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Bond line structures are useful and easy to read, however they are somewhat inadequate due to the mobility of electrons. Because electrons have no fixed location within a molecule or compound they can be delocalized throughout the structure; however, some spaces are more likely to have a higher electron density than others.

As a result organic chemist generally look at something called resonance. Resonance structures are used when one (Lewis) Structure cannot fully explain the bonding that occurs between the atoms. 2+ high energy structures are considered a part of the resonance hybrid which shows the overall delocalization. The more resonance structures a molecule has, the more stable it is. It is important to note that resonance structures only change the position of the electron not the structure of the molecule.

Delocalization and Resonance Rules

Structures should have the same number of electrons and lone pairs (LP)- no more, no less.

Resonance structures must follow the rules of Lewis Structures

Hybridization must remain the same

Only the electrons move - do not rearrange bonds

Resonance Hybrids

When drawing Resonance structures, use solid lines for bonds

A dotted line can be used to show a double or triple bond present

Draw only the LP found in all resonance structures

Stability Rules

Greater number covalent bonds = Greater stability due to complete octets

Most stable structure has the least number of formal charges

Stable structures have the least separation of formal charge

Negative charges on the atom with the higher electronegativity is more stable and positive charges on atom with the higher electropositivity is more stable.

Equivalent resonance structure attribute equally to stability.

Forward Arrow: reaction arrow, this is often used to show reactants forming into products.

Equilibrium Arrow: arrow that shows a reaction to be reversible; one arrow can be exaggerated in the more favored direction of EQ.

Resonance Arrow: shows that there is a different arrangement of electrons.

Curved arrows: show movement of electrons, either one or two. Used largely in mechanisms.

Broken arrows: used to show a reaction does not work.

Open Arrow (Retrosynthesis): shows the decomposition of a complex into simple starting material. Predominately used as a planning device for synthesis.

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The following will be gone over in greater detail at a later date. Right now this is just showing how molecular representations can be drawn as.

Thicker lines indicate a wedge for all molecular representation, this tells you that the atoms is point at you. Dashed lines indicate that the molecule is point away from you and mean the same for all molecular representation.

Fischer Projections: convention used to depict a stereoformula in 2D without destroying the chirality. As a result there are limitations with what can be done with fischer projections.Note that each intersection has a carbon present.

Hawthorne Projections: usually used for sugars and shows a 3D view of the molecule. In these structures C and H atoms are implicit.

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Partially Condensed: C-H bonds are not explicitly drawn and is only practical for small molecules.

Condensed: single bonds are not drawn and atoms are clustered together based on groups when applicable [e.g the 2 CH3 groups in isopropyl groups are written as (CH3)2.]

Molecular formula: does not provide structural details (e.g one formula can be long to multiple constitutional isomers which differ structurally).

bond lines: zigzag drawings were each point is a carbon and hydrogens are implied.

Valence Shell Electron Pair Repulsion (VSEPER): proposes that the geometric arrangement in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom. Repulsion causes electron pair to arrange themselves in 3D space to achieve maximal distance.

Steric Number: (Total Number of Atoms) + (Number of Lone Pair).

Cheat: Count the number of atoms attached to the central atom and the number of lone pairs (LP).

Steric Number 4 - Tetrahedral

0 LP - Tetrahedral geometry

1 LP - Trigonal Pyramidal geometry

2 LP - Bent geometry

Steric Number 3 - Trigonal Planar

1 LP - Bent Geometry

0 LP - Trigonal Planar

Steric Number 2 - Linear

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Hybridization: expansion on the valence bond theory and is essentially the notion that atomic orbitals fusing together to form newly hybridized orbitals, which can impact the molecular geometry and bonding properties of the compounds.

sp3: After the promotion of one of the 2s electrons there is a mixing of 3 p and 1 s orbitals. All four hybrid orbitals occupy the same energy level; however, they are lower in energy than the 2p orbital but higher than the 2s.

sp2: After the promotion of one electron in the s orbital to one of the 2p orbitals there is mixing of 1 s and 2 p orbitals. These equal energy orbitals are higher in energy than the 2s but lower than the 2p. The overlap of these orbitals results in sigma and pi bonds (sp2 orbital overlap is predominately in double bonds for organic chemistry).

sp: After the promotion of one electron in the s orbital to one of the 2p orbitals, there is mixing of 1 s and 1 p orbitals. The overlap of these orbitals results in sigma and pi bonds (sp orbital overlap is predominately in triple bonds for organic chemistry).

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Atomic Orbitals: region of space that electrons occupy. Due to the mobility of electrons, it is easier to consider electrons as a cloud instead of orbitals; however, orbitals--particularly s, p, and d--are often used to draw bonding interactions. You will remember in general chemistry that they possible orbitals are: s, p, d, f, g, h, i, j, etc. However for organic chemistry s and p orbitals are particularly important in bonding.

s orbitals: A spherically symmetric orbital around the nucleus of the atom. The further an electron is from the nucleus the larger the orbital becomes and the higher in energy the electron is. As a result, the energy trend goes from 1s < 2s < 3s < 4s...etc. S orbitals usually make up sigma bonds.

p orbitals: p orbitals are oriented in a specific direction - most commonly expressed as x, y, and z. p orbitals only start in the second energy level: 2p, 3p, 4 p, etc. P orbitals make up double and triple bonds (also known as pi bonds).

