water, pH and buffers

Life as we know it requires three primary ingredients: liquid water; a source of energy for metabolism; and the right chemical ingredients, primarily carbon, hydrogen, nitrogen, oxygen, phosphorus and sulfur. With this finding, Cassini has shown that Enceladus – a small, icy moon a billion miles farther from the sun than Earth – has nearly all of these ingredients for habitability. NASA: Ingredients for Life at Saturn’s Moon Enceladus

Water is a major component of cells and accounts for the high percentage of oxygen

humans are made uo if 55- 60 % of water 

• Water plays a central role in the chemistry of all life

•it is the solvent in which most biochemical reactions take place. •Proteins, polysaccharides, nucleic acids and membranes all assume their characteristic shapes in response to water •The chemical properties of water are related to the functions of biomolecules, entire cells, and organisms

•Important properties of water arise from its angled  ‘V’ shape [it is not linear] •Angle of 104.5o between two covalent bonds •Distribution of charge is asymmetrical •Polar O-H bonds due to uneven distribution of charge (oxygen (d-), hydrogen (d+)) =dipole

•Important consequence of the polarity of water molecule is that water molecules attract one another. •Ahydrogen bond is formed: •Between water molecules: a partially positive hydrogen atom attracts the partially negative oxygen atom of a second water molecule •Other molecules: Hydrogen bonds can form between electronegative atoms and a hydrogen attached to another electronegative atom

•A water molecule can, in some cases, form up to four hydrogen bonds

•Hexagonal lattice structure •Every water molecule is H -bonded to 4 others

Ice is less dense than water- it floats!

•Hydrophilic (water-loving) substances (polar and ionic (electrolytes)) readily dissolve in H2O •Polar water molecules align themselves around ions or other polar molecules •A molecule or ion surrounded by solvent molecules is solvated •When the solvent is water the molecules or ions are hydrated

(a) Electrostatic forces hold ions together in crystalline sodium

(b) Water molecules form solvation spheres around  Na+ and Cl-

•Solubility in water depends upon the ratio of polar to non-polar groups in a molecule •The larger the portion of polar groups (e.g. hydroxyl groups (-OH)) the more soluble the molecule is in water

•Hydrophobic (water-fearing) molecules are nonpolar •Hydrophobic effect - the exclusion of nonpolar substances by water (critical for protein folding and self-assembly of biological membranes) (Oil in water) •Amphipathic molecules have hydrophobic chains and ionic or polar ends.  Detergents (surfactants) are examples.

•Strongest bonds •Involve sharing of electrons •Carbon-carbon covalent bonds are common in biological systems

Weak non-covalent interactions are

important in:

• Stabilization of proteins and nucleic acids • Recognition of one biopolymer by another • Binding of reactants to enzymes

There are 4 major types of non-covalent forces: 

  (1) Hydrogen bonds

  (2) Hydrophobic interactions

  (3) Charge-charge interactions

  (4) Van der Waals forces

•Among the strongest of non-covalent interactions •H atom bonded to N, O, S can hydrogen bond to another electronegative atom (~0.2 nm distance) •Total distance between the two electronegative atoms is ~0.27 to 0.30 nm •In aqueous solution, water can H-bond to exposed functional groups on biological molecules

•Association of a relatively nonpolar molecule or group with other nonpolar molecules •The cumulative effects of many hydrophobic interactions can have a significant effect on the stability of a macromolecule

•Electrostatic interactions between two charged particles •Can be the strongest type of noncovalent forces •Can extend over greater distances than other forces •Charge repulsion occurs between similarly charged groups

•Salt bridges - attractions between oppositely-charged functional groups in proteins •Ion pairing - a salt bridge buried in the hydrophobic interior of a protein is stronger than one on the surface

•Weak  short range forces between atoms that are close together:
(a) Permanent dipoles of two uncharged   molecules
(b) Permanent dipole and an induced dipole   in a neighboring molecule •Although individually weak, many van der Waals interactions occur in biological macromolecules and  participate in stabilizing molecular structures

•Hydrogen bonds •Hydrophobic interactions •Charge-charge interactions •Van der Waals interactions

•Nucleophiles - electron-rich atoms or groups •Electrophiles - electron-deficient atoms or groups •Water is a relatively weak nucleophile •Due to its high cellular concentration, hydrolysis reactions in water are thermodynamically favored

In Acid-base reactions- hydrogen ions are added or removed from molecules.

The concentration of hydrogen ions in solution is expressed as the pH.

****pH is defined as the negative logarithm of the concentration of  H+

•Lower values are acidic fluids •Higher values are basic fluids

•Pure water consists of a low concentration of hydronium ions (H3O+) and an equal concentration of hydroxide ions (OH-) •Acids are proton donors (e.g. H3O+) and bases are proton acceptors (e.g. OH-)

•The slight ionization of water can be expressed in terms of an equilibrium constant:

  Keq =      [H+][OH-]

               [H2O]

     = Kw

Kw is known as the ion product of water.

Strong acids and bases dissociate completely in water

  HCl + H2O  Cl- + H3O+

•Cl- is the conjugate baseof HCl- can accept H •H3O+ is the conjugate acid of H2O – can donate a H

•Weak acids and bases do not dissociate   completely in H2O

•Prior to the Equivalence Point - The base reacts stoichiometrically with HA to yield A 

    pH = pKa + log[A-]

                                      [HA]   

•Equivalence Point - At this point, the moles of OH- added exactly equals the moles of the weak acid used to prepare the sample. 

   pH = pKa

•Past the equivalence point - The OH- added will stoichiometrically react with HA.  The pH can be computed from the OH- in excess of HA.

•Defines the pH of a solution in terms of:

  (1)  The pKa of the weak acid

  (2)  Concentrations of the weak acid (HA) and conjugate base (A-)

•Significant change in pH à disrupt molecular structure •Buffer capacity is the ability of a solution to resist changes in pH •Most effective buffering occurs where:  
  solution pH = buffer pKa •At this point: [weak acid] = [conjugate base] •Effective buffering range is usually at pH values equal to the pKa ± 1 pH unit

•Blood plasma of mammals has a constant pH, usually 7.4, which is regulated by a buffer system of:
 carbon dioxide/carbonic acid/bicarbonate •Buffer capacity depends upon equilibria between:
  (1) Gaseous CO2 (air spaces of the lungs)
  (2) Aqueous CO2 (dissolved in the blood)
  (3) Carbonic acid
  (4) Bicarbonate