Intro to Oranic Chemistry PDF

 ORGANIC CHEMISTRY Asst. Prof. Dr. Süleyman Aşır Faculty of Engineering 16H113 (Innovation) [email protected] ORGANIC CHEMISTRY NEAR EAST UNIVERSITY  ORGANIC CHEMIS...

 ORGANIC CHEMISTRY Asst. Prof. Dr. Süleyman Aşır Faculty of Engineering 16H113 (Innovation) [email protected] ORGANIC CHEMISTRY NEAR EAST UNIVERSITY  ORGANIC CHEMISTRY Textbook: “Organic Chemistry”, Solomons Fryhle Snyder Chapter 1 3  Life and the Chemistry of Carbon Compounds—We are Stardust What Is the Origin of the Element Carbon? How Did Living Organisms Arise? Chapter 1 4  Development of the Science of Organic Chemistry According to vitalism, organic compounds were only those that came from living organisms, and only living things could synthesize organic compounds through intervention of a vital force. Friedrich Wöhler, however, discovered in 1828 that an organic compound called urea (a constituent of urine) could be made by evaporating an aqueous solution of the inorganic compound ammonium cyanate. Chapter 1 5 Development of the Science of Organic Chemistry The synthesis of an organic compound, began the evolution of organic chemistry as a scientific discipline. Chapter 1 6  The Chemistry of... Natural Products The word “organic” is still used today by some people to mean “coming from living organisms” In science today, the study of compounds from living organisms is called natural products chemistry. Chapter 1 7  Atomic Structure The compounds we encounter in chemistry are made up of elements combined in different proportions. Elements are made up of atoms. An atom consists of a dense, positively charged nucleus containing protons and neutrons and a surrounding cloud of electrons. Chapter 1 8 Atomic Structure Each proton of the nucleus bears one positive charge; electrons bear one negative charge. Neutrons are electrically neutral; they bear no charge. Chapter 1 9  Atomic Structure Each element is distinguished by its atomic number (Z), a number equal to the number of protons in its nucleus. Because an atom is electrically neutral, the atomic number also equals the number of electrons surrounding the nucleus. Chapter 1 10  Isotopes Although all the nuclei of all atoms of the same element will have the same number of protons, some atoms of the same element may have different masses because they have different numbers of neutrons. Such atoms are called isotopes. Chapter 1 11  Isotopes For example, the element carbon has six protons in its nucleus giving it an atomic number of 6. Most carbon atoms also have six neutrons in their nuclei, and because each proton and each neutron contributes one atomic mass unit (1 amu) to the mass of the atom, carbon atoms of this kind have a mass number of 12 and are written as 12C. Chapter 1 12  Practice Problem 1.1 There are two stable isotopes of nitrogen, 14N and 15N. How many protons and neutrons does each isotope have? Chapter 1 13  Valence Electrons The most important shell, called the valence shell, is the outermost shell because the electrons of this shell are the ones that an atom uses in making chemical bonds with other atoms to form compounds. Chapter 1 14  Valence Electrons How do we know how many electrons an atom has in its valence shell? We look at the periodic table. The number of electrons in the valence shell (called valence electrons) is equal to the group number of the atom. For example, carbon is in group IVA and carbon has four valence electrons; oxygen is in group VIA and oxygen has six valence electrons. The halogens of group VIIA all have seven electrons. Chapter 1 15  Practice Problem 1.2 How many valence electrons does each of the following atoms have? (a) Na (b) Cl (c) Si (d) B (e) Ne (f) N Chapter 1 16  CHEMICAL BONDS: THE OCTET RULE Two major types of chemical bonds were proposed: 1. Ionic (or electrovalent) bonds are formed by the transfer of one or more electrons from one atom to another to create ions. 2. Covalent bonds result when atoms share electrons. Chapter 1 17  THE OCTET RULE The tendency for an atom to achieve a configuration where its valence shell contains eight electrons is called the octet rule. Chapter 1 18  Ionic Bonds Atoms may gain or lose electrons and form charged particles called ions. Ions form because atoms can achieve the electronic configuration of a noble gas by gaining or losing electrons. An ionic bond is an attractive force between oppositely charged ions. Chapter 1 19  Electronegativity Electronegativity is a measure of the ability of an atom to attract electrons. Electronegativity increases as we go across a horizontal row of the periodic table from left to right and it increases as we go up a vertical column Chapter 1 20 Electronegativity Chapter 1 21  Ionic Bonds An example of the formation of an ionic bond is the reaction of lithium and fluorine atoms. Ionic compounds, often called salts, form only when atoms of very different electronegativities transfer electrons to become ions. Chapter 