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Intermolecular forces: electrostatic attractive forces between molecules.

Dipole-Dipole: net dipole moment can attract or repel based on the partial charges. (See photo above)

Hydrogen Bonding: specific type of dipole-dipole. When connected to a electronegative atom H has a partial positive charge. When it is connected to a metal, it has a partial negative charge. The effect of hydrogen bonding can be dramatic because H is mall and therefore can get closer to other atoms.

London Dispersion forces: transient (moving) dipoles shift around due to the fact that electrons are constantly in motion. This movable dipoles can induce dipoles in neighboring molecules. These interactions are stronger with the larger the surface area.

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Principles to remember when filling Atomic orbitals

Aufbau: lowest energy level is filled first. 1s before 2s.

Pauli Exclusion: 2 electrons max for each orbital. Furthermore, the electrons must have opposite spin.

Hund’s Rule: one electron is placed in each degenerate orbital before you can pair up electrons.

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Formal Charge (FC): associated with any atom that does not exhibit the appropriate number of valence electrons. This means that the valence electrons (VE) are not what they normally are. +VE means that electrons were donated and -VE means that electrons were attained from another atom. FC of atoms are a useful for electron book keeping and helpful in determining reactivity of a compound. If you struggle to determine formal charge you can draw the lewis dot structure which can help you visualize it better. With enough practice one can determine Formal charge at a glance.

In neutral compounds, there can be individual atoms with charges. The molecule remains neural when the overall charge is 0. This means that the charges canceled each other out.

Charges are useful for insight on or predicting reactivity.

Unless you have a metal involved, the FC will either be neutral or +1/-1

There are two equations you can use to determine FC. This equations include:

FC = (Group Number) - (#electrons in Covalent bonds) - (#LP)

Group Number is the number of VE in the neutral atom

FC = (Group Number) - (#VE around atom in molecule)

This is the equation used above

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Lewis Structures are all about electrons. Each dot indicates a valence electron (electrons in the outer most orbital of the atom). If there are more than 4 valence electrons, you will need to pair them up (N-Ne groups).

Octet Rule: 8 is the magic number for Lewis Dot Structures. All 8 valences are in pairs. Aside from Nobel gases all atoms need to bond with other atoms to pair up their valence electrons. In the case of H-B the atoms have an exception to the octet rule, namely that they do not not pair so that that they have 8 valence electrons. In their neutral form: H is content with 2, Helium is content with 4, and B is content with 6. These atoms, however, will still pair to fill their orbitals, they just will not fill up to 8.

Unshared electron pairs are called lone pair. They do not participate in bonding with other atoms in Lewis Dot Structures. Lone pair in organic reactions however can act as nucleophiles to form bonds.

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Electronegativity (EN): measures of the atom’s ability to attract electrons.

Golden Rule: Fluorine is the most electronegative atom. Therefore the easiest way to remember the electronegativity trend is remembering where Fluorine is on the table (top right). Electronegativity increases to the right and up a column

.Bonds based on electronegativity:

Covalent bonds: two atoms share a pair of electrons. Way to remember is that “co” means jointly associated. Covalency is often greater between atoms with similar electronegativity. This sharing of electron density over more than two electrons is considered delocalization. EN < 0.5

Polar Covalent Bonds: In Polar Covalent bonds, electrons are shared unequally with the electron density being pulled towards the more electronegative atoms. EN between 0.5 -1.7.

Ionic bonds: The more electronegative atoms takes all of the electrons from the electropositive atom. In general, the electropositive atom is usually a metal which loses electrons to become cation, whereas the nonmetal accepts those electrons to become an anion. EN > 1.7

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Valence electrons are electrons in an atom's valence shell which can participate in bonding.

Valence bond Theory: Valence Bond theory describes covalent bond formation as well as the electronic structure of molecules. The theory assumes that electrons occupy atomic orbitals of individual atoms within a molecule and that the electrons of one atom are attracted to the nucleus of another atom. This theory is not actually accurate because it is the delocalized electrons that really participate in bonding. VB is not as accurate as Molecular Orbital (MO) theory but it is easier to conceptualize.

This attraction increases as the atoms approach one another until the atoms reach a minimum distance where the electron density begins to cause repulsion between the two atoms. This electron density at the minimum distance between the two atoms is where the lowest potential energy is acquired, and it can be considered to be what holds the two atoms together in a chemical bond.

Remember that electrons can behave like a wave and a particle so they have high mobility. When these electrons are being shared in bonds, there is an overlap in atomic orbitals

There are two types of interference prevalent that allow for bond sharing:.

Constructive Interference: large-amplitude wave (bond sharing)

Destructive interference: wave cancels out

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