Lehninger Principles of Biochemistry Lecture Connections PDF

Lecture Connections 2 | Water and Aqueous Solutions © 2009 W. H. Freeman and Company  CHAPTER 2 Water and Aqueous Solutions Learning goals: to understand – What kind of interactions occur between molecules – Why water is a good medium for life – Why nonp...

Lecture Connections 2 | Water and Aqueous Solutions © 2009 W. H. Freeman and Company  CHAPTER 2 Water and Aqueous Solutions Learning goals: to understand – What kind of interactions occur between molecules – Why water is a good medium for life – Why nonpolar moieties aggregate in water – How dissolved molecules alter properties of water – How weak acids and bases behave in water – How buffers work and why we need them – How water participates in biochemical reactions  Physics of Non-covalent Interactions Non-covalent interactions do not involve sharing a pair of electrons. Based on their physical origin, one can distinguish between Ionic (Coulombic) Interactions – Electrostatic interactions between permanently charged species, or between the ion and a permanent dipole Dipole Interactions – Electrostatic interactions between uncharged, but polar molecules Van der Waals Interactions – Weak interactions between all atoms, regardless of polarity – Attractive (dispersion) and repulsive (steric) component Hydrophobic Effect – Complex phenomenon associated with the ordering of water molecules around non-polar substances Examples of Noncovalent Interactions  Hydrogen Bonds Strong dipole-dipole or charge-dipole interaction that arises between an acid (proton donor) and a base (proton acceptor) Typically 4-6 kJ/mol for bonds with neutral atoms, and 6-10 kJ/mol for bonds with one charged atom Typically involves two electronegative atoms (frequently nitrogen and oxygen) Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic interaction. Ideally the three atoms involved are in a line Hydrogen Bonds: Examples  Importance of Hydrogen Bonds Source of unique properties of water Structure and function of proteins Structure and function of DNA Structure and function of polysaccharides Binding of a substrates to enzymes Binding of hormones to receptors Matching of mRNA and tRNA Biological Relevance of Hydrogen Bonds  Van der Waals Interactions Van der Waals interactions have two components: – Attractive force (London dispersion) Depends on the polarizability – Repulsive force (Steric repulsion) Depends on the size of atoms Attraction dominates at longer distances (typically 0.4-0.7 nm) Repulsion dominates at very short distances There is a minimum energy distance (van der Waals contact distance)  Origin of the London Dispersion Force Quantum mechanical origin Instantaneous polarization by fluctuating charge distributions Universal and always attractive Stronger in polarizable molecules Important only at a short range  Biochemical Significance of Van der Waals Interactions Weak individually –Easily broken, reversible Universal: –Occur between any two atoms that are near each other Importance – determines steric complementarity – stabilizes biological macromolecules (stacking in DNA) – facilitates binding of polarizable ligands  Water is the Medium for Life Life evolved in water (UV protection) Organisms typically contain 70-90% water Chemical reactions occur in aqueous milieu Water is a critical determinant of the structure and function of proteins, nucleic acids, and membranes  Structure of the Water Molecule Octet rule dictates that there are four electron pairs around an oxygen atom in water. These electrons are on four sp3 orbitals Two of these pairs covalently link two hydrogen atoms to a central oxygen atom. The two remaining pairs remain nonbonding (lone pairs) Water geometry is a distorted tetrahedron The electronegativity of the oxygen atom induces a net dipole moment Because of the dipole moment, water can serve as both a hydrogen bond donor and acceptor.  Hydrogen Bonding in Water Water can serve as both – an H donor and – an H acceptor Up to four H-bonds per water molecule gives water the – anomalously high boiling point – anomalously high melting point – unusually large surface tension Hydrogen bonding in water is cooperative. Hydrogen bonds between neighboring molecules are weak (20 kJ/mole) relative to the H–O covalent bonds (420 kJ/mol)  Water as a Solvent Water is a good solvent for charged and polar substances – amino acids and peptides – small alcohols – carbohydrates Water is a poor solvent for nonpolar substances – nonpolar gases – aromatic moieties – aliphatic chains  Water Dissolves Many Salts High dielectric constant reduces attraction between oppositely charged ions in salt crystal, almost no attraction at large (> 40 nm) distance Strong electrostatic interactions between the solvated ions and water molecules lowers the energy of the system Entropy increases as ordered crystal lattice is dissolved  Ice: Water in a Solid State Water has many different crystal forms; the hexagonal ice is the most common