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What is “organic chemistry” and why is it so important? • Organic chemistry is the study of the compounds of carbon. • Think about how organic compounds affect our daily life: • Our clothes – natural and synthetic fibers • Our medicines • Our food – carbohydrates, proteins, triglycerides • Oils, perfumes, paints, plastics, detergents, etc.
What is “organic chemistry” and why is it so important? alizarin: the first naturally occurring dye to be synthesized (1868) indigo: used to dye blue jeans ethanol: a fermentation product
“Vital force” theory Friedrich Wöhler (1828): A slight problem! The ammonium cyanate was synthesized from bones.......but the Kolbe synthesis of acetic acid in 1845 put the theory to rest!
Vitamin B12 The synthesis of vitamin B12 was finally completed in 1972 by Woodward and Eschenmoser after 10 years of work and the assistance of roughly one hundred graduate students.
“A production of amino acids under possible primitive earth conditions” S.L. Miller, Science, 117, 528 (1953) amino acids including glycine and alanine
Proteins In order to give you an idea of the vastness of organic chemistry, we will look at proteins, our final topic of study in CHEM 263! E. coli contains ~5,000 different chemical compounds of which 3,000 are proteins. Man contains ~2,000,000 different proteins. Biologists believe that there are in excess 10,000,000 of proteins which take part in the process of life! http://www.wisegeek.com/how-many-proteins-exist.htm
Angiotensin II Angiotensin II is a blood pressure regulating hormone. It contains 8 amino acid residues. It is possible to arrange these in 40,320 different ways only one of which corresponds to the hormone! Its structure is actually: Asp-Arg-Val-Tyr-Ile-His-Pro-Phe.
Organic chemistry A logical subject based on fundamental principles: Can we predict the reactivity of a group of atoms based on what we’ve learned in school chemistry courses? The answer is YES! You know a great deal about acetone so let us look at it.
Acetone Here is acetone! But…where are the carbons and hydrogens?
Acetone Here is acetone! But…where are the carbons?
Acetone But…where are the hydrogens?
Acetone - a base! Remember the curved-arrow convention. The tail points to the electron source and the head to the electron destination.
Models of chemical bonding These models are based on the premise that atoms react to produce the electronic configuration of a noble gas.
Models of chemical bonding: the ionic bond Consider the Li - F bond. Li 1s2 2s1 - gives He configuration on loss of e- F 1s2 2s2 2p5 - gives Ne configuration on gain of e- “Ions” are the structural units of ionic compounds and have strong electrostatic forces holding them together.
Models of chemical bonding: the covalent bond When atoms of similar electronegativities are bonded, complete electron transfer cannot take place. The noble gas configurations are attained by sharing electrons: Electronegativity - the ability of an atom to attract electrons. In the covalent bond, the atoms share electrons. The structural unit is the “molecule.”
Lewis structures 1. Find the total number of valence electrons of all of the atoms. 2. Use pairs of electrons to form bonds between all bonded pairs of atoms. 3. Distribute the remaining electrons to give each hydrogen a duet and atoms of the second period an octet. e.g. CH4 e.g. C2H4
Lewis structures (CH3)2CHCH2OH • Find the total number of valence electrons of all of the atoms. • 4 x 4 + 10 x 1 + 6 = 32 • 2. Use pairs of electrons to form bonds between all bonded pairs of atoms. 28 electrons
Lewis structures (CH3)2CHCH2OH 3. Distribute the remaining electrons to give each hydrogen a duet and atoms of the second period an octet.
Formal charge It is often necessary to include a formal charge when we draw a structure. Thus in drawing the hydronium ion we need to know where the charge is located. formal = valence - unshared electrons - 0.5 x shared electrons charge electrons in bonded atom in bonded atom in free atom H3O+ - for oxygen, formal charge = 6 - 2 - 0.5 x 6 = +1 What is the formal charge on nitrogen in the ammonium ion?
formal = valence - unshared electrons - 0.5 x shared electrons charge electrons in bonded atom in bonded atom in free atom Formal charge for atom 1 = 6 - 6 (non-bonded e-) - 0.5 x 2 = -1 Formal charge for atom 2 = 6 - 2 - 0.5 x 6 = +1 Formal charge for atom 3 = 6 - 4 - 0.5 x 4 = 0
Resonance Consider the carbonate ion, CO32-. We can draw three equivalent structures: In reality the ion is perfectly symmetric. All C-O bond lengths are identical and the negative charge is delocalized over the three oxygens. The structure is a hybrid of these three contributing structures.
The theory of resonance • Whenever a molecule can be represented by 2 or more structures which differ only in the arrangement of their electrons, there may be resonance: • The molecule is a hybrid of all the contributing structures and cannot be adequately represented by any one of these structures.
The theory of resonance • Resonance is important when these structures are of about the same stability. For example, However: • The hybrid is more stable than any of the contributing structures. This increase in stability is called the resonance energy.
How to draw resonance structures Draw all the resonance structures of:
Quantum mechanics In order to understand covalent bonding, we must start by studying the electronic structure of individual atoms. Schrödinger calculated mathematical expressions which describe the motion of electrons. These “wave equations” give the energy levels available to electrons as well as the relative probability of finding an electron associated with a given energy level at any point in space. Orbitals are 3-dimensional representations of these probabilities.
Atomic orbitals Atomic orbitals are described by three quantum numbers: The most important is the principal quantum number, n. It governs the energy of an orbital. It can be equal to any positive integer except for zero. The second is the angular momentum quantum number, l, whose value depends on n: l = 0, 1, 2, …. n-1. It determines the shape of the orbital. The third is the magnetic quantum number, ml, which governs the orientation of the orbital relative to the three axes. ml = -l, -l + 1 … 0 … l - 1, l.
Atomic orbitals Orbitals are classified (named) according to their values of n and l using a number and a letter. The number represents the value of n and the letter represents the value of l. Electrons in an orbital having l = 0 are called s electrons. Electrons in an orbital having l = 1 are called p electrons. Electrons in an orbital having l = 2 are called d electrons.
Atomic orbitals n = 1, l = 0, ml = 0 one 1s orbital n = 2, l = 0, ml = 0 one 2s orbital n = 2, l = 1, ml = -1, 0, or +1 three 2p orbitals!
1s 2s 2p
Phase signs When the value of a wave equation is calculated for a particular point in space relative to the nucleus, the result may be a positive number, a negative number, or zero. node Y (+) Y (-)
Electron configurations The aufbau principle: Orbitals of lowest energy are filled first. The Pauli exclusion principle: Orbitals can accommodate a maximum of two electrons but only if they are of opposite spin. Hund’s rule: One electron is placed in each degenerate orbital before adding a second electron to an orbital. The electronic configuration of carbon is therefore 1s2 2s2 2p1 2p1
Linear combination of atomic orbitals Representation of the formation of the H-H bond by the sharing of electrons by the two hydrogens and overlap of their singly occupied atomic orbitals.
Molecular orbitals no interaction maximum stability The H-H bond
Antibonding molecular orbitals When two atomic orbitals combine, they form two molecular orbitals. We have met the bonding molecular orbital. Here the atomic orbitals combine by addition. Thus orbitals of the same phase sign overlap. The antibonding molecular orbital is formed by interaction of orbitals of opposite phase sign: