Significance of free electrons vs lone pairs
free electrons - not bound to any specific atom
- essential for electrical and thermal conductivity
lone pairs - pair of valence electrons localized on a single atom, not involved in bonding
- influence molecular geometry
- nucleophiles
- enable intermolecular interactions like hydrogen bonding
- affect molecular polarity and dipole moments
Difference between covalent and non-covalent interactions
covalent - stable
- formed by sharing of outer electrons, due to overlapping orbitals
- strongest bond
- peptide bonds, disulfide bonds
- phosphodiester bonds
non-covalent - greater distance than in covalent bonds
- hydrogen bonds interactions
- electrostatic interactions
- van der waals interaction
- hydrophobic effects
Hydrogen bonds
- H bond forms between the hydrogen atom on the donor and the lone pair on the acceptor
- bond energy = 3-5 kcal/mol
- length = 3
- directional property - importance of orientation
- lone pairs on O and N act as H-bond acceptor
- sulfur on cysteine can act as a weak H-bond acceptor since it has lone pairs
- hydrogen bonding between bases in DNA, amide nitrogen hydrogen and carbonyl oxygen
- water molecules are cohesive, can H-bond to 4 other water molecules (1 for each hydrogen, 2 for each oxygen)
Electrostatic interactions
- salt bridges
- bond energy = 4.8 kcal/mol
- charged amino acid side chains participate in electrostatic interactions to facilitate protein folding
- interaction between positive and negative charge
Electrostatic interaction formula
E=K q1q2/Dr
- larger D = weaker interaction
- water has one of the highest D of any pure liquid
Van der Waals (weak) forces - dipole interactions
- interactions occur between electrically neutral molecules
- polar molecules have permanent dipoles
- non-polar molecules can interact via these forces -> each atom has an electron cloud that fluctuates, yielding a transient electric dipole
- dipole of one atom can influence the electron distribution in another
- london dispersion forces = weakest
- repulsion occurs between atoms that are brought too close to each other
Hydrophobic effect
- not a true force but a consequence of the energy need to insert a nonpolar molecule into water
- water naturally forms a hydrogen bond network, these must be broken to insert nonpolar molecules
- nonpolar form van der waal interaction among themselves and do not H-bond with water (difficult to solvate)
- organization required to dissolve nonpolar solutes (clathrates- form around nonpolar molecules)
Entropy in hydrophobic effects
- arrangement of water molecules around each hydrophobic molecule is thermodynamically unfavorable
- organization of water molecules reduced when hydrophobic molecules associate with each other -> increasing entropy (disorder) of the system (thermodynamically favored)
- “attraction” between non-polar molecules in aqueous environment = hydrophobic effect
- default to lower energy state
Protein folding and hydrophobic effect
- in an unfolded protein nonpolar or hydrophobic regions are exposed to water which organize itself around these regions
- in a folded protein hydrophobic regions often associate with each other and are shielded from the aqueous solvent
- less order is required for the water molecule to solubilize protein -> more favorable energetically
Why is water an excellent solvent
- water is very polar
- forms hydration spheres around ions, proteins, and other molecules
- can solubilize polar compounds by forming H-bonds with these solutes
- ability of water to hydrate ions is great than tendency of opposite ions to attract one another
