Bonding, Structural Formulas, and Molecular Shapes

Bonding, Structural Formulas, and Molecular Shapes

. . . . . . . . . . . . - . 109 109 109 109 . . . . . . . . - . Fundamentals of Organic Chemistry CHEM 109 Fundamentals of Organic Chemistry CHEM 109 For Students of Health Colleges Credit hrs.: (2+1)

King Saud University College of Science, Chemistry Department CHEM 109 CHAPTER 1. INTRODUCTION Syllabus 3 Introduction Types of chemical bonds: (Ionic and covalent bonds) - Atomic and molecular orbital: (sigma and pi bond) - Hybridization (sp3, sp2, sp) - Inductive effect, polarization, and Stability of carbocations - Classification of organic compounds and functional groups - Types of chemical reactions: (Substitution (Free radical - electrophilic - nucleophilic), Elimination, Oxidation and reduction reactions). Lectures (2) Aliphatic Hydrocarbons Classes of hydrocarbons: (saturated and unsaturated) Nomenclature: (IUPAC and common names) Isomerism: (Structural and Geometrical) - Physical properties of aliphatic hydrocarbons - Preparation of saturated hydrocarbons (Alkanes): (Hydrogenation of unsaturated hydrocarbons - Hydrolysis of alkyl Grignard reagent - Reaction of lithium dialkyl cuprates with alkyl halides) - Reactions of saturated hydrocarbons: (Halogenations) Preparation of Unsaturated hydrocarbons: (Alkenes and Alkynes): (Elimination reactions (Dehydration, dehydrohalogenation and dehalogenation reactions) and Saytzeff rule) Reactions of Unsaturated hydrocarbons: (Electrophilic addition reactions (Markovnikov's rule), hydrogenation, halogenation, hydrohalogenation, and hydration - Oxidation reactions Acidity of alkynes). Syllabus 4

Aromatic compounds Aromaticity: structure and bonding requirements and Hckel's rule - Nomenclature of aromatic compounds - Electrophilic aromatic substitution reactions: (Alkylation, acylation, halogenations, nitration and sulfonation) - Effects of substituents on electrophilic aromatic substitution reactions - Side-chain reactions: (Oxidation of alkylbenzenes). Lectures (2) Alcohols, Phenols and Ethers Structure, classifications and nomenclature - Physical properties - Preparation of alcohols and phenols: (Hydration of alkenes - Nucleophilic substitution reaction of alkyl halides - Reduction of aldehydes, ketones and acids - Addition of Grignard compounds to aldehydes and ketones) - Preparation of Phenols: (Benzene sulfonic acids) - Preparation of ethers (Williamson synthesis) - Reactions of Alcohols, Phenols and Ethers: (Salt formation of alcohols and phenols (Acidity of phenols and Reaction of Alcohols with Sodium metal) - Reactions of Alcohols and Ethers with Hydrogen halides - Conversion of Alcohols to alkyl halides Oxidation of alcohols - Electrophilic substitution reactions of phenols) - Alcohols with More Than One Hydroxyl Group; glycols. Lectures (4) 1st Midterm Exam Syllabus 5 Aldehydes and Ketones Structure and Nomenclature - Physical properties - Preparation of aldehydes and ketones: (Hydration of alkynes - Ozonolysis of alkynes - Friedel-Crafts acylation - Oxidation of alcohols) - Reactions of aldehydes and ketones: (Nucleophilic addition reaction (addition of hydrogen cyanide, Reduction, Grignard addition, addition of Alcohol (hemiacetal and acetal