1 22  Ionic Substances Ionic substances, because of their strong internal electrostatic forces, are usually very high melting solids, often having melting points above 1000 oC. In polar solvents, such as water, the ions are solvated, and such solutions usually conduct an electric current. Chapter 1 23  Practice Problem 1.3 Using the periodic table, which element in each pair is more electronegative? (a) Si, O (b) N, C (c) Cl, Br (d) S, P Chapter 1 24 Covalent Bonds and Lewis Structures Covalent bonds form by sharing of electrons between atoms of similar electronegativities to achieve the configuration of a noble gas. Molecules are composed of atoms joined exclusively or predominantly by covalent bonds. Chapter 1 25  Lewis structures A dash structural formula has lines that show bonding electron pairs and includes elemental symbols for the atoms in a molecule. Chapter 1 26  Lewis structures Two carbon atoms can use one electron pair between them to form a carbon–carbon single bond Chapter 1 27  Lewis structures Atoms can share two or more pairs of electrons to form multiple covalent bonds. Chapter 1 28  Lewis structures Carbon atoms can also share more than one electron pair with another atom to form a multiple covalent bond. Consider the examples of a carbon–carbon double bond in ethene (ethylene) and a carbon–carbon triple bond in ethyne (acetylene). Chapter 1 29  Lewis structures Ions, themselves, may contain covalent bonds. Consider, as an example, the ammonium ion. Chapter 1 30  PRACTICE Problem 1.4 Consider the following compounds and decide whether the bond in them would be ionic or covalent. (a) KCl (b) F2 (c) PH3 (d) CBr4 Chapter 1 31  How To Write Lewis Structures 1. Lewis structures show the connections between atoms in a molecule or ion using only the valence electrons of the atoms involved. 2. For main group elements, the number of valence electrons a neutral atom brings to a Lewis structure is the same as its group number in the periodic table. Chapter 1 32  How To Write Lewis Structures 3. If the structure we are drawing is a negative ion (an anion), we add one electron for each negative charge to the original count of valence electrons. If the structure is a positive ion (a cation), we subtract one electron for each positive charge. 4. In drawing Lewis structures we try to give each atom the electron configuration of a noble gas. Chapter 1 33  Problem 1.1 Write the Lewis structure of CH3F. Chapter 1 34  Problem 1.2 Write a Lewis structure for methylamine (CH3NH2). Chapter 1 35  Practice Problem 1.5 Write the Lewis structure of (a) CH2F2 (difluoromethane) and (b) CHCl3 (chloroform). Chapter 1 36  Practice Problem 1.6 Write the Lewis structure of CH3OH. Chapter 1 37  How To Write Lewis Structures 5. If necessary, we use multiple bonds to satisfy the octet rule (i.e., give atoms the noble gas configuration). The carbonate ion (CO32-) illustrates this: Chapter 1 38  Problem 1.3 Write the Lewis structure of CH2O (formaldehyde). Chapter 1 39  Exceptions to the Octet Rule In second row elements fewer electrons are possible In higher rows other orbitalsare accessible and more than 8 electrons around an atom are possible Chapter 1 40  Formal Charges and How To Calculate Them F = Z - (1/2)S - U F: Formal charge Z: number of valence electrons S: number of shared electrons U: number of unshared electrons Chapter 1 41 Formal Charges Chapter 1 42 Formal Charges Chapter 1 43  Practice Problem 1.10 Write a Lewis structure for each of the following negative ions, and assign the formal negative charge to the correct atom: (a) CH3O- (b) NH2- (c) CN- (d) HCO2- (e) HCO3- (f) HC2- Chapter 1 44 A Summary of Formal Charges Chapter 1 45  Practice Problem 1.11 Assign the proper formal charge to the colored atom in each of the following structures: Chapter 1 46  ISOMERS Different compounds that have the same molecular formula Constitutional isomers are different compounds that have the same molecularf ormula but differ in the sequence in which their atoms are bonded—that is, their connectivity. Constitutional isomers usually have different physical properties (e.g., melting point, boiling point, and density) and different chemical properties (reactivity). Chapter 1 47 ISOMERS C3H6O Chapter 1 48  Problem 1.6 There are two constitutional isomers with the formula C2H6O. Write structural formulas for these isomers. Chapter 1 49 STRUCTURAL FORMULAS Chapter 1 50  Dash Structural Formulas Atoms joined by single bonds can rotate relatively freely with respect to one another. H H HH H H H H H HO H C C H H C O C C H C O C C H C H H H H HH H H H Equivalent dash formulas for propyl alcohol Chapter 1 51  Dash Structural Formulas Dash structural formulas show what is called the connectivity of the atoms. Constitutional isomers have different connectivities and, therefore, must have different structural formulas. Chapter 1 52 Condensed Structural Formulas Chapter 1 53  Problem 1.8 Write a condensed structural