Understanding Dipole Moments, Hydrogen Bonding, and Protein Structure in Aqueous Solutions, Study notes of Biology

Download Understanding Dipole Moments, Hydrogen Bonding, and Protein Structure in Aqueous Solutions and more Study notes Biology in PDF only on Docsity! BIBC 100 Handout 1 10-2-02 Andrew Hires Aqueous Solutions, Amino Acids & Secondary Protein Structure Dipole Moments: Why do they exist? Not all atoms are created equal. Some like electrons more than others (Each atom has a different sets of orbitals for electrons to occupy, and each set has slightly different binding energies.) When two atoms interact, they compete for the electrons that orbit around each of their nuclei. Linus Pauling codified the relative measure of electron attracting power, which he termed electronegativity. The higher the electronegativity, the more the atom can pull on orbiting electrons of neighboring atoms. Here is a periodic table style chart of common elements. Electronegativity H = 2.1 X x x x x x Li = 1.0 Be = 1.5 B = 2.0 C = 2.5 N = 3.0 O = 3.5 F = 4.0 Na = 0.9 Mg = 1.2 Al = 1.5 Si = 1.8 P = 2.1 S = 2.5 Cl = 3.0 K = 0.8 Ca = 1.0 Ga = 1.6 Ge = 1.8 As = 2.0 Se = 2.4 Br = 2.8 Rb = 0.8 Sr = 1.0 In = 1.7 Sn = 1.8 Sb = 1.9 Te = 2.1 I = 2.5 Cs = 0.7 Ba = 0.9 Tl = 1.8 Pb = 1.9 Bi = 1.9 Po = 2.0 At = 2.2 If two atoms are identical and form a chemical bond, i.e. H2, they share the 2 electrons between them evenly. But if the two atoms are different, i.e., Lithium and Hydrogen, the more electronegative (hydrogen) sequesters the two available electrons more towards it. Here is a plot of electron densities (in A, solid indicates high density, and mesh indicates the edge of the electron cloud. In B, color indicates partial charge at that point) Here is a plot of water. Oxygen is the second most electronegative element, clocking in a 3.5units, vs. Hydrogen’s 2.1. Hence, oxygen pulls the electrons towards it. Although the overall charge of the molecule is 0, there is a dipole moment, with the oxygen side partially negative, and the hydrogen side partially positive. We call this a polar molecule. H2 is evenly charged all around and is a non-polar molecule. Dipole moments can be calculated for entire molecules, side chains, or individual bonds. δ− δ− 2δ+ What are dipole moments? The dipole moment is a simple expression of how far the electrons are getting pulled from the “center” of the atoms. Instead of talking about each electron’s displacement, we lump them together and model as two equal point charges of opposite sign separated by a distance. The greater the displacement of the electrons, the larger the distance between our two imaginary point charges. µ=Qr where µ is the dipole moment Q is the charge r is the distance For H2O, we get a value for µ of 1.85 Debye, the equivalent of 1 electron’s worth of charge getting pulled ~0.35 angstroms from the center of the water molecule, towards the oxygen. Non-covalent Forces Hydrogen Bonding Since the electrons around the hydrogens in water are getting pulled towards the oxygen, the hydrogens are not fully satisfied. They try to make up for this by attracting electrons from regions of other molecules that are rich in electron density. In liquid water, there are plenty of other electron rich sources, the oxygens of other water molecules! Ice The attractive force between the electron dense oxygen to another molecule’s hydrogen is weaker than the intramolecular chemical bonds, but still significant. This Hydrogen Bonding is strongest and most stable when the hydrogen is located in a straight line between to electron rich regions. Hydrogen bonding occurs primarily to oxygens or nitrogens, not carbons. It can act inter- or intramolecularly, and it governs the structure of liquid water, ice, protein folds and many other intra and intermolecular structures. Van der Waals Force The electrons in a dipole are unevenly distributed. This unevenness can influence the shape of the electron cloud in neighboring molecules, inducing a slight attractive force between the two molecules. This Van der Waal’s force is proportional to 1/r6, while the electrostatic repulsion of electron clouds is proportional to 1/r12. Hence there is a Van der Waal’s radius around every atom, where the attractive and repulsive forces cancel. At low temperature, this force even works on non-polar molecules. The electrons in the cloud around a molecule are in constant flux. Briefly, more electrons may exist on one side of an atom than on the other. This transient dipole can set up other transient dipoles in neighboring atoms leading to a slight attraction. The Van der Waal’s force acts inter- and intramolecularly.