Formation), addition of ammonia and amine derivatives). Lectures (3) Carbohydrates Definitions and Classification (monosaccharides, disaccharides and polysaccharides) Monosaccharides: (Nomenclature - Structure (Optical isomerism, cyclic structure, Fischer Projection, Haworth Formulas)) - Reactions of Monosaccharides: (Reduction and oxidation of monosaccharides) Disaccharides: (Maltose, Cellobiose, Sucrose and Lactose) Polysaccharides: (Cellulose and Starch) Lectures (4) Carboxylic acids and Their Derivatives Structure and Nomenclature - Physical properties - Acidity of Carboxylic acids - Preparation: (Hydrolysis of nitrile - Carbonation of Grignard reagents) - Reactions of carboxylic acids: (Salt Formation - Ester, amide, anhydride, and acid chloride formation). Syllabus 6 Amines Structure of amines - Nomenclature of amines - Physical properties of amines - Basicity of amines - Preparation of amines: (Reduction of nitro compounds, nitriles and amides Alkylation of ammonia) - Reactions of amines: (Sulfa drugs - Diazonium salts (Formation and Replacement reactions) Lectures (2) 2nd Midterm Exam. Amino Acids, Peptides, and Proteins - Sources, classification and Structure - The acidbase Properties of Amino Acids - Reactions of amino acids: (The Ninhydrin Reaction, Peptides - Sanger reaction - Formation of an amide

linkage (The peptide bond: Proteins)) - Structure of proteins. Lectures (4) Nucleic Acids Chemical Structure: (General structure (Nucleoside, Nucleotide and Nucleic acids) - DNA; structure - RNA; structure and types). Lectures (2) Final Exam. References 7 Organic chemistry: A short course by I Harold Hart, David J. Hart and Leslie E. Craine, Houghton Mifflin Company, USA, 2012. Elements of Organic Chemistry (second edition) is written by Isaak Zimmerman and Henry Zimmerman and published by Macmillan Publishing Co., Inc. New York in 1983. : - - : - : /.. - | 8 9 Bonding, Structural Formulas, and Molecular Shapes Organic Chemistry: Definition 10

o The word Organic can be a biological or chemical term, in biology it means anything that is living or has lived. The opposite is NonOrganic. o Organic Chemistry is unique in that it deals with vast numbers of substances, both natural and synthetic. The clothes, the petroleum products, the paper, rubber, wood, plastics, paint, cosmetics, insecticides, and drugs o But, from the chemical makeup of organic compounds, it was recognized that one constituent common to all was the element carbon. o Organic chemistry is defined as the study of carbon/hydrogencontaining compounds and their derivatives. The Uniqueness of Carbon 11 o What is unique about the element carbon? o Why does it form so many compounds? The answers lie in structure of the carbon atom. The The position of carbon in the periodic table. o These factors enable it to form strong bonds with other carbon atoms and with other elements (hydrogen, oxygen, nitrogen, halogens,etc). o Each organic compound has its own characteristic set of physical and chemical properties which depend on the structure of the molecule.

Atomic Structure 12 o Atoms consist of three main particles: neutrons (have no charge), protons (positively charged) and electrons (negatively charged). Neutrons and protons are found in the nucleus. Electrons are found outside the nucleus. Electrons are distributed around the nucleus in successive shells (principal energy levels). o Atom is electrically neutral. i.e. Number of electrons = Number of protons o Atomic number of an element is the number of protons. Atomic Structure 13 o The energy levels are designated by capital letters (K, L, M, N, ..) or whole numbers (n). o The maximum capacity of a shell = 2n2 electrons. n = number of the energy level. o For example, the element carbon (atomic number 6) 6 electrons are distributed about the nucleus as Shell K L M N Number of electrons 2 4 0 0

Atomic Structure 14 Valance Electrons: ElectronDot Structures o Valance Electrons are those electrons located in the outermost energy level (the valance shell). o Electron-dot structures The symbol of the element represents the core of the atom. The valance electrons are shown as dots around the symbol. Chemical Bonding 15 o In 1916 G.N. Lewis pointed out that: The noble gases were stable elements and he ascribed their lack of reactivity to their having their valence shells filled with electrons. 2 electrons in case of helium. 8 electrons for the other noble gases. o According to Lewis, in interacting with one another atoms can achieve a greater degree stability by ofrearrangement of the valence

electrons to acquire the outer-shell structure of the closest noble gas in the periodic table. Chemical Bonding 16 17 Chemical Bonding A) Ionic Bonds o Elements at the left of the periodic table give up their valance electrons and become +ve charged ions (cations). o Elements at the right of the periodic table gain the electrons and become -ve charged ions (anions). o Ionic bond The electrostatic force of attraction between oppositely charged ions. A + B A + B x x Electron donor Electron acceptor atom

atom A + x B Electrostatic attraction Cation A Anion x B Ionic bond o The majority of ionic compounds are inorganic substances. Chemical Bonding 18 Electronegativity Measures The Ability of An Atom To Attract Electrons

