Lecture 1 Pharmaceutical Organic Chemistry PO101 PDF

Principals of Organic Chemistry ‫مبادئ الكيمياء العضوية‬  Chapter 1B Carbon Compounds and Chemical Bonds chapter 1 2  Introduction of Organic Chemistry The chemistry of the compounds of carbon. The human body is largely composed of organi...

Principals of Organic Chemistry ‫مبادئ الكيمياء العضوية‬  Chapter 1B Carbon Compounds and Chemical Bonds chapter 1 2  Introduction of Organic Chemistry The chemistry of the compounds of carbon. The human body is largely composed of organic  compounds. Organic chemistry plays a central role in medicine,  bioengineering etc. Common Elements in Organic Compounds  The structure of atom Simplified structure of an atom structure of carbon atom  electron shells a) Atomic number = number of Electrons b) Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells. c) Electron shells determine how an atom behaves when it encounters other atoms  Electrons are placed in shells according to rules: 1) The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons C would like to Gain 4 electrons N would like to O would like to Gain 3 electrons Gain 2 electrons  Chemical Bonds: The Octet Rule Octet Rule Atoms form bonds to produce the electron configuration of a noble gas (because the electronic configuration of noble gases is particularly stable). For most atoms of interest this means achieving a valence shell configuration of 8 electrons corresponding to that of the nearest noble gas. Atoms close to helium achieve a valence shell configuration of 2 electrons. Atoms can form either ionic or covalent bonds to satisfy the octet rule. A chemical bond is a force of attraction that holds two atoms together to fill its electron shells with 8 electrons (octet rule). Kinds of chemical bonds: 1. Ionic bonds 2. Covalent bonds 3. Metallic bonds chapter 1 10  IONIC BOND bond formed between two ions by the transfer of electrons 11  Formation of Ions from Metals  Ionic compounds result when metals react with nonmetals  Metals lose electrons to match the number of valence electrons of their nearest noble gas  Positive ions form when the number of electrons are less than the number of protons Group 1 metals  ion 1+ Group 2 metals  ion 2+ Group 13 metals  ion 3+ 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.  Learning Check  A. X would be the electron dot formula for 1) Na 2) K 3) Al  B. X would be the electron dot formula  1) B 2) N 3) P  Theories of Covalent Bonding 11.1 Valence bond (VB) theory and orbital hybridization 11.2 The mode of orbital overlap and types of covalent bonds 11.3 Molecular orbital (MO) theory and electron delocalization  The Covalent Bond A covalent bond is a chemical bond formed between (nonmetal + nonmetal) or (metalloide + nonmetal) in which two or more electrons are shared by two atoms. chapter 1 16  The Covalent Bond Covalent bonds occur between atoms of similar electronegativity (close to each other in the periodic table) Atoms achieve octets by sharing of valence electrons Molecules result from this covalent bonding Valence electrons can be indicated by dots (electron-dot formula or Lewis structures) but t time-consuming The usual way to indicate the two electrons in a bond is to use a line (one line = two electrons) chapter 1 17  NONPOLAR COVALENT BONDS when electrons are shared equally H2 or Cl2  The Covalent Bond Oxygen Atom Oxygen Atom Oxygen Molecule (O2) chapter 1 19  The Covalent Bond when two atoms share a pair of electrons. P+1 P+1 chapter 1 20  The Covalent Bond when two atoms share a pair of electrons. P+1 P+1 It’s like both atoms have a filled orbital. chapter 1 21  The Covalent Bond The sharing of a pair of electrons between 2 atoms. (or even 2 or 3 pairs of electrons). H2  The Covalent Bond The sharing of a pair of electrons between 2 atoms. Cl2 – Covalent Bonds – Non polar covalent Bond è Atoms achieve octets by sharing of valence electrons è Molecules result from this covalent bonding è Valence electrons can be indicated by dots (electron-dot formula or Lewis structures) but this is time-consuming è The usual way to indicate the two electrons in a bond is to use a line (one line = two electrons) Chapter 1 24 POLAR COVALENT BONDS when electrons are shared but shared unequally H2O  2-Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich electron poor region region e- poor e- rich H F H F d+ d- 26 9.5  3- The Metallic Bond Metallic Bond is a bond found in metals; holds metal atoms together very strongly. Formed