a projected 289,000 admissions in the year 2030. Other sources estimate up to 1.5 million Americans will have an osteoporotic-related fracture each year. The cost of treating these people is also enormous, in 1991 Medicare spent an estimated $2.9 billion for treatment and out-patient care of hip fractures, this number can only be expected to rise.

Buffering agent

A buffering agent is a weak acid or base used to maintain the acidity (pH) of a solution near a chosen value after the addition of another acid or base. That is, the function of a buffering agent is to prevent a rapid change in pH when acids or bases are added to the solution. Buffering agents have variable properties—some are more soluble than others; some are acidic while others are basic. As pH managers, they are important in many chemical applications, including agriculture, food processing, biochemistry, medicine and photography.

Carbonic acid

Not to be confused with carbolic acid, an antiquated name for phenol.Carbonic acid is a chemical compound with the chemical formula H2CO3 (equivalently OC(OH)2). It is also a name sometimes given to solutions of carbon dioxide in water (carbonated water), because such solutions contain small amounts of H2CO3. In physiology, carbonic acid is described as volatile acid or respiratory acid, because it is the only acid excreted as a gas by the lungs. It plays an important role in the bicarbonate buffer system to maintain acid–base homeostasis.

Carbonic acid, which is a weak acid, forms two kinds of salts: the carbonates and the bicarbonates. In geology, carbonic acid causes limestone to dissolve, producing calcium bicarbonate, which leads to many limestone features such as stalactites and stalagmites.

It was long believed that carbonic acid could not exist as a pure compound. However, in 1991 it was reported that NASA scientists had succeeded in making solid H2CO3 samples.

Collecting duct system

The collecting duct system of the kidney consists of a series of tubules and ducts that physically connect nephrons to a minor calyx or directly to the renal pelvis. The collecting duct system participates in electrolyte and fluid balance through reabsorption and excretion, processes regulated by the hormones aldosterone and vasopressin (antidiuretic hormone).

There are several components of the collecting duct system, including the connecting tubules, cortical collecting ducts, and medullary collecting ducts.

Equilibrium constant

For experimental methods and computational details see Determination of equilibrium constants.The equilibrium constant of a chemical reaction is the value of its reaction quotient at chemical equilibrium, a state approached by a dynamic chemical system after sufficient time has elapsed at which its composition has no measurable tendency towards further change. For a given set of reaction conditions, the equilibrium constant is independent of the initial analytical concentrations of the reactant and product species in the mixture. Thus, given the initial composition of a system, known equilibrium constant values can be used to determine the composition of the system at equilibrium. However, reaction parameters like temperature, solvent, and ionic strength may all influence the value of the equilibrium constant.

A knowledge of equilibrium constants is essential for the understanding of many chemical systems, as well as biochemical processes such as oxygen transport by hemoglobin in blood and acid-base homeostasis in the human body.

Stability constants, formation constants, binding constants, association constants and dissociation constants are all types of equilibrium constants.

Fetal scalp blood testing

Fetal scalp blood testing is a technique used in obstetrics during labor to confirm whether fetal oxygenation is sufficient.

The procedure can be performed by creating a shallow cut by a transvaginally inserted blood lancet, followed by applying a thin pipe to the site that samples blood by capillary action.

Two constituents that are commonly tested by this method are pH and lactate, both being indicators of acid base homeostasis. A low pH and high level of lactate indicate that there is acidosis, which in turn is associated with hypoxia.

pH and lactate appear to have the same sensitivity in indicating hypoxia during labour. Analysis of pH requires a relatively large amount of blood (30–50 μl), and sampling failure rates of 11–20% have been reported. Analysis of lactate only requires 5 μl of blood.

Kidney

The kidneys are two bean-shaped organs found in vertebrates. They are located on the left and right in the retroperitoneal space, and in adult humans are about 11 centimetres (4.3 in) in length. They receive blood from the paired renal arteries; blood exits into the paired renal veins. Each kidney is attached to a ureter, a tube that carries excreted urine to the bladder.

The nephron is the structural and functional unit of the kidney. Each human adult kidney contains around 1 million nephrons, while a mouse kidney contains only about 12,500 nephrons. The kidney participates in the control of the volume of various body fluid compartments, fluid osmolality, acid-base balance, various electrolyte concentrations, and removal of toxins. Filtration occurs in the glomerulus: one-fifth of the blood volume that enters the kidneys is filtered. Examples of substances reabsorbed are solute-free water, sodium, bicarbonate, glucose, and amino acids. Examples of substances secreted are hydrogen, ammonium, potassium and uric acid. The kidneys also carry out functions independent of the nephron. For example, they convert a precursor of vitamin D to its active form, calcitriol; and synthesize the hormones erythropoietin and renin.

