Nature of Molecules and Intermolecular Interactions

Nature of molecules and intermolecular interactions Covalent Bond (GC) • Lewis Electron Dot formulas: ▪ Formulas are drawn in such a way that an octet is achieved on each atom. Exceptions include boron column (form 3 bonds and have a six-tet), large elements (3rd row and below such as the 10-tet P in PO43- and the 12-tet S in SO42-), and radicals (compounds with an odd # total electrons that result in a single, unpaired electron) ▪ Rules of thumb • Carbon: 4 bonds total (meaning 4 total bonds, can be either 4 single bonds or two double bonds) and no lone pairs e.g. CH4, CO2 • Oxygen: O can be o O: 2 bonds total, 2 lone pairs e.g. H2O, O2 o O1-: 1 bond, 3 lone pairs, formal charge of -1 o O1+: 3 bonds, 1 lone pair, formal charge of +1 • Nitrogen: N can be o N: 3 bonds total, 1 lone pair e.g. Amine or ammonia NH3 o N+: 4 bonds, 0 lone pair, formal charge of +1 e.g. NH4+ • Halogens: 1 bond, 3 lone pairs. e.g. CCl4 • Hydrogen: 1 bond, 0 lone pair (exception to octet rule) • Carbocation: C+ has 3 bonds, no lone pairs, formal charge +1 • Carbanion: C- has 3 bonds, 1 lone pair, formal charge -1 • Boron: 3 bonds, 0 lone pairs (exception to octet rule). e.g. BH3 o Resonance structures ▪ When there are more than 1 satisfactory Lewis structures for a molecule, they are called resonance structures ▪ You can visualize the molecule “shifts” between each of its resonance structures really fast, spending more time in the more stable resonance structures. ▪ Stable properties: • Octet rule is satisfied in every atom (except for boron group and hydrogen) • No formal charges • If there must be formal charges, like charges are apart and unlike charges are close together o Formal charge ▪ Formal charge = valence electron # in the unbonded atom - electron # in the bonded atom ▪ Electron # in the bonded atom = dots around the atom + lines connected to the atom ▪ Common formal charges: • Oxygen with only a single bond: -1 • Oxygen with no bond but have an octet: -2 (oxygen usually exists as the diatomic O2 and have a double bond to themselves) • Carbon with only 3 bonds: either +1 if carbocation or -1 if carbanion • Nitrogen with 4 bonds: +1 • Halogen with no bonds, but have an octet: -1 (Halogens usually exist as a diatomic and have a single bond to themselves such as Cl2) • Boron with 4 bonds: -1. e.g. BH4o Lewis acids and bases ▪ Lewis acid accept electron pairs. They don’t have lone pairs on the central atom. e.g. BF 3 ▪ Lewis base donate electron pairs. They have lone pairs on their central atom. e.g. NH 3 • Partial ionic character o Covalent bonds between atoms with dissimilar EN’s have a partial ionic character • o Role of electronegativity in determining charge distribution ▪ The more EN atom receives a partial negative charge ▪ The less EN atom receives a partial positive charge o Dipole Moment ▪ Molecules with asymmetrical partial charge distribution have a dipole moment e.g. H2O has a dipole moment because the molecule is bent and the oxygen-side of the molecule is partially negative ▪ Dipole moment depends on charge and distance • The greater the EN difference, the greater the charge and hence the dipole moment • The greater the distance separating the charges, the greater the dipole moment ▪ Molecules with symmetrical partial charge distribution do not have dipole moments. e.g. CCl4 does not have a dipole moment because the partially negative chlorine atoms are arranged symmetrically in a tetrahedron. The symmetry cancels out their individual dipole moments ▪ Things with a dipole moment are said to be polar σ and π bonds o Hybrid orbitals: sp3, sp2, sp and respective geometries ▪ Hybrid orbitals: combination of sigma and pi orbitals. Average out characteristics of original orbitals. Equal energy levels • Sigma bonds formed in overlap of hybrid orbitals, pi bonds formed in overlap of p orbitals ▪ Determining hybrid orbitals formed by an atom: count number of sigma bonds and lone pairs. Match this number up to sum of superscripts in the hybrid name (sp2, sp3, etc.) ▪ Orbital character: superscripts determine what percent character the orbital is. sp is 50% s, 50% p. sp2 is 33% s, 66% p • The more s character = the more stable • Pi bonds are less stable than sigma bonds o Valence shell electron pair repulsion and the prediction of shapes of molecules (e.g., NH3, H2O, CO2) ▪ VSEPR (valence shell electron pair repulsion) theory: Electrons in an orbital minimize