[BCH 261] - Midterm Exam Guide - Ultimate 36 pages long Study Guide!

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BCH 261
MIDTERM EXAM
STUDY GUIDE
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Chapter 2  Water
The attractive forces between water molecules & the slight tendency of water to ionize are of crucial
importance to structure and function of biomolecules.
2.1 Weak Interactions in Aqueous Systems
Hydrogen bonds between water molecules provide cohesive forces that make water a liquid at
room temperature (3.4 hydrogen bonds per water molecule)
Favored is the extreme ordering of molecules typical crystalline water (ice).
Polar biomolecules dissolve readily in water because they can replace water-water interactions
with more energetically favorable water-solute interactions.
Non-polar biomolecules interfere with water-water interactions but are unable to form water-
solute interactions therefore nonpolar molecules are poorly soluble in water.
In aqueous solutions nonpolar molecules tend to cluster together.
Hydrogen bonds, ionic bonds, hydrophobic & van der Waals interactions are individually weak,
but collectively have significant influence on the three-dimensional structures of proteins,
nucleic acids, polysaccharides & membrane lipids.
Hydrogen Bonding Gives Water Its Unusual Properties
Water has higher melting point, boiling point, and heat of vaporization than most other common
solvents (Table 2–1).
Unusual properties are a consequence of attractions between adjacent water molecules
(hydrogen bonding)  giving liquid water greater internal attraction & ability to stay together.
Electron structure of the H2O molecule reveals the cause of these intermolecular attractions.
Each hydrogen atom of a water molecule shares electron pair with central oxygen atom.
Geometry of the molecule dictated by the shapes of the outer electron orbitals of the oxygen
atom, which are similar to the sp3 bonding orbitals of carbon (see Fig. 1–14).
These orbitals describe a rough tetrahedron, with a hydrogen atom at each of two corners and
unshared electron pairs at the other two corners (Fig. 2–1a). The H-O- H bond angle is 104.5o ,
slightly less than the 109.5o of a perfect tetrahedron (because of crowding by the nonbonding
orbitals of the oxygen atom)
The oxygen nucleus attracts electrons more strongly than does the hydrogen nucleus (a proton);
oxygen is more electronegative.
The sharing of electrons between H and O is therefore unequal the electrons more often in the
vicinity of the oxygen atom than of the hydrogen.
The result of unequal electron sharing = electric dipoles in the water molecule, one along each
of the H- O bond (hydrogen= partial positive charge, Oxygen atom= partial negative charge)
As a result, there is an electrostatic attraction between the oxygen atom of one water molecule
and the hydrogen of another (Fig. 2–1c), called a hydrogen bond. (Book represents H bond with
3 parallel blue lines as in Figure 2–1c.)
Hydrogen bonds are relatively weak. Those in liquid water have a bond dissociation energy (the
energy required to break a bond) of about 23 kJ/mol, O-H has bond dissociation of 470 kJ/mol
in H2O and 348 kJ/mol for a covalent C-C bond.
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Hydrogen bond is about 10% covalent, due to overlaps in bonding orbitals, and about 90%
electrostatic. At room temperature thermal energy of aqueous solution (the kinetic energy of
motion of the individual atoms and molecules) is of same order of magnitude as that required to
break hydrogen bonds.
When water is heated, increase in temperature reflects the faster motion of individual water
molecules. At any given time, most of the molecules in liquid water are engaged in hydrogen
bonding (but the lifetime of each hydrogen bond is just 1 to 20 picoseconds)
Upon breakage of one hydrogen bond, another hydrogen bond forms, with the same partner or a
new one, within 0.1 ps.
The sum of all the hydrogen bonds between H2O molecules confers great internal cohesion on
liquid water Extended networks of hydrogen-bonded water molecules can form bridges
between solutes (proteins and nucleic acids, for example) allowing the larger molecules to
interact with each other over distances of several nanometers without physically touching.
The nearly tetrahedral arrangement of orbitals about oxygen atom (Fig. 2–1a) 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, however, water molecules are
disorganized and in continuous motion each molecule forms hydrogen bonds with an average
of only 3.4 other molecules.
In ice on the other hand, each water molecule is fixed & forms hydrogen bonds with a full
complement of four other water molecules  yields a regular lattice structure (Fig. 2–2).
Breaking a sufficient proportion of hydrogen bonds to destabilize the crystal lattice of ice
requires much thermal energy accounts for the relatively high melting point of water (Table
2–1). When ice melts or water evaporates, heat is taken up by the system:
H2O (solid)  H2O (liquid) releases 5.9 kJ/ mol heat NRG (enthalpy)
H2O (liquid)  H2O (gas) releases 44.0 kJ/ mol heat NRG (enthalpy)
During melting or evaporation entropy of the aqueous system increases as more highly
ordered arrays of water molecules relax into the less orderly hydrogen bonded arrays in liquid
water or 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. Recall free-energy change (delta G)
must have a negative value for a process to occur spontaneously: delta G = delta H - T delta S,
where delta G represents the driving force, delta H the enthalpy change from making and
breaking bonds, and delta S the change in randomness.
Because delta H is positive for melting & evaporation, it is clearly the increase in entropy (delta
S) that makes delta G negative and drives these transformations.
Water Forms Hydrogen Bonds with Polar Solutes
Hydrogen bonds are not unique to water. Readily form between an electronegative atom (the
hydrogen acceptor, usually oxygen or nitrogen with a lone pair of electrons) and a hydrogen
atom covalently bonded to another electronegative atom (the hydrogen donor) in the same or
another molecule (Fig. 2–3).
Hydrogen atoms covalently bonded to carbon atoms do not participate in hydrogen bonding
because carbon is only slightly more electronegative than hydrogen & thus the C- H bond is
only very weakly polar. Distinction explains why Butanol has relatively high boiling point of
117oC, whereas butane (CH3(CH2)2CH3) has a boiling point of only -0.5 o C.
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