Lecture-1-pharm D clinical-Introduction.pdf

Pharmaceutical Organic Chemistry I Level 1 Pharm D clinical 2024-2025 Lecture (1) Introduction 1 Course Description: Pharmaceutical Organic Chemistry I Code: PC102C Overall Aims of Course By the end of this course the students wil...

Pharmaceutical Organic Chemistry I Level 1 Pharm D clinical 2024-2025 Lecture (1) Introduction 1 Course Description: Pharmaceutical Organic Chemistry I Code: PC102C Overall Aims of Course By the end of this course the students will be able to demonstrate knowledge of the basic concepts of organic chemistry effectively, and to carry simple organic experiments safely and under the supervision of their instructors. 2 Syllabus Introduction Hybridization Resonance Alkanes Alkenes Alkynes Alcohol and Ethers Alkyl halides Aromaticity 3 Course Staff: Dr. Sandra Nabil Email: [email protected] Dr. May Adel Email: [email protected] 4 Building On: Basic Knowledge of atoms, electrons, charges … just the basics. 6 Reaching To: Atom structure. Electrons distribution (Aufbau, Pauli Exclusion and Hund’s rules). Chemical Bonds. Atomic orbitals & Molecular orbitals. 7 An element’s properties depend on the structure of its atoms. Each element consists of unique atoms. An atom is the smallest unit of matter that still retains the properties of an element. 8  Atoms are composed of subatomic particles. Neutrons (no electrical charge) Protons (positive charge) Neutrons and protons form the atomic nucleus. Electrons (negative charge) Electrons form a cloud around the nucleus. Neutron mass and proton mass are almost identical. Electrons has negligible mass 9 10 Atomic Number and Atomic Mass Atomic number (Z) is the number of protons which is equal to the number of electrons. It determines the atom’s identity Mass number is the sum of protons plus neutrons in the nucleus. Atomic mass (A), the atom’s total mass, can be approximated by the mass number. 11 To summarize: 12 The Energy Levels of Electrons 13 Electronic Structure of Atoms: Electrons are concentrated about the nucleus in regions of 3D space called principle energy levels and identified by the principle quantum numbers (n) with values 1,2,3,4,5,6 & 7. Each principle energy level (shell), can contain up to 2n2 electrons. 1st shell can contain 2n2 (2 x 1= 2) electrons 2nd shell can accommodate up to eight electrons {2n2 (2 x 2= 8)} 3rd shell 18 electrons 4th shell 32 electrons and so on. 14 Each principle energy shell is subdivided into regions of space called orbitals and each energy shell consists of a specific number of orbitals: The first principal energy level contains only a single orbital called 1s orbital.(2 electrons) The second principle level contains: ✓ one s orbital (2s) ✓ three p orbitals (2px, 2py, 2pz) These orbitals are designated (2s, 2px, 2py, 2pz). (8 electrons) The third principle level contains: ✓ one s orbital (3s) ✓three p orbitals (3px, 3py, 3pz) ✓ five d orbitals (3dx, 3dy, 3dz, 3dxy, 3dzy) These orbitals are designated (3s, 3px, 3py, 3pz, 3dx, 3dy, 3dz, 3dxy, 3dzy). & so on. 16 17 Electron Distribution and Chemical Properties The chemical behavior of an atom is determined by the distribution of electrons in electron shells. Valence electrons are those in the outermost shell, or valence shell. The chemical behavior of an atom is mostly determined by the valence electrons. Elements with a full valence shell are chemically inert. 18 19  Neon, with two filled shells (10 electrons) (a) Electron-distribution diagram First shell Second shell 20  Neon, with two filled shells (10 electrons) (a) Electron-distribution diagram First shell Second shell (b) Separate electron orbitals 2 electrons 1s orbital 21  Neon, with two filled shells (10 electrons) (a) Electron-distribution diagram First shell Second shell (b) Separate electron orbitals x y 2 electrons 2 electrons z 6 electrons 1s orbital 2s orbital Three 2p orbitals 22  Neon, with two filled shells (10 electrons) (a) Electron-distribution diagram First shell Second shell (b) Separate electron orbitals x y 2 2 6 z 1s orbital 2s orbital Three 2p orbitals (c) Superimposed electron orbitals 1s, 2s, and 2p orbitals 23  In writing electronic configurations, we follow the Aufbau principle, Hund’s rule, and the Pauli exclusion principle. The Aufbau principle (German for building up) orbitals fill order of increasing energy from lowest to highest energy. In other words, electrons always go into orbitals with the lowest possible energy. Orbitals which are of exactly the same energy, such as the 2pX, 2pY, and 2pZ orbitals, are said to be degenerate. 25  Hund’s rule says that when electrons go into degenerate orbitals, they occupy them singly before pairing begins. Finally, the Pauli exclusion principle states that only electrons with opposite spins can occupy the same orbital. In other words, if two electrons must go into the same orbital, they must be paired. 26 Given a periodic table, all we need to know to write the electronic configuration for a given atom is the atomic number Z, which tells us the number of electrons in the neutral atom. 27 Ionic, Covalent, coordinate and Polar Bonds G. N. Lewis proposed that an atom is most stable if its outer shell is either filled or contains eight electrons. According to Lewis’s theory, an atom will give up, accept, or share electrons in order to achieve a filled outer shell or an outer shell that contains eight electrons. This theory has come to be called the octet rule. 28  Electrons in inner shells are called core electrons & they do not participate in chemical bonding. Electrons in the outermost shell are called valence electrons. Helium, neon & argon belong to noble gases. They have an extremely stable “closed-shell” electron configuration and are very unreactive. 