CHEM 326 Organic Chemistry I MWF 12:00 pm - 12:50 pm, T 8:00 am – 8:50 am CH 101

1. CHEM 326 Organic Chemistry I MWF 12:00 pm - 12:50 pm, T 8:00 am – 8:50 am CH 101 Instructor:Dr. Grigoriy Sereda Office Hours: CH 108 MWF 11:00 am – 12:00 pm Tel. 677-6190 e-mail: gsereda@usd.edu Web: http://www.usd.edu/~gsereda

2. Common knowledge: Organic reactions take place and they are useful in everyday life The purpose of the course: answer the following questions: • Which types of organic reactions may occur? (Classification and nomenclature) 2. How do these reactions proceed? (Reaction mechanisms) 3. How to make use of these reactions? (Organic synthesis)

4. Why to take this course: 1. You will learn many practical applications of organic chemistry, that will be very useful in you career in chemistry or chemistry-related areas such as medical science. 2. Organic chemistry is a very important part of culture, because it expands the Human’s ability to explore the world, change the world, and develop intellectually.

5. Features of this course: 1. Material is built up from one topic to another 2. Unlimited help from me 3. No curving. Your grade will not depend on performance of other students in the class either way To enjoy the course: 1. Ask me for help if something is even slightly not clear 2. Do your homeworks To get grumpy with the course: 1. Expect "bailouts" 2. Ignore my help 3. Trust rumors from students who failed in the past.

6. 1. Chemical Structure and Chemical Bonding Required background: General chemistry Essential for: 1. Writing organic structures and reactions 2. Writing and understanding organic structures and reaction mechanisms 3. Predicting of molecular geometry and charge distribution in a molecule 4. Critical application of scientific theories

7. Outline 1. Subject of organic chemistry 2. Chemical formulas and chemical equations 3. Reaction schemes 4. Electrolytic dissociation 5. Surfaces of potential energy 6. Electronic structure of atoms 7. Chemical structures in organic chemistry

8. 1. Subject of organic chemistry Organic chemistry- is chemistry of most of the compounds, containing carbon. Compounds with at least one carbon - carbon bond - are for sure organic. Compounds with one carbon are considered as organic sometimes. The term “Organic” was applied to substances, isolated from living things by Jons Berzelius (Beginning of the 19th century). These compounds were though to be arisen from a “Vital force”, responsible for the process of life. 1828 - Friederich Wohler converted inorganic ammonium cyanate to organic urea. 1856-1863 - Marcellin Berthelot synthesized organic acetylene from inorganic compounds The theory of vital force was ruled out.

9. Organic compounds are champions in many respects: 1. Most recognizable by smell (found in vine): 2. Most recognizable by taste (2 mg/100 million liters of water):

10. 3. The most toxic: Aflatoxin B1: LD50 = 0.4 mg/kg (Compare: NaCN: LD50 = 15 mg/kg) 4. One of the largest number of polymorphs (different crystalline forms of the same compound): (JACS, 2000, 122, 585) #1 - Red prisms #2 - Yellow prisms #3 - Orange plates #4 - Orange needles #5 - Yellow needles #6 - Orange-red prisms

11. Outline 1. Subject of organic chemistry 2. Chemical formulas and chemical equations 3. Reaction schemes 4. Electrolytic dissociation 5. Surfaces of potential energy 6. Electronic structure of atoms 7. Chemical structures in organic chemistry

12. 2. Chemical formulas and chemical equations Formula unit - is a molecule (for compounds, consisting of molecules) or the smallest fragment with the same ratio between atoms, as in the compound. Chemical formula shows how many atoms of each type are present in the formula unit. Examples: • Na2SO4 means: “One formula unit of the compound Na2SO4contains two atoms of sodium, one atom of sulfur and four atoms of oxygen”. 2, 1, and 4 – chemical indices 2. C2H6O means: “One formula unit of the compound C2H6O contains two atoms of carbon, six atoms of hydrogen and one atom of oxygen”.

