Basic Cell Chemicals PDF

Basic cell Chemicals - B.L.T. Balasuriya Important molecules in a cell  Water  The principal fluid medium of the cell is water, which is present in most cells, except for fat cells, in a concentration of 70 % to 85 %.  Many cellular chemicals are dissolved in the wat...

Basic cell Chemicals - B.L.T. Balasuriya Important molecules in a cell  Water  The principal fluid medium of the cell is water, which is present in most cells, except for fat cells, in a concentration of 70 % to 85 %.  Many cellular chemicals are dissolved in the water. Others are suspended in the water as solid particulates.  Chemical reactions take place among the dissolved chemicals or at the surfaces of the suspended particles or membranes.  Ions  The most important ions in the cell are K+, Mg2+,PO43-,SO42-,HCO3- and smaller quantities of Na+,Cl-,Ca2+  The ions provide inorganic chemicals for cellular reactions. Also, they are necessary for operation of some of the cellular control mechanisms.  Ions acting at the cell membrane are required for transmission of electrochemical impulses in nerve and muscle fibers.  Proteins  After water, the most abundant substances in most cells are proteins, which normally constitute 10% to 20 % of the cell mass.  These can be divided into two types: - structural proteins e.g. - intracellular filaments Extracellular proteins (collagen, elastin fibers) - functional proteins e.g. - Enzymes, hormones, antibodies  Lipids  Especially important lipids are phospholipids and cholesterol (2 % of the total cell mass).  Phospholipids and cholesterol are mainly insoluble in water – form the cell membrane and intracellular membrane barriers that separate the different cell compartments.  Some cells contain large quantities of triglycerides, also called neutral fat.  In the fat cells, triglycerides often account for as much as 95 % of the cell mass.  The fat stored in these cells represents the body’s main “storehouse” - can later be dissoluted and provide energy wherever in the body it is needed.  Carbohydrates  Most human cells do not maintain large stores of carbohydrates; (about 1 % of their total mass ) but increases to as much as 3 % in muscle cells, occasionally 6% in liver cells.  Dissolved glucose is always present in the surrounding extracellular fluid.  A small amount of carbohydrate is always stored in the form of glycogen - an insoluble polymer of glucose. Types of Bonding  The different types of chemical bonding are determined by how the valence electrons are shared among the bonded atoms. - covalent bonding - ionic bonding Covalent bonds  The valence electrons are shared as pairs between the bonded atoms.  Pure covalent bonding only occurs when two nonmetal atoms of the same kind bind to each other. Polar covalent bonding  When two different nonmetal atoms are bonded or a nonmetal and a metal are bonded, then the bond is a mixture of covalent and ionic bonding called polar covalent bonding.  Polar molecules will have an overall dipole which can be represented with a dipole arrow (pointing to the more electronegative end of the molecule).  The quantitative measure of a molecule’s polarity is called its dipole moment. Ionic bonding  the valence electrons are completely transferred from one atom to the other atom.  Ionic bonds occur between metals and nonmetals when there is a large difference in electronegativity. The Hydrogen bond  Is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule.  Usually the electronegative atom is oxygen, nitrogen, or fluorine, which has a partial negative charge. Acids and bases  The Bronsted-Lowry Theory - An acid is a proton (hydrogen ion) donor. -A base is a proton (hydrogen ion) acceptor.  The Arrhenius Theory - Acids are substances which produce hydrogen ions in solution. - Bases are substances which produce hydroxide ions in solution. Neutralisation happens because hydrogen ions and hydroxide ions react to produce water. pH  It commonly expressed as negative log of the molar concentrations of hydrogen ions H + (really hydronium ions H30+) in solution.  So a solution of HCl with a pH of 2.0 has a concentration of hydronium ions of 1x 10 -2 (1/100). pH @ 25 0C  Pure water has [H+]=10-7 and thus pH=7.  Acids have a high [H+] and thus a low pH.  Bases have a low [H+] and thus a high pH. Ways to measure pH  pH meter  Electrode measures H+ concentration  Must standardize (calibrate) before using.  Indicator dyes and test strips  Less precise  Each indicator is only good for a small pH range (1-2 pH units)  But may be good for field usage, or measuring small volumes, or dealing with noxious samples. Why is pH important in biology?  pH affects solubility of many substances.  pH affects structure and function of most proteins - including enzymes.  Many cells and organisms (esp. plants and aquatic animals) can only survive in a specific pH environment.  