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BCH2011: Textbook summary - Lecture 1 + 2

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Monash University

LECTURE 1 + 2 Weak Interactions in Aqueous Systems: Hydrogen bonds between water molecules provide the cohesive forces that make water a liquid in room temperature and ice with a highly ordered arrangement of molecules at cold temperatures. Polar biomolecules dissolve readily in water because they can replace water- water interactions with more energetically favourable water-solute interactions. Non-polar biomolecules are poorly soluble in water because they interfere with water-water interactions but are unable to form water-solute interactions. In aqueous solutions, nonpolar molecules tend to cluster together. Hydrogen Bonding Gives Water its Unusual Properties: Each hydrogen atom of a water molecule shares an electron pair with the central oxygen atom. The oxygen nucleus attracts electrons more strongly than does the hydrogen nucleus (a proton); that is, oxygen is more electronegative. This means that the shared electrons are more often in the vicinity of the oxygen atom than of the hydrogen. The result of this unequal electron sharing is two electric dipoles in the water molecule, one along each of the H-O bonds; each hydrogen atom bears a partial positive charge (S+), and the oxygen atom bears a partial negative charge equal in magnitude to the sum of the two partial positives (2S-). As a result, there is an electrostatic attraction between the oxygen atom of one water molecule and the hydrogen of another, called a hydrogen bond. Hydrogen bonds are fairly weak. Those in liquid water have a bond dissociation energy. At room temperature, the thermal energy of an aqueous solution (the kinetic energy of motion of the individual atoms and molecules) is of the same order of magnitude as that required to break hydrogen bonds. When water is heated, the increase in temperature reflects the faster motion of individuals water molecules. At any given time, most of the molecules in liquid water are hydrogen bonded; when one hydrogen bond breaks, another hydrogen bond forms, with the same partner or a new one. The apt phase ‘flickering clusters’ has been applied to the short-lived groups of water molecules interlinked by hydrogen bonds in liquid water. Extended networks of hydrogen bonded water molecules also form bridges between solutes (proteins, nucleic acids, etc.) that allow the larger molecules to interact with each other. The nearly tetrahedral arrangement of the orbitals about the oxygen atom allows each water molecule to form hydrogen bonds with as many as four neighboring water molecules. In liquid water at room temperature and atmospheric pressure, water molecules are disorganized and in continuous motion, so that each molecule forms hydrogen bonds with an average of only 3.4 other molecules. In ice, each water molecule is fixed in space and forms hydrogen bonds with a full complement of four other water molecules to yield a regular lattice structure. Hydrogen bonds account for the relatively high melting point of water, because much thermal energy is required to break a sufficient proportion of hydrogen bonds to destabilize the crystal lattice of ice. When ice melts or water evaporates, heat is taken up by the system. During melting or evaporation, the entropy of the aqueous system increases as highly ordered arrays of water molecules in ice relax into the less orderly hydrogen-bonded arrays in liquid water or into the wholly disordered gaseous state. At room temperature, both the melting of ice and the evaporation of water occur spontaneously; the tendency of the water molecules to associate through hydrogen bonds is outweighed by the energetic push toward randomness. The free-energy change (∆G) must have a negative value for a process to occur spontaneously. ∆G = ∆H - T∆S, where ∆G represents the driving force, and ∆S the change in randomness. Because ∆H is positive for melting and evaporation, it is clearly that increase in entropy (∆S) that makes ∆G negative and drives these changes. Water forms Hydrogen Bonds with Polar Solutes: Hydrogen bonds readily form between an electronegative atom (the hydrogen acceptor, usually oxygen or nitrogen) and a hydrogen atom covalently bonded to another electronegative atom (the hydrogen atom) in the same or another molecule. Hydrogen atoms covalently bonded to carbon atoms do not participate in hydrogen bonding, because carbon is only slightly more electronegative than hydrogen and thus the C-H bond is very weakly polar. Uncharged but polar biomolecules such as sugars dissolve readily in water because of the stabilizing effect of hydrogen bonds between the hydroxyl groups or carbonyl oxygen of the sugar and the polar water molecules. Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic interaction, which occurs when the hydrogen atom and the two atoms that share it are in a straight line – that is, when the acceptor atom is in line with the covalent bond between the donor atom and H. This arrangement puts the positive charge of the hydrogen ion directly between the two partial negative charges. Hydrogen bonds are thus highly directional and capable of holding two hydrogen-bonded molecules or groups in a specific geometric arrangement. Water Interacts Electrostatically with Charged Solutes: Water is a polar solvent. It readily dissolves most biomolecules, which are generally charged or polar compounds; compounds that dissolve easily in water are hydrophilic. Nonpolar solvents are poor solvents for polar biomolecules but easily dissolve those that are hydrophobic – nonpolar molecules such as lipids and waxes. Water dissolves salts such as NaCl by hydrating and stabilizing the Na+ and Cl- ions, weakening the electrostatic interactions between them and thus counteracting their tendency to associate in a crystalline lattice. Water also readily dissolves charged biomolecules, including compounds with functional groups such as (COO-), (NH3+), and phosphate esters or anhydrides. Water replaces the solute-solute hydrogen bonds linking these biomolecules to each other with solute-water hydrogen bonds, thus screening the electrostatic interactions between solute molecules. Entropy Increases as Crystalline Substances Dissolve: As a salt such as NaCl dissolves, the Na+ and Cl- ions leaving the crystal lattice acquire far greater freedom of motion. The resulting increase in entropy of the system is largely responsible for the ease of dissolving salts such as NaCl in water. Nonpolar Gases are Poorly Soluble in Water: The molecules of the biologically important gases CO2, O2, and N2 are nonpolar. In O2 and N2, electrons are shared equally by both atoms. In CO2, each C=O bond is pola
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