between atoms of metallic elements. Electron clouds around atoms. Good conductors at all states, very high melting points Examples; Na, Fe, Al, Au, Co, Cu  Orbital: a region in space where the probability of finding an electron is large Atomic Orbitals (AOs): The region of space where one or two electrons of an isolated atom are likely to be found. Molecular Orbitals (MOs) :The region of space where one or two electrons of a molecule are likely to be found. chapter 1 28  Linus Pauling Proposed that valence atomic orbitals in the molecule are different from those in the isolated atoms Mixing of certain combinations of atomic orbitals generates new atomic orbitals Process of orbital mixing = hybridization; generates hybrid orbitals  Hybrid Orbitals Key Points The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed. Types of Hybrid Orbitals sp sp2 sp3 sp3d sp3d2 – Bonding Molecular Orbitals (Ymolec) AOs combine by addition (the AOs of the same phase sign overlap) The value of Y2 (electron probability density) in the region between the two nuclei increases The two electrons between the nuclei serve to attract the nuclei towards each other This is the ground state (lowest energy state) of the MO chapter 1 31 Molecular Orbital of Hydrogen antibonding orbital bonding orbital chapter 1 32  The energy of electrons in the bonding orbitals is substantially less than the energy of electrons in the individual atoms – The energy of electrons in the antibonding orbitals is substantially more In the ground state of the hydrogen molecule electrons occupy the lower energy bonding orbital only chapter 1 33  Bonding Molecular Orbital The overlapping of two hydrogen 1s atomic orbitals 1s 1s Ø1 Ø2 Ψ+ = Ø1 + Ø2 atomic atomic bonding orbital orbital orbital The overlapping of two hydrogen 1s waves chapter 1 34 – Antibonding molecular orbital (Y *molec) Formed by interaction of AOs with opposite phase signs Electrons in the antibonding orbital avoid the region between the two nuclei Repulsive forces between the nuclei predominate and electrons in antibonding orbitals make nuclei fly apart chapter 1 35  Antibonding Molecular Orbital The overlapping of two hydrogen 1s atomic orbitals 1s 1s Ø1 Ø2 Ψ- = Ø1 - Ø2 atomic atomic antibonding orbital orbital orbital The overlapping of two hydrogen 1s waves node chapter 1 36 MO s* antibonding s bonding MO MO s* antibonding s bonding MO Formation of (sigma σ) bond H2 HCl Cl2 chapter 1 39 a) Sigma bond formation by s – s overlap s * antibonding MO Diagram of sigma bond 1s 1s formation by s – s overlap s bonding MO chapter 1 40 b) Sigma bond formation by s – p overlap s * antibonding MO p Diagram of sigma bond s formation by s – p overlap s bonding MO chapter 1 41 c) Sigma bond formation by p – p overlap Energy Diagram of sigma bond formation by p – p overlap chapter 1 42 Formation of (pi π) bond  * antibonding MO p p  bonding MO chapter 1 43 Structural Theory Central Premises Valency: atoms in organic compounds form a fixed number of bonds. Carbon can form one or more bonds to other carbons. 44  The Structure of Methane and Ethane: sp3 Hybridization The structure of methane with its four identical tetrahedral bonds cannot be adequately explained using the electronic configuration of carbon Hybridization of the valence orbitals (2s and 2p) provides four new identical orbitals which can be used for the bonding in methane Orbital hybridization is a mathematical combination of the 2s and 2p wave functions to obtain wave functions for the new orbitals chapter 1 45  When one 2s orbital and three 2p orbitals are hybridized four new and identical sp3 orbitals are obtained – When four orbitals are hybridized, four orbitals must result – Each new orbital has one part s character and 3 parts p character – The four identical orbitals are oriented in a tetrahedral arrangements – The antibonding orbitals are not derived in the following diagram The four sp3 orbitals are then combined with the 1s orbitals of four hydrogens to give the molecular orbitals of methane Each new molecular orbital can accommodate 2 electrons chapter 1 46 hybrid orbitals – sp, sp2, or sp3 formation of s bond  The sp3 hybrid orbitals in CH4 VSEPR predicts a tetrahedral shape Figure 11.4 Molecule geometry of CH4 4 sp3 hybrid orbitals 10.4 The three molecular shapes of the tetrahedral electron-group arrangement Examples: CH4, SiCl4, SO42-, ClO4- Examples: Examples: NH3 H2O PF3 OF2 ClO3 SCl2 H 3 O+ Figure 10.8 – Ethane (C2H6) The carbon-carbon