Acid base balance in human body and buffer system

Acid base balance in human body, buffer system, chemical bonds

1. Acid Base Balance in Human Body Dr. Vinod Kumar Patil Assistant Professor Department of Food and Nutrition Faculty of Science Khaja Bandanawaz University Kalaburagi, Karnataka

2. Maintenance of the internal environment is one of the vital functions (circulation or respiration). Maintaining a stable anion and cation concentrations in blood plasma is denoted as ISOIONIA . Maintaining of constant proton (H) concentration. Maintenance of stable pH, also called ISOHYDRIA , is one of the basic components of the internal environment. pH is used for express concentration of the protons: pH = – log c(H) Plasma and extracellular space concentrations of the protons are held in very narrow physiologic range of the pH is 7.36-7.44. Value of pH higher than 7.44 in arteries is denoted as ALKALEMIA , pH lower than 7.36 is ACIDEMIA .

3. Extensive deviations of pH value can cause serious consequences . For example change of protein structure (i.e. enzymes), membranes permeability, and electrolyte distribution. Value of pH in arterial blood higher than 7.8 , resp. lower than 6.8 are incompatible with life as they apply for arterial blood. This corresponds to fact that there is 2.5 fold difference between intracellular and arterial H concentration. This concentration gradient drives the movement of H + from cells to blood. Therefore it is not surprising that venous pH and pH of interstitial fluid is lower (i.e. more acidic) than arterial pH. Approximate value is 7.35 . Source of acids in the body is chiefly metabolism , source of bases is predominantly nutrient.

4. Three types of reactions can be distinguished from point of view of the acid-base balance. (1) Proton-productive , (2) Proton-consumptive , (3) Proton-neutral . 1. PROTON-PRODUCTIVE REACTIONS Anaerobic glycolysis in muscles and erythrocytes Glucose → 2 CH 3 CHOHCOO - + 2 H + b) Ketogenesis – production of ketone bodies Fatty acids → ketone bodies + n H + c) Lipolysis TAG → 3 FA + glycerol + 3 H + d) Ureagenesis CO 2 + 2 NH 4 + → Urea + H 2 O + 2 H + 2 ) PROTON-CONSUMPTIVE REACTIONS Gluconeogenesis: 2 Lactate + 2 H + → Glucose Neutral and dicarboxylic amino acids oxidation

5. 3) PROTON-NEUTRAL REACTIONS a) Complete glucose oxidation b) Lipogenesis from glucose Human (healthy or not) every day produces great quantities of acids - source of protons. Organism is acidified by these processes: 1) Complete oxidation: Carbon skeleton → CO 2 + H 2 O → HCO 3 - + H + 2) Incomplete oxidation: Carbohydrates → Glucose → Pyruvate, Lactate + H + Triacylglycerol → Fatty acids, Ketone bodies + H + Phospholipids → Phosphate + H + Proteins → Amino acids → Sulphate, Urea + H +

6. Acids can be divided into two groups: V olatile acids (respiratory acids). Non-volatile acids (metabolic acids). The most important volatile acid is carbonic acid (H 2 CO 3 ), produced by reaction of carbon dioxide (CO 2 is acid-forming oxide) with water. 15,000 – 20,000 mmol CO 2 (therefore same amount of carbonic acid) is produced every day. Respiratory system is very efficiently eliminates it and justifies the term volatile acid.

7. Two groups are distinguished among non-volatile acid: (1) O rganic , and (2) I norganic . 1 mmol/kg of body weight is produced every day . Non-volatile acid could be either (1) metabolised , or (2) excreted (using mainly kidneys). Organic non-volatile acids are for example: (1) lactic acid , (2) fatty acids , (3) ketone bodies ( acetoacetic acid, β- hydroxybutyric acid ). C ontinually produced by metabolism (incomplete oxidation of TAG, carbohydrates, proteins). As organic non-volatile acids are products of metabolism in normal conditions they are oxidized completely to CO 2 and H 2 O. Therefore they have no influence on proton overall balance.

8. Inorganic non-volatile acids are: (1) H 2 SO 4 (sulphuric acid is produced by oxidation of sulfhydryl groups – e.g. in amino acids that contain sulphur, i.e. cysteine, methionine). (2) H 3 PO 4 (phosphoric acid is produced by hydrolysis of phosphoproteins, phospholipids, nucleic acids). Inorganic non-volatile acids are predominantly excreted in urine. ATP production is coupled with H + production . Human body is evolutionary capable to handle acid load .

9. Systems responsible for maintenance of the acid-base balance: 1) Chemical buffering systems Chemical buffering systems deal with pH deviations in common metabolism, chemical buffers act immediately (acute regulation) only in the short-term . 2) Respiratory system Respiratory system regulates carbon dioxide and is able to change pCO 2 by its elimination or retention. Respiratory centre is in brainstem and reacts in 1-3 minutes . 3) Kidneys Their role in acid-base balance is very complex and react in hours-days .