formula for the compound that follows: Chapter 1 54 Bond-Line Formulas Chapter 1 55 Chapter 1 56  Problem 1.9 Write the bond-line formula for Chapter 1 57  Practice Problem 1.14 Write each of the following condensed structural formulas as a bond-line formula: Chapter 1 58  Three-Dimensional Formulas  A dashed wedge ( ) represents a bond that projects behind the plane of the paper  A solid wedge ( ) represents a bond that projects out of the plane of the paper  An ordinary line ( ) represents a bond that lies in the plane of the paper Chapter 1 59  H H H H C H H OR C C etc. H C H H H H H Ethane H H Br C OR C OR C etc. H H H Br Br H H H H Bromomethane Chapter 1 60  Br H H OH Cl H Examples of bond-line formulas that include three-dimensional representations H NH 2 Br H HO An example involving An example involving trigonal planar geometry linear geometry Chapter 1 61  Resonance Theory In a Lewis structure, we draw a well-defined location for the electrons in a molecule. In many molecules and ions (especially those containing p bonds), more than one equivalent Lewis structure can be drawn which represent the same molecule Chapter 1 62  We can write three different but equivalent structures, 1–3 O O O C C C O O O O O O 1 2 3 63 Chapter 1  O O O becomes becomes C C C O O O O O O 1 2 3  Structures 1–3, although not identical on paper, are equivalent; all of its carbon– oxygen bonds are of equal length 64 Chapter 1  Resonance theory states that whenever a molecule or ion can be represented by two or more Lewis structures that differ only in the positions of the electrons, two things will be true: None of these structures, which we call resonance structures or resonance contributors, will be a realistic representation for the molecule or ion. None will be in complete accord with the physical or chemical properties of the substance The actual molecule or ion will be better represented by a hybrid (average) of these structures 65 Chapter 1  Resonance structures, then, are not real structures for the actual molecule or ion; they exist only on paper O O O C C C O O O O O O 66 Chapter 1  It is also important to distinguish between resonance and an equilibrium In an equilibrium between two or more species, it is quite correct to think of different structures and moving (or fluctuating) atoms, but not in the case of resonance (as in the carbonate ion). Here the atoms do not move, and the “structures” exist only on paper. An equilibrium is indicated by and resonance by 67 Chapter 1  How to Write Resonance Structures  Resonance structures exist only on paper. Although they have no real existence of their own, resonance structures are useful because they allow us to describe molecules and ions for which a single Lewis structure is inadequate 68 Chapter 1  We write two or more Lewis structures, calling them resonance structures or resonance contributors. We connect these structures by double-headed arrows , and we say that the real molecule or ion is a hybrid of all of them 69 Chapter 1  We are only allowed to move electrons in writing resonance structures These are resonance structures H2C CH2 CH CH2 This is not a proper resonance structure of 1 and 2 because a hydrogen atom has been moved 70 Chapter 1  All of the structures must be proper Lewis structures H H C O H H This is not a proper resonance structure of methanol (hypervalent carbon having 5 bonds!!) 71 Chapter 1  The energy of the resonance hybrid is lower than the energy of any contributing structure. Resonance stabilizes a molecule or ion. This is especially true when the resonance structures are equivalent. Chemists call this stabilization resonance stabilization. If the resonance structures are equivalent, then the resonance stabilization is large 72 Chapter 1  The more stable a structure is (when taken by itself), the greater is its contribution to the hybrid 73 Chapter 1  The Use of Curved Arrows: How to Write Resonance Structures  Curved arrows show the direction of electron flow in a reaction mechanism point from the source of an electron pair to the atom receiving the pair always show the flow of electrons from a site of higher electron density to a site of lower electron density never show the movement of atoms. Atoms are assumed to follow the flow of the electrons 74 Chapter 1  Examples HO H NOT HO H NOT O O N N H H H C H C H  H  O O O H3C O H+ OH H3C O +H H 75 Chapter 1  Examples (1) Benzene (2) Carboxylate ion (RCOO-) O R O (3) Ozone (O3) O O O 76 Chapter 1  How to Decide When One Resonance Structure Contributes More to the Hybrid  The more covalent bonds a structure has, the more stable it is  Charge separation decreases stability Resonance structures for formaldehyde Four O O Three covalent covalent C C bonds H H H H bonds more stable less stable 77 Chapter 1  Structures in which all the atoms have a complete valence shell of electrons (i.e., the noble gas structure) are more stable 78 Chapter 1 Atomic Orbitals and Electron