bond is made from overlap of two sp3 orbitals to form a s bond The molecule is approximately tetrahedral around each carbon chapter 1 53  The representations of ethane show the tetrahedral arrangement around each carbon – a. calculated electron density surface b. ball-and-stick model c. typical 3-dimensional drawing Generally there is relatively free rotation about s bonds – Very little energy (13-26 kcal/mol) is required to rotate around the carbon- carbon bond of ethane chapter 1 54 Examples of Sigma Bond Formation 55  The Structure of Ethene (Ethylene) : sp2 Hybridization Ethene (C2H2) contains a carbon-carbon double bond and is in the class of organic compounds called alkenes – Another example of the alkenes is propene The geometry around each carbon is called trigonal planar – All atoms directly connected to each carbon are in a plane – The bonds point towards the corners of a regular triangle – The bond angle are approximately 120o chapter 1 56  The sp2 hybrid orbitals in BF3 VSEPR predicts a trigonal planar shape Figure 11.3  There are three s bonds around each carbon of ethene and these are formed by using sp2 hybridized orbitals The three sp2 hybridized orbitals come from mixing one s and two p orbitals – One p orbital is left unhybridized The sp2 orbitals are arranged in a trigonal planar arrangement – The p orbital is perpendicular (orthoganol) to the plane chapter 1 58  Overlap of sp2 orbitals in ethylene results in formation of a s framework – One sp2 orbital on each carbon overlaps to form a carbon-carbon s bond; the remaining sp2 orbitals form bonds to hydrogen The leftover p orbitals on each carbon overlap to form a bonding  bond between the two carbons A  bond results from overlap of p orbitals above and below the plane of the s bond – It has a nodal plane passing through the two bonded nuclei and between the two lobes of the  molecular orbital chapter 1 59  The bonding  orbital results from overlap of p orbital lobes of the same sign The antibonding * orbital results from overlap of p orbital lobes of opposite sign – The antibonding orbital has one node connecting the two nuclei and another node between the two carbons The bonding  orbital is lower in energy than the antibonding orbital – In the ground state two spin paired electrons are in the bonding orbital – The antibonding *orbital can be occupied if an electron becomes promoted from a lower level ( e.g. by absorption of light) chapter 1 61  The s orbital is lower in energy than the  orbital – The ground state electronic configuration of ethene is shown chapter 1 62  s bond remaining p orbitals from sp or sp2  bond hinders rotation Planar molecule (each carbon about the carbon-to- is trigonal planar) with  cloud carbon bond above and below the plane formation of  bond  Restricted Rotation and the Double Bond There is a large energy barrier to rotation (about 264 kJ/mol) around the double bond – This corresponds to the strength of a  bond – The rotational barrier of a carbon-carbon single bond is 13-26 kJ/mol This rotational barrier results because the p orbitals must be well aligned for maximum overlap and formation of the  bond Rotation of the p orbitals 90o totally breaks the  bond chapter 1 67  The Structure of Ethyne (Acetylene): sp Hybridization Ethyne (acetylene) is a member of a group of compounds called alkynes which all have carbon-carbon triple bonds – Propyne is another typical alkyne The arrangement of atoms around each carbon is linear with bond angles 180o chapter 1 68  The sp hybrid orbitals in gaseous BeCl2 atomic orbitals hybrid VSEPR orbitals predicts a linear shape Figure 11.2 orbital box diagrams  The carbon in ethyne is sp hybridized – One s and one p orbital are mixed to form two sp orbitals – Two p orbitals are left unhybridized The two sp orbitals are oriented 180o relative to each other around the carbon nucleus – The two p orbitals are perpendicular to the axis that passes through the center of the sp orbitals chapter 1 70  In ethyne the sp orbitals on the two carbons overlap to form a s bond – The remaining sp orbitals overlap with hydrogen 1s orbitals The p orbitals on each carbon overlap to form two  bonds The triple bond consists of one s and two  bonds chapter 1 71  Depictions of ethyne show that the electron density around the carbon-carbon bond has circular symmetry – Even if rotation around the carbon-carbon bond occurred, a different compound would not result chapter 1 72 – Bond Lengths of Ethyne, Ethene