Renal physiology is the study of kidney function. Nephrology is the medical specialty which addresses diseases of kidney function: these include chronic kidney disease, nephritic and nephrotic syndromes, acute kidney injury, and pyelonephritis. Urology addresses diseases of kidney (and urinary tract) anatomy: these include cancer, renal cysts, kidney stones and ureteral stones, and urinary tract obstruction.Procedures used in the management of kidney disease include chemical and microscopic examination of the urine (urinalysis), measurement of kidney function by calculating the estimated glomerular filtration rate (eGFR) using the serum creatinine; and kidney biopsy and CT scan to evaluate for abnormal anatomy. Dialysis and kidney transplantation are used to treat kidney failure; one (or both sequentially) of these are almost always used when renal function drops below 15%. Nephrectomy is frequently used to cure renal cell carcinoma.

PCO2

The pCO2, PCO2, or is the partial pressure of carbon dioxide (CO2), often used in reference to blood, but also used in oceanography to describe the partial pressure of CO2 in the Ocean, and in life support systems engineering and underwater diving to describe the partial pressure in a breathing gas. Usually the arterial blood is the relevant context; the symbol for in arterial blood is . Measurement of in the systemic circulation indicates the effectiveness of ventilation at the lungs' alveoli, given the diffusing capacity of the gas. It is a good indicator of respiratory function and the closely related factor of acid–base homeostasis, reflecting the amount of acid in the blood (without lactic acid).

PH

In chemistry, pH () is a scale used to specify how acidic or basic a water-based solution is. Acidic solutions have a lower pH, basic solutions have a higher pH. At room temperature, pure water is neither acidic nor basic and has a pH of 7.

The scale is logarithmic. It is approximately the negative of the base 10 logarithm of the molar concentration (measured in units of moles per liter) of hydrogen ions. More precisely it is the negative of the base 10 logarithm of the activity of the hydrogen ion. At 25 °C, solutions with a pH less than 7 are acidic and solutions with a pH greater than 7 are basic. The neutral value of the pH depends on the temperature, being lower than 7 if the temperature increases. Pure water is neutral (pH 7) at 25 °C. Contrary to popular belief, the pH value can be less than 0 or greater than 14 for very strong acids and bases respectively.Measurements of pH are important in agronomy, medicine, chemistry, water treatment, and many other applications.

The pH scale is traceable to a set of standard solutions whose pH is established by international agreement.

Primary pH standard values are determined using a concentration cell with transference, by measuring the potential difference between a hydrogen electrode and a standard electrode such as the silver chloride electrode.

The pH of aqueous solutions can be measured with a glass electrode and a pH meter, or an indicator.

There are three current theories used to describe acid–base reactions: Arrhenius, Bronsted-Lowry and Lewis when determining pH.

Respiratory alkalosis

Respiratory alkalosis is a medical condition in which increased respiration elevates the blood pH beyond the normal range (7.35–7.45) with a concurrent reduction in arterial levels of carbon dioxide. This condition is one of the four basic categories of disruption of acid–base homeostasis.

Van Slyke determination

The Van Slyke determination is a chemical test for the determination of amino acids containing a primary amine group. It is named after the biochemist Donald Dexter Van Slyke (1883-1971).One of Van Slyke's first professional achievements was the quantification of amino acids by the Van Slyke determination reaction. To quantify aliphatic amino acids, the sample is diluted in glycerol and then treated with a solution of sodium nitrite, water and acetic acid. The resulting diazotisation reaction produces nitrogen gas which can be observed qualitatively or measured quantitatively.Van Slyke Reaction: R-NH2 + HONO → ROH + N2 + H2O In addition, Van Slyke developed the so-called Van Slyke apparatus, which can be used to determine the concentration of respiratory gases in the blood, especially the concentration of sodium bicarbonate. This was of high importance to be able to recognize a beginning acidosis in diabetic patients as early as possible, in order to start alkali treatment. The Van Slyke apparatus became a standard equipment in clinical laboratories around the world and the results of Van Slyke's research are still used today to determine abnormalities in the acid-base homeostasis. Later on, Van Slyke further improved his apparatus, increasing its accuracy and sensitivity. Using the new method, he was able to further investigate the role of gas and electrolyte equilibria in the blood and how they change in response to respiration.

Water ionizer

A water ionizer (also known as an alkaline ionizer) is a home appliance which claims to raise the pH of drinking water by using electrolysis to separate the incoming water stream into acidic and alkaline components. The alkaline stream of the treated water is called alkaline water. Proponents claim that consumption of alkaline water results in a variety of health benefits, making it similar to the alternative health practice of alkaline diets. Such claims violate basic principles of chemistry and physiology. There is no medical evidence for any health benefits of alkaline water.The machines originally became popular in Japan and other far eastern countries before becoming available in the U.S. and Europe.

Diagnostic
Disease
Therapy
See also
Volume status
Electrolyte
Acid–base
Renal function
Hormones
Acid-base balance
Other
Blood composition
Other

This page is based on a Wikipedia article written by authors (here).
Text is available under the CC BY-SA 3.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.