13. Chemical equations show how many formula units react and how many formula units are formed in the reaction. The number of each type of atoms in the left part of the equation must be equal to the number of atoms of the same type in the right part of the equation. Examples: CH4 + 2O2 = CO2 + 2H2O means: “One molecule of the compound CH4 reacts with two molecules of the compound O2 and produces one molecule of the compound CO2 and two molecules of the compound H2O”. 1, 2, 1 and 2 – chemical coefficients

14. Outline 1. Subject of organic chemistry 2. Chemical formulas and chemical equations 3. Reaction schemes 4. Electrolytic dissociation 5. Surfaces of potential energy 6. Electronic structure of atoms 7. Chemical structures in organic chemistry

15. 3. Reaction schemes One-step chemical schemes show, which compounds react, which compounds are formed, which compounds and conditions affect the reaction. - At least one organic compound must be present in the left part of the scheme - All other reacting compounds are usually called “reagents” and shown either in the left part of the scheme or over the reaction arrow. - Chemical coefficients are not shown, except if you need to show some stoichiometric features of the reaction. - Some reaction products, especially inorganic ones, are often not shown. Example:

16. If there is only one shown reaction product, for the next step, write the reagents over the reaction arrow. Otherwise the reagents for the next step will be confused with the reaction products for the previous step, which makes the whole reaction scheme wrong (zero points on a test). If there is more than one shown reaction product, there are two ways to write the next step: 1. Start each step from a new line, rewriting the product of the preceding step as the starting material form the next step.

17. 2. Draw the next reaction arrow from the product of the preceding step, which reacts further, and write the reagent over the reaction arrow.

18. Outline 1. Subject of organic chemistry 2. Chemical formulas and chemical equations 3. Reaction schemes 4. Electrolytic dissociation 5. Surfaces of potential energy 6. Electronic structure of atoms 7. Chemical structures in organic chemistry

19. 4. Electrolytic dissociation 1. Acids produce H+ and the anion of the conjugated base. Example: HCl = H+ + Cl- Presentation in reaction schemes: HCl (or another acid), H+, H3O+ 2. Hydroxides of alkali and alkali earth metals produce OH- and the cation of the corresponding metal. Example: NaOH = Na+ + OH- Presentation in reaction schemes: NaOH (or another hydroxide), OH- 3. Salts produce cations of the metal and anions of the rest of the molecule. Example 1: Na2Cr2O7 = 2Na+ + Cr2O72- Presentation in reaction schemes: Na2Cr2O7, Cr2O72- (reactive anion) Example 2: Hg(CH3COO)2 = Hg2+ + 2CH3COO- Presentation in reaction schemes: Hg(CH3COO)2, Hg2+ (reactive cation)

20. Outline 1. Subject of organic chemistry 2. Chemical formulas and chemical equations 3. Reaction schemes 4. Electrolytic dissociation 5. Surfaces of potential energy 6. Electronic structure of atoms 7. Chemical structures in organic chemistry

21. 5. Surfaces of potential energy The potential energy of a molecule (a set of reacting molecules) depends on many geometrical parameters (distances between atoms, angles between bonds, and dihedral angles between planes). Minima on the surface of potential energy are called intermediates. Maxima on the surface of potential energy are called transition states.

22. The surfaces of potential energy are used to describe many molecular events, such as: 1. A reaction proceeds from the intermediate A to the intermediate B through the transition state C. The larger the energy difference between A and C, the slower the reaction. If B has a higher energy, than A, the reaction consumes heat. In the opposite case, the reaction produces heat. Reaction step – is the path of the chemical system from one intermediate (or starting materials) to the next one (or final products). Reaction mechanism – is a sequence of reaction steps.

23. Outline 1. Subject of organic chemistry 2. Chemical formulas and chemical equations 3. Reaction schemes 4. Electrolytic dissociation 5. Surfaces of potential energy 6. Electronic structure of atoms 7. Chemical structures in organic chemistry

24. 6. Electronic structure of atoms The wave-like nature of electrons keeps them from falling on atomic nuclei. 3D-mathematical description of these waves are called orbitals Electronic configuration - is a list of all occupied atomic orbitals in the system with their electron populations Examples:Hydrogen (1e): (1s)1; Carbon (6e): (1s)2(2s)2(2px)1(2py)1 Graphical representation of electronic configurations: Each orbital is represented by a dash at the corresponding level of energy. An electron, characterized by (s) is represented by an up headed arrow (s = 1/2) or a down headed arrow (s = -1/2). Thus an electronic configuration can be represented by arrows-electrons, distributed on (described by) “shelves”- orbitals. Only one or two electrons can be placed on one atomic orbital. Doublet and Octet rules: For elements of the 1st period, a maximum of 2 valence electrons (1s2) may be located around the nucleus. For elements of the 2nd period a maximum of 8 valence electrons ((2s)2 (2p)6 ) may be located around the nucleus.