In the plasma of healthy individuals, pH is slightly alkaline, maintained in the narrow range of 7.35 to 7.45.  Conversely, gastric fluid pH can be quite acidic (on the order of 2.0) and pancreatic secretions can be quite alkaline (on the order of 8.0).  Enzymatic activity and protein structure are frequently sensitive to pH; in any given body or cellular compartment, pH is maintained to allow for maximal enzyme/protein efficiency. Acidosis and alkalosis Buffer Solution Buffer Solution An acid/base equilibrium system that is capable of maintaining a relatively constant pH even if a small amount of strong acid or base is added. a) Components of a buffer solution: a mixture of a weak acid and its conjugate base e.g. acetic acid & sodium acetate (HA & A–) HA(aq) H+(aq) + A–(aq) or, ammonium chloride & ammonia (NH4+ & NH3) NH4+(aq) H+(aq) + NH3(aq)  pH of a Buffer Solution (b) pH of buffer solution: (pH ≈ pKa of HA) HA(aq) H+ + A–(aq) Initial “M” 0 “A” Change –x +x +x Equil M–x x A+x [H+][A–] x(A + x) Ka = = [HA] (M – x) Since Ka is usually small (< ~ 10–4): M – x ≈ M and A+x≈A thus, Ka ≈ x(A)/M or simply, Ka[HA] (general buffer system) [H+] ≈ [A–] and [H+] ≈ Kamoles HA moles A–  Henderson-Hasselbach Equation Ka[HA] (general buffer system) [H ] ≈ + [A–] Again assuming that x is small, you can take the log of both sides and rearrange to get the Henderson-Hasselbach equation; [base ] pH = pKa + log [acid] Buffering Action (LeChatelier’s Principle)  The major equilibrium in a buffer system is: HA(aq) H+(aq) + A–(aq) 1. What if some acid [H+] is added? Since H+ is a product of above equilibrium, the reaction will shift in reverse, so that [HA] will increase and [A–] will decrease But, as long as some A– remains, the system is still a mixture of HA and A–. Therefore, the general equation, [H+] ≈ Ka[HA]/[A–] still applies. Thus, [H+] is still ≈ Ka.  Buffering Action (cont’) 2. What if some base (OH-) is added instead? The added base (OH-) will neutralize some of the acid HA and produce more of the conjugate base A–. HA + OH– --> H2O + A– [HA] will decrease and [A-] will increase But, as long as some HA remains, the system is still a mixture of HA and A–, and ∴ still a buffer solution. buffer capacity: is the amount of acid or base you can add without causing a large change in pH. buffer range: the pH values for which a buffer system is the most effective. (usually ±1 pH unit of either side of pKa)  Important buffers in the human body  There are three important buffers in the human body. - The protein buffer system  Proteins are enormous molecules constructed from smaller compounds called amino acids.  One of these amino acids called histidine , is a weak base and is readily converted into its conjugate acid.  The conjugate acid of histidine has a pKa of 6.0 , but histidine is incorporated in to a protein, its pka can be as a high as 7.0.  Thus protein that contain histidine can be effective buffers around a pH of 7.  The range of proteins can act as buffers.  Protein buffers occur both inside and outside cells. Protein buffers are particularly important inside cells, maintaining intracellular fluid at a constant pH.  - The phosphate buffer system  The pka of phosphoric acid is 2.1, which is too law to serve as an effective buffer around a pH of 7.  However removing one proton from phosphoric acid produces H2PO4-, which has a pka of 7.2.  There for a buffer that contains H 2PO4- and HPO42- has a close pH to 7.2 and the pH can be adjusted as needed by changing the relative concentrations of these two ions.  This system works with the protein buffer to maintain the pH of intracellular fluid at an appropriate level.  - The carbonic acid buffer system  Is the most important extracellular buffer.  H2CO3 and bicarbonate ions are the primary buffering agents in blood plasma, which must be maintained at a pH of 7.4.  However the pka of carbonic acid at body temperature (370C) is only 6.1, significantly bellow the plasma pH.  Raising the buffer pH to 7.4 requires that the concentrations of HCO3- be 20 times larger than that of H2CO3.  Since very high concentrations of HCO 3- would disrupt the osmotic balance between the plasma and the intracellular fluids.  The only expedient is to maintain the concentration of H2CO3 at a very law level.  The normal concentrations of carbonic acid and bicarbonate in plasma are 0.0012M and 0.024 M respectively. Tutorial -3 1. One of the important buffer system in living cells is the phosphate buffer. I. What two substances make up this buffer ? II. Which of these substances neutralizes acid ? III. Which of these substances neutralizes bases? 2. Explain why the pH of your blood plasma goes up when you breath too fast? 3. Explain why proteins that contain histidine can help maintain the pH of body fluids around 7