Increasing electronegativity H F O N 4 3.5 3 B Be Li 2.5 2 1.5 1 S Si Al

3 2.5 2.1 1.8 1.5 Mg Na 1.2 0.9 Br K 2.8 0.8 Li + F electron transfer + Li+

F Li+ He configuration F Ne configuration Cl Li P C Decreasing electronegativity 2.1 +

F Li+ F ionic bond 19 B) Covalent Bonds Chemical Bonding o Elements that are close to each other in the periodic table attain the stable noble gas configuration by sharing valence electrons between them. o Covalent bond The chemical bond formed when two atoms share one pair of electrons. o A shared electron pair between two atoms or single covalent bond, will be represented by a dash (-). 20 Chemical Bonding B) Covalent

Bonds Example s H2 H H H + H Cl + Cl2 Cl H C H H or H H Cl Cl

C lone pair or Cl or N Cl H H methyl amime N H H H N C

C O H H H ethanol H lone pair H C Cl H N N2 N H

H methane H C H each H shares two electrons (He configuration) H H H or H chloromethane lone pair 21 Chemical Bonding

B) Covalent Bonds o In molecules that consist of two like atoms; the bonding electrons are shared equally (both the same o When two unlikeatoms atoms; have electronegativity). the bonding electrons are no longer shared equally (shared unequally). A Polar Covalent Bond A bond, in which an electron pair is unequally.atom assumes a partial negative charge The more shared electronegative and the less electronegative atom assumes a partial positive charge. C

O or C O 22 Chemical Bonding B) Coordinate Covalent Bonds o There are molecules in which one atom supplies both electrons to another atom in the formation of a covalent bond. o For example; H H N H H Ammonia (Lewis base) + H Hydrogen ion

(Lewis acid) H N H H Ammonium ion o Lewis base The species that furnishes the electron pair to form a coordinate covalent bond. o Lewis acid The species that accepts the electron pair to complete its Chemical Bonding 23 How Many Bonds to an Atom? Number Covalence The number of covalent bonds that an atom can form with other atoms. i.e. the covalence number is equal to the number of electrons needed to fill its valance shell. Element Number of Number of electrons valence electrons in filled valence shell H

1 2 1 C 4 8 4 N 5 8 3 O 6 8 2 F, Cl, Br, I 7 8 1 Covalence number 24 Atomic Orbitals Shapes of Organic Molecule

Orbital Picture of Covalent Bo o An atomic orbital represents a specific region in space in which an electron is most likely to be found. o Atomic orbitals are designated in the order in which they are filled by the letters s, p, d, and f. o Example s: K shell has only one 1s orbital. L shell has one 2s and three 2p (2px, 2py and 2pz). z o An s orbital is spherically shaped electron cloud with the atoms nucleus and its center. x y s orbital 25 Shapes of Organic Molecule Orbital Picture of Covalent Bo Atomic Orbitals

o A p orbital is a dumbbell-shaped electron cloud with the nucleus between the two lobes. o Each p orbital is oriented along one of three perpendicular coordinate axes (in the x, y, or z direction). y y px orbital x x x y z z z py orbital pz orbital 26

Shapes of Organic Molecule Orbital Picture of Covalent Bo Atomic Orbitals o An energy level diagram of atomic orbitals. 3px 3py 3pz 3s Energy content of orbital increases 2px 2py 2pz 2s 1s

o When filling the atomic orbitals, keep in mind that (1) An atomic orbital contain no more 2 electrons. (2) Electrons fill orbitals of lower energy first. is filled by 2 electrons until all the orbitals of equal (3) No orbital energy have at least one electron. 27 Shapes of Organic Molecule Orbital Picture of Covalent Bo Atomic Orbitals o The electronic configuration of carbon (atomic number 6) can be represented 2as 2 1 1s 2s sp x2p1y or 1s22s22p2 2px Energy content of atomic orbital 2py 2pz