their energy by moving as far away from other electrons as possible, causes bond angles and geometric conformations ▪ Lone pairs are ignored when determining molecular shapes o Structural formulas for molecules involving H, C, N, O, F, S, P, Si, Cl o Delocalized electrons and resonance in ions and molecules ▪ • Resonance structures result from electrons not being fixed in position (that’s why you “push” electrons when drawing resonance structures ▪ When electrons are not fixed in position, they are delocalized electrons ▪ For all practical purposes, resonance and electron delocalization mean the same thing ▪ In ions, resonance and electron delocalization occurs to “distribute” the charge around ▪ In molecules, resonance and electron delocalization occurs in aromatic rings and conjugated double bonds Multiple bonding o Effect on bond length and bond energies ▪ Multiple bonding decreases bond length ▪ Multiple bonding increases bond energy o Rigidity in molecular structure • ▪ Multiple bonding increases rigidity in molecular structure ▪ Single bonds can rotate but double and triple bonds can’t ▪ Even partial double bonds like those found in the peptide bond prevents free rotation Stereochemistry of covalently bonded molecules (OC) o Isomers: same molecular formula, different structural formula ▪ Structural (constitutional) isomers • Same molecular formula, but different connectivity • Positional isomers: structural isomers that have same functional groups positioned differently • Functional isomers: structural isomers that have different functional groups ▪ Geometric isomers • Same molecular formula, same connectivity, different orientation across double bond • When both sides of the double bond contain the same 2 groups, then cis and trans are used • Cis = same side, Trans = opposite sides • When different groups are attached to either side, Z and E is used • Z is when higher priority groups same side across double bond • E is when higher priority groups on different sides across double bond ▪ Stereoisomers (e.g., Diastereomers, enantiomers, cis/trans isomers) • 2# chiral - #meso = different stereoisomers • Enantiomers are mirror images of each other. That means ALL chiral centers in one enantiomer is reversed in the other • You can’t have stereoisomers if you don’t have a chiral center • Diastereomers - more than one chiral center, inversion of stereochemistry on some but not all of its chiral centers. e.g. (R)-(R) vs (R)-(S) • Meso compounds may have chiral centers, but as a molecule, they are achiral and optically inactive • Meso compounds reduce the total number of stereoisomers • Stereoisomers have the same chemical properties • Enantiomers have the same physical properties • Diastereomers have different physical properties • Note: in biological molecules, people use D and L for R and S, respectively ▪ Conformational isomers: have some molecular formula, same connectivity, same stereochemistry, but can rotate about a single bond to switch between different conformations • Technically, conformational isomers are not really isomers because you don’t have to break any bonds to convert from one conformation to another. They are more accurately called conformers • Conformers about a single bond o Eclipsed ▪ Syn-periplanar: highest torsional strain, most unstable, bulky groups eclipse each other ▪ Anticlinal eclipsed: high torsional strain, unstable, bulky groups eclipse hydrogens o Staggered ▪ Gauche: low torsional strain, stable, bulky groups 60˚ staggered ▪ Anti: lowest torsional strain, most stable, bulky groups 180˚ staggered o Single bonds will rotate such that it achieves the most stable conformation • Conformers of cyclohexose o Chair: most stable, everything is staggered o Twist boat: less stable, things are not completely eclipsed o Boat: least stable, everything is eclipsed o Hexose rings will twist and turn to achieve the most stable conformation • Torsional strain: the strain due to eclipsing of groups across a single bond • Steric interactions o Axial: most unstable because axial groups are orientated with a high degree of clashing o Equatorial: most stable because the equatorial groups are orientated away from one another o Bulky groups like to be in the equatorial position • Most stable conformation: completely staggered (chair), with bulky groups in the equatorial position • Least stable conformation: completely eclipsed (boat), with bulky groups in the axial position o