and Ethane The carbon-carbon bond length is shorter as more bonds hold the carbons together – With more electron density between the carbons, there is more “glue” to hold the nuclei of the carbons together The carbon-hydrogen bond lengths also get shorter with more s character of the bond – 2s orbitals are held more closely to the nucleus than 2p orbitals – A hybridized orbital with more percent s character is held more closely to the nucleus than an orbital with less s character – The sp orbital of ethyne has 50% s character and its C-H bond is shorter – The sp3 orbital of ethane has only 25% s character and its C-H bond is longer chapter 1 73  Summary of Concepts from Quantum Mechanics – Atomic Orbital(AO): region in space around a nucleus where there is a high probability of finding an electron – Molecular Orbital (MO): results from overlap of atomic orbitals – Bonding Orbitals: when AOs of same sign overlap – Antibonding Orbitals: when AOs of opposite sign overlap – The energy of electrons in a bonding orbital is less than the energy of the individual atoms – The energy of electrons in an antibonding orbitals is more chapter 1 74 – The number of molecular orbitals formed equals the number of the atomic orbitals used – Hybridized orbitals are obtained by mixing the wave functions of different types of orbitals Four sp3 orbitals are obtained from mixing one s and three p orbitals – The geometry of the four orbitals is tetrahedral – This is the hybridization used in the carbon of methane Three sp2 orbitals are obtained from mixing one s and two p orbitals – The geometry of the three orbitals is trigonal planar – The left over p orbital is used to make a  bond – This is the hybridization used in the carbons of ethene Two sp orbitals are obtained from mixing one s and one p orbital – The geometry of the two orbitals is linear – The two leftover p orbitals are used to make two  bonds – This is the hybridization used in the carbons of ethyne Sigma (s) bonds have circular symmetry when viewed along the bond axis Pi () bonds result from sideways overlap of two p orbitals chapter 1 75 Hybridization – is the mixing of two or more atomic orbitals in an atom (usually central atom) to form a new set of hybrid orbitals (same in shape and energy).  Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have different shape and energy from original atomic orbitals.  Number of hybrid orbitals is equal to the number of pure atomic orbitals used in the hybridization process.  Covalent bonds are formed by: a. Overlap of hybrid orbitals with atomic orbitals b. Overlap of hybrid orbitals with other hybrid orbitals  Hybridization of Carbon sp3 sp2 sp CH4 H2C = CH2 HC CH chapter 1 77 sp3 hybridization in Methane CH4 excitation becomes Ground state excited state Molecule geometry of CH4 4 sp3 hybrid orbitals  sp2 hybridization in Ethene C2H4 energy sp2 hybridization in Ethene C2H4 120 0 sp Hybridization in Acetylene C2H2 2p 2p hydridization sp 2s E 1s 1s 180 0 85 What is the hybridization in each carbon atom? 4 2 5 6 3 1 7 sp3 – 2, 5 sp2 – 1, 3, 4 sp – 6, 7 tetrahedral trigonal planar linear What is the hybridization in each carbon atom? 4 3 2 1 88 89 90 91 92 93 94 95 96 97 98 99 100 101  Which set of approximate bond angles at C1, C2, and N of the following molecule indicates the correct shape? a) C1 120°, C2 120°, N 120° b) C1 109.5°, C2 120°, N 109.5° c) C1 109.5°, C2 120°, N 120° d) C1 120°, C2 109.5°, N 109.5° 102  Which set of hybridization states of C1, C2, and N of the following molecule is correct? a) C1 sp2, C2 sp3, N sp3 b) C1 sp3, C2 sp2, N sp3 c) C1 sp3, C2 sp2, N sp2 d) C1 sp2, C2 sp2, N sp3 103  Which species does not contain an sp3-hybridized atom? a) BH3 b) BH4- c) NH3 d) NH4+ Which molecule contains an sp-hybridized atom? a) HCO2H b) HNO3 c) HNO2 d) HCN 104  Which of the following statements is wrong? a) When two orbitals overlap in-phase with each other, a bonding molecular orbital forms. b) When two orbitals overlap out-of-phase with each other, an antibonding molecular orbital forms. c) When one of two atoms connected by a σ bond rotates about the bond axis, orbital overlap is lost. d) When one of two atoms connected by a π bond rotates about the bond axis, orbital overlap is los 105  Correct answer: d) C1 120°, C2 109.5°, N 109.5° Correct answer: a) C1 sp2, C2 sp3, N sp3 Correct answer: a) BH3 Correct answer: d) HCN c) When one of two atoms connected by a σ bond rotates about the bond axis, orbital overlap is lost. 106