1. Structure and Bonding

A Review of Needed Material

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Organic Chemistry

“Organic” – until mid 1800’s referred to compounds from living sources (mineral sources were “inorganic”)

Wöhler in 1828 showed that urea, an organic compound, could be made from a inorganic materials

Organic compounds are those based on covalently bonded carbon and study of their structures and reactions.

NH4Cl AgNCO H2N-C-NH2

OAgCl+ heat +

Ammoniumchloride

Silvercyanate

Urea Silverchloride

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Atomic Structure

Structure of an atom Nucleus = assigned a positive charged, very dense,

protons and neutrons in it and small (10-15 m), deflects an alpha particle shot at it.

Electrons = therefore are negatively charged, occupy allowed energies (sometimes called an electron cloud) (10-10 m across) around nucleus … and are not “orbiting” the nucleus … Ahhhhh see page 4!

Diameter is about 2 10-10 m (200 picometers (pm)) [the unit angstrom (Å) is 10-10 m = 100 pm]

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Isotopes

Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers

Nuclear “spin” occurs when an odd number of nuclear particles are present.

The spin has a specific energy when placed in a magnetic field.

This energy is characteristic of the isotope and its environment … and can be detected, characterized, and used to determine which atoms are bonded to each other in a pure substance! NMR - Yaaahooo

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Atomic Structure: Orbitals

Quantum mechanics: describes electron energies by a wave equation A Wave function, , is a solution of wave equation A plot of 2 describes probable electron density The Shapes that come from all this math are

called orbitals ( 2 ). We pick the solutions that work and through the rest out.

We can mathematically combine these solutions to make other orbitals at our whim – again, we use the ones that work and leave the rest out:

HYBRID ORBITALS!

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It’s the Filled Orbitals – Stupid! Orbitals have an energy order to them … I hate the word “SHELL” Each primary quantum number denotes increasing energy of the

electrons within them and contain different numbers and kinds of orbitals

Each orbital can be occupied by two electrons n =1 contains one s orbital, denoted 1s, holds only two electrons n = 2 contains one s orbital (2s) and three p orbitals (2p), eight

electrons n = 3 contains an s orbital (3s), three p orbitals (3p), and five d orbitals

(3d), 18 electrons

Eliminate your idea of “filled shells” because n = 3 shell is very happy not being filled and only having 8 electrons. It is the filled orbitals that cause stability. Filled s and p orbitals are the ones we are most interested in with organic chemistry.

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Shapes of Atomic Orbitals

Four different kinds of orbitals are useful to discuss s, p, d, and f s and p orbitals most important in organic chemistry

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Chemical Bonding Theory

Observation: Carbon makes four bonds and has yet to be observed to ever have 5 bonds that are isolable.

Observation: 2-chlorobutane has two isomers

Explanation: Carbons with 4 bonds = tetrahedral

Note that a wedge indicates a bond is coming forward

Note that a dashed line indicates a bond is behind the page

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Chemical Bonds

Atoms form bonds because the compound that results is more stable than the separate atoms.

Bond Energy =

Bond Length

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Valence Bond Theory

Valence electrons are what make bonds. …

Lewis structures show valence electrons of an atom as dots

First bond cylindrically symmetrical, sigma () bond

Second and third bonds … called pi () bonds …

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Number of Covalent Bonds

Atoms with one, two, or three valence electrons form one, two, or three bonds and the empty p orbital

does not bond directly but is available to accept electron density – Lewis Acid!

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Number of Covalent Bonds

Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s

and p orbitals = octet rule

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Valence Electrons of Boron

Boron has three valence electrons (2s2 2p1), forming three bonds (BF3)

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Valence Electrons of Carbon

Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4)

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Valence of Nitrogen

Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3)

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Valence Electrons of Oxygen

Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O)

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Hybrid orbitals – Reorganization of orbital energies! In CH4, all C–H bonds are identical (tetrahedral) Each C–H bond has a strength of 438 kJ/mol and bond

length of 110 pm Bond angle: each H–C–H is 109.5°, the tetrahedral

angle.

Explanation: sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent orbitals (sppp = sp3), Pauling (1931) arranged in a tetrahedral array.

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Pictures – hybrid orbitals

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Bonding Ethane

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Bonding Ethene

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Bonding in Ethyne

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What about lone pairs?

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Molecular Orbital Theory

Helps us render the idea that the electrons are lowering in energy when bonds form. Best in 2 electron scenerios!

Bonding

Anti-bonding

But really … I have never had to use these ideas to remember or derive anything in all 10+ organic chemistry classes I have taken or with any of the organic synthesis I have accomplished.2 patents32 anticancer compounds synthesized4 new synthetic pathways7 novel reactions yielding >93% yield from others work

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What works!

• Thermodynamics, • Bond Energies - Will it even work• Equilibrium – Can you get enough of it

• Kinetics• Is the energy barrier to high?• Which of the “X” possible reactions will win?

• Polarity in molecules• Who likes whom?• Consonant vs. Dissonant synthesis

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Summary Organic chemistry – chemistry of carbon compounds

Atom: positively charged nucleus surrounded by negatively charged electrons

Electronic structure of an atom described by wave equation Electrons occupy orbitals around the nucleus with specific energies. Different orbitals have different allowed energies and different shapes

s orbitals are spherical, p orbitals are dumbbell-shaped

Covalent bonds - electron pair shared between two atoms are lower in energy

Valence bond theory – Valence electrons to the bonding

Sigma () bonds - Circular cross-section and are formed by head-on interaction

Pi () bonds – “dumbbell” shape from sideways interaction of p orbitals

Atoms use hybrid orbitals to form bonds in organic molecules. tetrahedral geometry has four sp3 hybrid orbitals planar geometry uses three equivalent sp2 hybrid orbitals and one unhybridized p orbital linear uses two equivalent sp hybrid orbitals with two unhybridized p orbitals