Configuration 79 Chapter 1  Electron Configurations  The relative energies of atomic orbitals in the 1st & 2nd principal shells are as follows: Electrons in 1s orbitals have the lowest energy because they are closest to the positive nucleus Electrons in 2s orbitals are next lowest in energy Electrons of the three 2p orbitals have equal but higher energy than the 2s orbital Orbitals of equal energy (such as the three 2p orbitals) are called degenerate orbitals 80 Chapter 1  Aufbau principle Orbitals are filled so that those of lowest energy are filled first  Pauli exclusion principle A maximum of two electrons may be placed in each orbital but only when the spins of the electrons are paired 81 Chapter 1  Hund’s rule When we come to orbitals of equal energy (degenerate orbitals) such as the three p orbitals, we add one electron to each with their spins unpaired until each of the degenerate orbitals contains one electron. (This allows the electrons, which repel each other, to be farther apart.) Then we begin adding a second electron to each degenerate orbital so that the spins are paired 82 Chapter 1  83 Chapter 1 The Structure of Methane and Ethane: sp3 Hybridization Ground state of a carbon atom C 1s 2s 2px 2py 2pz 84 Chapter 1  Hybridization sp3 sp3 covalent hybridized bond H carbon H C H H (line bond structure) 85 Chapter 1  Hybridization sp3 H H C H H (Lewis structure) 109o H Tetrahedral structure C H H H Carbon with (3-D stucture) 4 bonds 86 Chapter 1  The Structure of Methane Chapter 1 87 The Structure of Ethane Chapter 1 88 Chapter 1 89  The Structure of Ethene (Ethylene): sp2 Hybridization  sp2 1  bond + 1 p bond o ~120 H H ~120o C C sp2 hybridized carbon H H o ~120 90 Chapter 1  sp2 3-Dimensional View p-bond (p-orbitals overlap) H H C C H H Planar structure Carbon with (3 + 1p) bonds 91 Chapter 1  92 Chapter 1  93 Chapter 1 An sp2-hybridized carbon atom Chapter 1 94  Restricted Rotation and the Double Bond  There is a large energy barrier to rotation associated with groups joined by a double bond ~264 kJ/mol (strength of the p bond) To compare: rotation of groups joined by C-C single bonds ~13-26 kJ/mol 95 Chapter 1  Cis -Trans Isomerism  Stereochemistry of double bonds H R H identical to H R H H H H H i H3C identical to Pr H H i Pr CH3 96 Chapter 1 H3C different from CH3 CH3 (trans) CH3 (cis) Restricted rotation of C=C 97 Chapter 1  – Cis-Trans System Useful for 1,2 disubstituted alkenes e.g. (1) H Br Cl Cl Br vs. H H H trans -1-Bromo- cis-1-Bromo- 2-chloroethane 2-chloroethane 98 Chapter 1 (2) H vs. H H H trans -3-Hexene cis -3-Hexene (3) Br Br Br vs. Br trans -1,3- cis -1,3- Dibromopropene Dibromopropene 99 Chapter 1  The Structure of Ethyne (Acetylene): sp Hybridization  sp 1  bond + 2 p bond H C C H 180o sp2 hybridized carbon Linear structure Carbon with (2 + 2 p) bonds 100 Chapter 1  101 Chapter 1  102 Chapter 1  103 Chapter 1  104 Chapter 1  Hybridizing three p orbitals with one s orbital yields four sp3 orbitals and they are tetrahedral Hybridizing two p orbitals with one s orbital yields three sp2 orbitals and they are trigonal planar Hybridizing one p orbital with one s orbital yields two sp orbitals, a linear molecule Chapter 1 105  A sigma () bond (a type of single bond) is one in which the electron density has circular symmetry when viewed along the bond axis A pi (p) bond, part of double and triple carbon–carbon bonds, is one in which the electron densities of two adjacent parallel p orbitals overlap sideways to form a bonding pi molecular orbital Chapter 1 106 16. How to Predict Molecular Geometry: The Valence Shell Electron Pair Repulsion Model  Valence shell electron pair repulsion (VSEPR) model: 1) We consider molecules (or ions) in which the central atom is covalently bonded to two or more atoms or groups Chapter 1 107 2) We consider all of the valence electron pairs of the central atom—both those that are shared in covalent bonds, called bonding pairs, and those that are unshared, called nonbonding pairs or unshared pairs or lone pairs 108 Chapter 1 3) Because electron pairs repel each other, the electron pairs of the valence shell tend to stay as far apart as possible. The repulsion between nonbonding pairs is generally greater than that between bonding pairs 109 Chapter 1 4) We arrive at the geometry of the molecule by considering all of the electron pairs, bonding and nonbonding, but we describe the shape of the molecule or ion by referring to the positions of the nuclei (or atoms) and not by the positions of the electron pairs Chapter 1 110 Methane 111 Chapter 1  Ammonia  A tetrahedral arrangement of the electron pairs explains the trigonal pyramidal arrangement of the four atoms. The bond angles are 107° (not 109.5°) because the nonbonding pair occupies more space than the bonding pairs 112 Chapter 1 Water 113 Chapter 1 Boron Trifluoride Chapter 1 114 Beryllium Hydride Chapter 1 115 Carbon Dioxide Chapter 1 116 Chapter 1  END OF CHAPTER 1  118 Chapter 1