Hexagonal ice forms a regular lattice, and thus has a low entropy Hexagonal ice has lower density than liquid water; ice floats  The Hydrophobic Effect Refers to the association or folding of non- polar molecules in the aqueous solution Is one of the main factors behind: – Protein folding – Protein-protein association – Formation of lipid micelles – Binding of steroid hormones to their receptors Does not arise because of some attractive direct force between two non-polar molecules  Solubility of Polar and Non- polar Solutes Why are non-polar molecules poorly soluble in water? Low Solubility of Hydrophobic Solutes can be Explained by Entropy Bulk water has little order: - high entropy Water near a hydrophobic solute is highly ordered: - low entropy Low entropy is thermodynamically unfavorable, thus hydrophobic solutes have low solubility  Origin of the Hydrophobic Effect (1) Consider amphipathic lipids in water Lipid molecules disperse in the solution; nonpolar tail of each lipid molecule is surrounded by ordered water molecules Entropy of the system decreases System is now in an unfavorable state Origin of the Hydrophobic effect (2) Non-polar portions of the amphipathic molecule aggregate so that fewer water molecules are ordered. The released water molecules will be more random and the entropy increases. All non-polar groups are sequestered from water, and the released water molecules increase the entropy further. Only polar “head groups” are exposed and make energetically favorable H-bonds.  Hydrophobic Effect Favors Ligand Binding Binding sites in enzymes and receptors are often hydrophobic Such sites can bind hydrophobic substrates and ligands such as steroid hormones Many drugs are designed to take advantage of the hydrophobic effect  Colligative Properties Some properties of solution — boiling point, melting point, and osmolarity — do not depend strongly on the nature of the dissolved substance. These are called colligative properties Other properties — viscosity, surface tension, taste, and color, among other — depend strongly on the chemical nature of the solute. These are non-colligative properties. Cytoplasm of cells are highly concentrated solutions and have high osmotic pressure Effect of Extracellular Osmolarity  Ionization of Water H+ + OH- H 2O  O-H bonds are polar and can dissociate heterolytically Products are a proton (H+) and a hydroxide ion (OH-) Dissociation of water is a rapid reversible process Most water molecules remain un-ionized, thus pure water has very low electrical conductivity (resistance: 18 M cm) The equilibrium H2O H+ + OH- is strongly to the left Extent of dissociation depends on the temperature  Proton Hydration Protons do not exist free in solution. They are immediately hydrated to form hydronium (oxonium) ions A hydronium ion is a water molecule with a proton associated with one of the non-bonding electron pairs Hydronium ions are solvated by nearby water molecules The covalent and hydrogen bonds are interchangeable. This allows for an extremely fast mobility of protons in water via “proton hopping” Proton Hopping  Ionization of Water: Quantitative Treatment Concentrations of participating species in an equilibrium process are not independent but are related via the equilibrium constant [H + ] [OH- ] H2O  H + OH + - Keq = ———— [H2O] Keq can be determined experimentally, it is 1.8 10-16 M at 25 °C [H2O] can be determined from water density, it is 55.5 M Ionic product of water:  -  14 2 K w  K eq [H 2 O] [H ][OH ] 1 10 M In pure water [H+] = [OH-] = 10-7 M  What is pH? pH is defined as the negative pH = -log[H+] logarithm of the hydrogen ion concentration. Simplifies equations K w [H  ][OH - ] 1 10  14 M 2 The pH and pOH must always add to 14  log[ H  ]  log[OH - ] 14 pH can be negative ([H+] = 6 M) In neutral solution, [H+] = [OH-] pH  pOH 14 and the pH is 7 pH Scale: 1 unit = 10-fold  Dissociation of Weak Electrolytes: Principle O Keq O Weak electrolytes H3C + H2O + H3O+ H3C dissociate only partially OH O- in water K a  K eq [H 2 O] Extent of dissociation is determined by the acid [H  ][CH 3COO - ] dissociation constant Ka Ka  1.74 10  5 M [CH 3COOH] We can calculate the  [CH 3COOH] pH if the Ka is known. [ H ] K a  [CH 3COO  ] But some algebra is needed! Dissociation of Weak Electrolytes: Example What is the final pH of a solution when 0.1 moles of acetic acid is adjusted to 1 L of water? O O Ka H3C H 3C + H+ We assume that the OH O- 0.1 – x x x only source of H+ is the weak acid [ x ][ x ] Ka  1.74 10  5 M [0.1 - x] To find the [H+], a x 2 1.74 10  6  1.74 10  5 x quadratic equation x 2  1.74 10  5 x  1.74 10  6 0 must be solved. x = 0.001310, pH = 2.883 Dissociation of Weak Electrolytes: Simplification O The equation can be Ka O H3C H 3C + H+ simplified if the amount of OH O- dissociated species is much 0.1 – x x x less than the amount of 0.1 x x undissociated acid [ x ][ x ] Approximation works for Ka  1.74 10  5 M [0.1] sufficiently weak acids and bases x 2 1.74 10  6 Check that x