Pharm. Organic Chemistry 1 Lecture 1 _2024 Final PDF

# Pharm. Organic Chemistry 1 ## New Mansoura University - Faculty of Pharmacy - Pharm D Program - Semester 1 - PMC 102 - Pharm D Program ## Lecture 1: Introduction - Prof. / Shahenda Metwally El-Messery # What is Organic Chemistry? Organic chemistry started as the chemistry of life. Then it be...

# Pharm. Organic Chemistry 1 ## New Mansoura University - Faculty of Pharmacy - Pharm D Program - Semester 1 - PMC 102 - Pharm D Program ## Lecture 1: Introduction - Prof. / Shahenda Metwally El-Messery # What is Organic Chemistry? Organic chemistry started as the chemistry of life. Then it became the chemistry of carbon compounds. Now it is both. It is the chemistry of the compounds of carbon along with other elements such as are found in: - Living things (made of organic chemicals). - Proteins - DNA (controls genetic make-up). - Foods, medicines. # Some organic Chemicals - DNA - An image is shown of a DNA molecule. - Medicines - active Pharmaceutical Ingredients - Excipients - Fuels - Pigments - An image is shown of a blue powder. # Section 1 - Atoms, Electrons and Orbitals # Section 2 - Structure Determines Properties # Section 3 - Chemical Bonding # Section 4 - Polar Covalent Bonds, Electronegativity and Hybridization # Atoms, Electrons, and Orbitals - **Particles and Symbols of the Atom** - **Neutron:** Mass = 1.67 * 10^-27 kg, Charge = 0.00 C - **Electron:** Mass = 9.11 * 10^-31 kg, Charge = -1.60 *10^-19 C - **Proton:** Mass = 1.67 * 10^-27 kg, Charge = +1.60 * 10^-10 C > The number of protons in the nucleus is called the atomic number (Z). - **Atomic Structure: Orbitals** > Shapes of Atomic Orbitals for Electrons - There are four different kinds of orbitals for electrons based on those derived for a hydrogen atom. - **Orbitals and Shells** - The first shell contains one s orbital, denoted 1s, holds only two electrons. - The second shell contains one s orbital (2s) and three p orbitals (2p), eight electrons. - The third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons. - **p-Orbitals** - In each shell there are three perpendicular p orbitals, px, py, and pz, of equal energy. - Lobes of a p orbital are separated by region of zero electron density, a node. - **Atomic Structure: Electron Configurations** - **Ground state electron configuration:** Z electrons (Z = atomic number of the atom) are placed serial into the orbitals according to the following guidelines. - **Aufbau principle:** electrons go into lowest energy orbitals first. - **Pauli principle:** each orbital only contains a maximum of two electrons with anti-parallel spin. - **Hund’s rule:** When there are orbitals of equal energy in subshells to fill, one electron is added to each degenerated orbital before a second orbital is added to one of them. - **Electron Configuration** - With electron configuration elements are represented numerically by the number of electrons in their shells and number of shells. - For example, Nitrogen - configuration = 2,5 - **Electron configuration and Orbital diagram** > Give the electronic configuration for the following: - Na+ - Cl- - Mg+2 - Na11, Mg12, Cl17 - e- + Cl - Cl- - 1s2 2s2 2p6 3s2 3p6 # Structure Determines Properties - An image is shown of a chemical molecule and a space shuttle. # Electronegativity and Chemical Bonding - **Electronegativity** > Electronegativity is defined as the ability of an atom to attract electrons to itself. - **Electronegative atoms:** attract electrons strongly, hold their electrons tightly, and tend to take on electrons. - **Electropositive atoms:** attract electrons weakly and may give up electrons. > Electronegativity helps us predict the relative reactivity of analogous compounds. Recalling trends in electronegativity is extremely important! # Chemical Bonding - **A - Ionic Bonds** - **Ionic Bonding and Electron Transfer** > One can imagine ionic bonds as arising from the transfer of an electron from a metal to a nonmetal. > For example, gaseous sodium can transfer an electron to gaseous chlorine... > The transfer is endothermic but the formation of solid NaCl is strongly exothermic. - **Ionic Bonding in Organic Chemistry** > Carbon atoms rarely form ions; hence, ionic bonds are rare in organic chemistry. > Covalent bonds involving the sharing of electrons between nonmetal atoms are much more common. > Ions do appear in salts of C, N, O, and H which are nucleophilic (electron-donating) at the nonmetal atom. - **B - Covalent bonds, Lewis Formulas, and