25. Outline 1. Subject of organic chemistry 2. Chemical formulas and chemical equations 3. Reaction schemes 4. Electrolytic dissociation 5. Surfaces of potential energy 6. Electronic structure of atoms 7. Chemical structures in organic chemistry

26. 7. Chemical structures in organic chemistry The Structural theory in organic chemistry was developed by Kekule, Couper and Butlerov in 1861. According to this theory, atoms in organic molecules are connected by chemical bonds in a certain order, according to their valency (number of bonds an atom can form), rather than randomly clumped with one another. 3D-mathematical description of electronic waves in molecules are called molecularorbitals Only one or two electrons can be placed on one molecular orbital According to the valence bond approximation, each covalent bond is provided by a pair of electrons (bonding electrons) acting as “glue” between two bounded atomic nuclei. Some molecular orbitals are mostly located outside of the “between nuclei” areas. Electrons on those orbitals most often belong to a certain atom, do not contribute to covalent bonds, and are called “lone (non-bonding) electron pairs”. Lewis structure – is a chemical structure, where each bonding or non-bonding electron pair is represented by a pair of dots, and each single electron is represented by a single dot.

27. Different elements have different abilities to hold electrons around their nuclei. Such ability is called Electronegativity In the Periodic table, electronegativity of elements increases from the left to the right and less prominently - from the bottom to the top. What is especially important for us: O>N>C>H

28. Actual and Formal Charges Actual charges are real and can be obtained by physical measurements. The negative charge goes to more electronegative atoms, the positive charge goes to less electronegative atoms. For example, in NH4+, the actual positive charge is located on less electronegative hydrogens. Formal Charge – is the difference between the nuclear charge, corresponding to the valence electrons and the number of electrons, contributing to the atomic charge. One way to figure out the formal charge is to write a reaction scheme, or it can be calculated, using the definition.

29. Rules for Writing Structural Formulas • Always show the formal charge 2. Don’t show non-valence electrons 3. Show unshared electrons, if it’s essential for what you are writing the structure for. 4. If you have shown any of shared or unshared electrons, you must show all unshared electrons. 5. The doublet for the elements of the 1st period and the octet for the elements of the 2nd period may not be exceeded, but not necessarily complete.

30. Molecular geometry 1. Bond lengths depend on which atoms are connected and how they are connected. It’s explained well in your textbooks. 2. Bond angles. The rule for prediction: All electron pairs (shared or unshared) around a nucleus are arranged so that they are as far from one another as possible. (Exception: double bonds behave like one electron pair, because those pairs can not be separated). Limitations: There must not be unpaired (single) electrons in the system. Possible Cases:

31. Physical methods to determine molecular geometry: X-Ray crystallography – in the solid state Electron diffraction and Microwave spectroscopy – in the gas phase

32. Hybridization Depending on the geometrical configuration around the atom (positioning of electron pairs), more than one atomic orbital may be significantly involved in the formation of a chemical bond or a lone electron pair. Numbers of the involved atomic orbitals are described in terms of hybridization. For instance, tetrahedral configuration such as in methane involves one s- and three p- orbitals, which corresponds to the sp3-hybridization. Hybridization is not a real process, but rather a mathematical procedure that derives hybrid orbitals (green on this picture by the linear combination of atomic orbitals). Since chemical bonds described in terms of hybridization share atomic orbitals, they significantly influence each other. The concept of hybridization was suggested by Pauling in 1928 and won him Nobel prize in 1954

33. Resonance Structures Since Lewis structures are based on electron pairs arbitrary assigned to pairs of atoms, they often do not describe the molecule correctly. We need to describe such problematic structures with a set of structures rather than with a single structure. Such set of Lewis structures is called a set of resonance structures. Don’t confuse chemical equilibrium and resonance structures. The actual molecule is planar, but doesn’t have such strong charge separation as would be predicted by the structure on the right.