29 1- Ionic Bonds Ionic compounds are formed when an atom on the left side of the periodic table (an electropositive element) transfers one or more electrons to an atom on the right side of the periodic table (an electronegative element). 30  Sodium has one valence electron while chlorine has seven valence electrons, when sodium metal and chlorine gas are mixed, each sodium atom transfers an electron to a chlorine atom, and crystalline sodium chloride (table salt) is formed as a result. The positively charged sodium ions and negatively charged chloride ions held together by the attraction of opposite charges (electrostatic attractions) to form an ionic bond. 31 2- Covalent Bonds Instead of giving up or acquiring electrons, an atom can achieve a filled outer shell by sharing electrons. For example, two fluorine atoms can each attain a filled shell of eight electrons by sharing their unpaired valence electrons. A bond formed as a result of sharing electrons is called a covalent bond. 32 More examples Bond-pair electrons Lone-pair electrons 33 H Cl  Types of covalent bonds: Acc. to bond polarity: 1) Polar covalent bond. 2) Non-polar covalent bond. Acc. to orbital overlap: 1) Sigma (σ) bond. 2) Pi (π) bond. 34 Electronegativity: The tendency of an atom to draw electrons in a covalent bond toward itself is called electronegativity. An electronegative atom attracts electrons; an electropositive atom donates electrons. EN increases across a row in the periodic table and decreases in going down a column. The most EN atom is fluorine. 35 36  Polar covalent bonds In F-F covalent bond, the atoms that share the bonding electrons are identical so they share the electrons equally. Such a bond is called a nonpolar covalent bond. In contrast, the bonding electrons in hydrogen chloride, water, and ammonia are more attracted to one atom than another because the atoms that share the electrons in these molecules are different and have different ENs. This results in a polar distribution of charge and a polar covalent bond. 37  Indicate the bond type in the following molecules: NaF Ionic bond HCl Polar covalent bond Cl2 Non polar covalent bond CH4 Non polar covalent bond 38  Understanding bond polarity is critical to understanding how organic reactions occur, because a central rule that governs the reactivity of organic compounds is that electron-rich atoms or molecules are attracted to electron-deficient atoms or molecules. 39 3- Coordinate (Dative covalent) bonds A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. When a molecule having atom surrounded by only 6 electrons and a molecule rich by electrons through non bonding (unshared pairs) electrons reacts together, a coordinate bond will occur by complete sharing of 2 electrons from the donor atom while the electron poor atom will act as acceptor for these 2 electrons 40 H+ NH3 NH4+ 41  Bond & Lone pairs Bond pair electrons are those shared in a bond. Lone pair electrons are those not used in bonding (nonbonding electrons). Bond-pair electrons When you draw a Lewis structure, make sure that: Hydrogen atoms are surrounded by no more than two electrons C, O, N, and halogen (F, Cl, Br, I) atoms are surrounded by no more than eight electrons—they must obey the octet rule. (Bond pair + Lone pair = 8 except for hydrogen). 44  H Cl Unshared or lone pair (LP) shared or bond pair 45  Draw Lewis structures, showing all valence electrons, for the following covalent molecules: (a) CO2 (b) NH3 (c) CH3Cl H a) O C O b) H N H c) H C Cl H H 47 Chemical Bonding Theories: Lewis Theory: Lewis proposed that an atom is most stable if its outer shell is either filled or contains eight electrons. According to Lewis’s theory, an atom will give up, accept, or share electrons in order to achieve a filled outer shell or an outer shell that contains eight electrons. This theory has come to be called the octet rule. Such electron movement/sharing results in a bond & a new molecule is formed. 53 Molecular Orbital theory: The Lewis model, tells us only part of the story. A drawback of the model is that it treats electrons like particles and does not take into account their wavelike properties. According to MO theory, covalent bonds result from the combination of atomic orbitals (AO) to form molecular orbitals (MO)— orbitals that belong to the whole molecule rather than to a single atom. Like atomic orbitals, molecular orbitals have specific sizes, shapes, and energies (differ from those of AO) and describes the volume of space around the nucleus of an atom where an electron is likely to be found 55 56 57 Types of bonds: Sigma bond: Overlap of two s orbitals or end-on overlap of two p orbitals. There is more but when we take the next part. 58 End-on overlap of two p orbitals to form a σ bonding molecular orbital and 59 a σ* antibonding molecular orbital. Pi bond: Side-to-side overlap of two p orbitals. Side-to-side overlap of two parallel p orbitals to form a  bonding 60 molecular orbital and a * antibonding molecular orbital.  Sigma Pi Sigma 61 62