2s 1s Energy level diagram for carbon. 28 Shapes of Organic Molecule Orbital Picture of Covalent Bo Molecular Orbitals o A covalent bond consists of the overlap between two atomic orbitals to form a molecular orbital. o Example: Molecular orbital of H2 H + H H Two 1s atomic orbitals H

Overlap H H One bonding sigma molecular orbital 29 Molecular Orbitals Shapes of Organic Molecule Orbital Picture of Covalent Bo o Sigma bonds ( bonds) can be formed from The overlap of two s atomic orbitals. The end-on overlap of two p atomic orbitals. The overlap of two an s atomic orbital with a p atomic orbital. o pi bonds ( bonds) can be formed from the side-side overlap between two p atomic orbitals. Bond Energy and Bond Length 30

o A molecule is more stable than the isolated constituent atoms. This stability is apparent in the release of energy during the formation of the molecular bond. o Heat of formation (bond energy) The amount of energy released when a bond is formed. o Bond dissociation energy The amount of energy that must be absorbed to break a bond. o Bond length The distance between molecular structure. nuclei in the Hybridization (Alkanes sp3) 31 o In the case of alkanes sp3, the three 2p orbitals of the carbon atom are combined with its 2s orbital to form four new orbitals called "sp3" hybrid orbitals. o Four hybrid orbitals were required since there are four atoms attached to the central carbon atom. o These new orbitals will have an energy slightly

above the 2s orbital and below the 2p orbitals as shown in the following illustration. o Notice that no change occurred with the 1s orbital. o Regular tetrahedron angles of 109.5. with all H-C-H bond Metha ne Hybridization (Alkanes sp3) 32 Hybridization (Alkanes sp2) 33 o In the case of alkenes sp2, the 2s orbital is combined with only two of the 2p orbitals (since we only need three hybrid orbitals for the three groups. thinking of groups as atoms and non-bonding pairs) forming three hybrid orbitals called sp2 hybrid orbitals.

o The other p-orbital remains unhybridized and is at right angles to the trigonal planar arrangement of the hybrid orbitals. o The trigonal planar arrangement has bond angles of 120. Ethene (Ethyl ene) Hybridization (Alkanes sp) 34 o In the case of alkynes sp, the 2s orbital is combined with only one of the 2p orbitals to yield two sp hybrid orbitals. o The two hybrid orbitals will be arranged as far apart as possible from each other with the result being a linear arrangement. o The two unhybridized p-orbitals stay in their respective positions (at right angles to each other) and perpendicular to the linear molecule (180). Ethyne (Acetyle ne) Inductive Effect 35

o Inductive effect can be defined as the permanent displacement of electrons forming a covalent bond (sigma bonds ) towards the more electronegative element or group. o The inductive effect is represented by the symbol, the arrow pointing towards the more electronegative element or group of elements. (+ I) effect if the substituent electron-donating O (- I) effect if the substituent electron-withdrawing R Electron-donating substituents (+I): -CH3, -C2H5, -NH2, OH, OCH3, . Bond Polarity and Dipole Moment () 36 o Dipole moment (depends on the inductive effect). o A bond with the electrons shared equally between two atoms is called a nonpolar bond like in Cl-Cl and C-C bond in ethane. o A bond with the electrons shared unequally between two different elements is called a polar bond. o The bond polarity is measured by its dipole moment (). o Dipole moment () defined to be the amount of charge separation ( + and ) multiplied by the bond length. Functional Groups 37

Functional Group is a reactive portion of an organic molecule, an atom, or a group of atoms that confers on the whole molecule its characteristic properties. Class Alkane Alkene Alkyne Alkyl halide Alcohol Ether Aldehyde Ketone Carboxylic acid Ester Amine General

formula RH Functional group C C (single Specific H3C CH3 H2C = CH2 R CH = CH2 bond) C = C (double C C R C CH bond) (triple HC CH -bond) RX X (X = F, Cl, Br, H3C - Cl -I) R OH OH H3C - OH R O R - C- O C H3C O CH3 R

O C O C R O C R O C OR R H R R NH2 OH O C C

H O C C O C OH O C OR C NH2 O H C H, H3C O C O H3C C H CH3 O O ,

H C OH H3C C OH O H C OCH3 O H3C C OCH3 H3C NH2

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