10. 4) Liver Liver is pivotal organ of the energetic metabolism it also have important influence on the acid-base balance. Liver is the most important tissue where ammonium is detoxified in both (1) urea cycle, and (2) glutamine synthesis. Which one of these fates of ammonium is favoured closely depends on status of the acid-base balance: a) NH 4 + → urea + 2H + → acidification of the body CO 2 + 2 NH 4 + → CO(NH 2 ) 2 + 2H + + H 2 O H + + HCO 3 + → H 2 O + CO 2 ( consumption of bicarbonate ) b) NH 4 + → glutamine synthesis → H + is not produced, glutamine is taken up by the kidneys. In the kidney is H + excreted as NH 4 + . 5) Myocardium Myocardium influences acid-base balance through lactate and ketone bodies oxidation.

11. Buffering systems Buffers are substances capable of releasing and binding H + . Short-term and acute changes in acid-base balance can be balanced by buffers . Each buffer keeps its particular pH . This pH could be calculated by means of the Henderson-Hasselbalch equation: pH = pK + log [conjugated base]/[acid] Henderson-Hasselbalch equation for bicarbonate buffer (HCO 3 - /CO 2 ): pH = pK H 2 CO 3 + log ([HCO 3 - ] / [H 2 CO 3 ]) pH = pK H 2 CO 3 + log ([HCO 3 - ] / α x pCO 2 ) pH = pK ± 1 is range where buffers work optimally. The ratio in bicarbonate buffer is 20:1 (HCO 3 - : CO 2 )

12. There are several buffer systems in the body. The most important include: Bicarbonate buffer (H 2 CO 3 - /CO 2 ) H aemoglobin buffer (in erythrocytes) P hosphate buffer Proteins A mmonium buffer Their importance differs as it depends on localization.

13. Localization Buffer Commentary Interstitial fluid (ISF) Bicarbonate Buffers metabolic acids Phosphate Low concentration – limited significance Proteins Low concentration – limited significance Blood Bicarbonate Buffers metabolic acids Haemoglobin Buffers CO (carbonic acid production) Plasma proteins Minor Phosphate Low concentration – limited significance Main buffer systems according to body compartments

14. Intercellular fluid (ICF) Proteins Significant buffer Phosphate Significant buffer Urine Phosphate Responsible for majority of the titratable urine acidity Ammonium Significant: elimination of ammonium nitrogen and protons; cation

15. Buffer Plasma Erythrocytes Together HCO 3 - / CO 2 35 % 18 % 53 % Hb / Hb-H + – 35 % 35 % Plasma proteins 7 % – 7 % Inorganic phosphate 1 % 1 % 2 % Organic phosphate - 3% 3% Total 43% 57 % 100% Blood buffers and their buffer capacity

16. The total number of protons in the nucleus of an atom gives us the atomic number of that atom , represented by letter ‘Z’ . All the atoms of a particular element have the same number of protons, and hence the same atomic number. The number of protons and neutrons combine to give us the mass number of an atom, represented by letter ‘A’. As both protons and neutrons are present in the nucleus of an atom, they are together called nucleons .

17. Chemical bonds  form when electrons can be simultaneously close to two or more nuclei, but beyond this, there is no simple, easily understood  theory that would not only explain why  atoms  bind together to form molecules, but would also predict the three-dimensional structures of the resulting compounds as well as the energies and other properties of the bonds themselves. Ionic bond , type of linkage formed from the electrostatic attraction between oppositely charged  ions  in a  chemical compound . Such a bond forms when the valence (outermost) electrons of one atom are transferred permanently to another atom. The atom that loses the electrons becomes a positively charged ion (cation), while the one that gains them becomes a negatively charged ion (anion).

18. Covalent Bonding Formation of an ionic bond by complete transfer of an electron from one atom to another is possible only for a fairly restricted set of elements. Covalent bonding, in which neither atom loses complete control over its valence electrons , is much more common. In a  covalent bond  the electrons occupy a region of space  between  the two nuclei and are said to be  shared  by them.

19. Hydrogen bond  (or  H-bond ) is primarily an  electrostatic  force of attraction between a  hydrogen  (H) atom which is  covalently bonded  to a more  electronegative  "donor" atom or group ( Dn ), and another electronegative atom bearing a  lone pair  of electrons—the hydrogen bond acceptor (Ac). Hydrogen bonds can be  inter molecular  (occurring between separate molecules) or  intra molecular  (occurring among parts of the same molecule).

20. Vander Waals force  is a distance-dependent interaction between  atoms  or  molecules . Unlike  ionic  or  covalent bonds , these attractions do not result from a  chemical electronic bond ; they are comparatively weak and therefore more susceptible to disturbance. The van der Waals force quickly vanishes at longer distances between interacting molecules. Van der Waals forces include attraction and repulsions between  atoms ,  molecules , as well as other  intermolecular forces .