Polarization of light, specific rotation ▪ Light is an electromagnetic wave; they are waves of electric and magnetic fields (in phase, but perpendicular to each other and also to the direction of propagation) ▪ Normal light has the EM fields in all directions (in a 360˚ circle perpendicular to the direction of propagation) ▪ Polarized light has EM fields all in one direction ▪ Specific rotation: chiral molecules containing a single enantiomer will rotate polarized light (to varying degrees) either to the left or to the right. This is why chiral molecules are said to be “optically active” ▪ Left rotation: (-) or l or levorotatory ▪ Right rotation: (+) or d or dextrorotatory ▪ Caution: (+) or (-) does NOT correspond to R/S configurations ▪ Caution: d and l is NOT the same as D and L. Upper case letters denote absolution configurations in sugars o Absolute and relative configuration • Steps in assigning (R) and (S) o a) Is the carbon center chiral? o b) Assign priorities according to Cahn-Ingold-Prelog rules o c) Turn molecule such that lowest priority group is at the back o d) Rotate from 1st to 2nd to 3rd priority group like a steering wheel. It’s (R) if you end up turning right and (S) if you end up turning left • Relative configuration is defined in relationship to another chiral center. The direction that a molecule rotates plane-polarized light is the prime example ▪ Conventions for writing R and S forms • If only 1 chiral center o (R/S)-molecule, where R/S is the absolute configuration and molecule is the name of the compound • If more than 1 chiral center o (#R/S, #R/S)-molecule, where # is the carbon number (in ascending order), R/S is absolute configuration, and molecule is name of compound ▪ Conventions for writing E and Z forms • If only 1 double bond o (E/Z)-molecule, where E/Z is the geometric configuration across the double bond, and molecule is the name of the compound • If more than 1 double bond ▪ o (#E/Z, #E/Z)-molecule, where # is the carbon number (the smaller one in the double bond, in ascending order), and molecule is the name of the compound Cahn-Ingold-Prelog rules for assigning priority • Assign atom with higher MW greater priority • e.g. CH(OH)2 vs CH2F o Doesn’t matter how many oxygen atoms there are, because fluorine has greater MW so fluorine has higher priority Liquid Phase - Intermolecular Forces (GC) • Hydrogen bonding o Weak interaction between a partially positive H and a partially negative atom o Technically, hydrogen bonds are a special type of dipole-dipole interaction o Hydrogen bonding increases the boiling point o Partially positive H are also called hydrogen bond donors. They are hydrogens that are bonded to either F, O, or N o Partially negative atoms are also called hydrogen bond acceptors. They are most commonly F, O, or N o Do ethers form hydrogen bonds with other ethers? No, because ethers do not have a partially positive H (donor) o The more polar a bond is, the stronger the hydrogen bond. The H-F bond is the most polar, followed by the H-O bond, and lastly the H-N bond • Dipole Interactions o All polar molecules exhibit dipole-dipole interactions. This is where the polar molecules align such that opposites attract o Dipole-dipole interactions increase the boiling point, though not as significantly as hydrogen bonding. o Dipole interactions are stronger the more polar the molecule is o Ion-dipole interactions are similar to dipole-dipole interactions but it’s stronger because it is no longer an interaction involving just partial charges. Instead, it is an interaction between a full charge (ion) and a partial charge (dipole) o Ion-dipole interactions get stronger when you have larger charge magnitude of the ion, and large polarity of the dipole molecule • Van der Waals’ Forces (London dispersion forces) o Also called dispersion forces o Dispersion forces exists for all molecules but are only significant for non-polar molecules. For polar molecules, dipole forces are predominant o Dispersion forces result from induced and instantaneous dipoles o Induced dipoles: when a polar molecule interacts with a non-polar molecule, then polar molecule induces a dipole in the non-polar molecule o Instantaneous dipoles: Non-polar molecules have randomly fluctuating dipoles that tend to align with one another from one instant to the next o Dispersion forces get stronger for larger molecules. For example, decane (C10H22) has a stronger dispersion force than ethane (C2H6)