the Octet Rule** > Covalent bonds involve the sharing of electrons between two atoms. > Atoms share electrons to achieve a more stable electron configuration. > Maximum stability is reached when an atom achieves a full valence shell, isoelectronic with the nearest noble gas. > Electrons may not be shared equally - covalent bonds may be polarized! - **Covalent Bonding in H2** > A neutral hydrogen atom needs one more valence electron to achieve a full valence shell. > Two hydrogens sharing their valence electrons achieve a full n = 1 shell. > In H2, the electron configuration of each hydrogen atom is analogous to that of helium, the first noble gas. > Two dots or a line between two atoms denotes the sharing of two electrons between the atoms in a covalent bond. - **Covalent Bonding in F2** > A neutral fluorine atom needs one more valence electron to achieve a full valence shell. > Two hydrogens sharing their valence electrons achieve a full n = 2 shell containing 8 electrons. > In F2, the electron configuration of each hydrogen atom is analogous to that of neon, the second noble gas. > The six electrons on the periphery of each fluorine are nonbonding electron pairs or lone pairs. # Polar Covalent Bonds, Electronegativity and Hybridization - **Polar Covalent Bonds** > Two atoms of different electronegativities share electrons unequally in a covalent bond. The result is a polar covalent bond. > The greater the difference in electronegativity, the more polarized the bond. > Polar covalent bonds tend to be sites of reactivity in organic molecules. - **Electrostatic Potential Maps** > The model of molecular charge as a dipole is a simplification; in reality molecules contain a spatial distribution of charge. > An electrostatic potential (ESP) map shows the distribution of charge over a molecule. > Nonpolar molecules contain unpolarized bonds and/or a symmetric distribution of charge - **Periodic Trends in Electronegativity** > TABLE 1.3 Selected Values from the Pauling Electronegativity Scale | Period | Group Number 1A | Group Number 2A | Group Number 3A | Group Number 4A | Group Number 5A | Group Number 6A | Group Number 7A | | ---------- | ---------- | ---------- | ---------- | ---------- | ---------- | ---------- | ---------- | | 1 | H 2.1 | | | | | | | | 2 | Li 1.0 | Be 1.5 | B 2.0 | C 2.5 | N 3.0 | O 3.5 | F 4.0 | | 3 | Na 0.9 | Mg 1.2 | Al 1.5 | Si 1.8 | P 2.1 | S 2.5 | Cl 3.0 | | 4 | K 0.8 | Ca 1.0 | | | | | Br 2.8 | | 5 | | | | | | | I 2.5 | - **Theories of Covalent Bonding** > **1 - Valence and Hybridization** The number of atoms which are typically bonded to a given atom is termed the valence of that atom. | Bonds | Valence | | ----------- | ---------- | | One | Monovalent | | Four | Tetravalent | | Three | Trivalent | | Two | Divalent | | One | Monovalent | > **2 - Orbital Hybridization** > Mixing of two or more atomic orbitals to form a new set of hybrid orbitals. > Mix at least 2 non-equivalent atomic orbitals (e.g. s and p). > Hybrid orbitals have very different shape from original atomic orbitals. - **sp3 Hybridization** > **sp3 Orbitals and the Structure of Methane, CH4** - Carbon has four valence electrons (2s2 2p2). - In CH4, all C-H bonds are identical (tetrahedral). - sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3). - **VSEPR Model** > VSEPR model can be used to predict the bond angles in hybridized compounds. > Outer Shell of valence electrons surrounds an atom. > Valence electrons may form single, double, or triple bonds, or they may be unshared. > Negative charged region of space created and repulsion occurred so various regions of electron density around an atom spread out. - **NH3:** bond angle is 107.3° not 109.5° because The unshared pair of electrons on the nitrogen repels adjacent electron pairs more strongly than do bonding pairs of electrons. - **H2O:** bond angle is 104.5° not 109.5° because The 2 unshared pair of electrons on the oxygen. - **H2S:** bond angle is 92° not 109.5° because The 2 unshared pair of electrons on the sulfur. Bond angle is not like in H2O because of the difference in electronegativity between O and S atoms (O > S). - **sp2 Hybridization** > **sp2 Orbitals and The Structure of Ethylene, CH2=CH2** - sp² hybrid orbitals: 2s orbital combines with two 2p orbitals, giving three orbitals (spp = sp²). - sp² orbitals are in a plane with 120°angles, which makes them Trigonal Planar. - The remaining p orbital is perpendicular to the plane. - Two sp2 hybridized orbitals overlap to form a sigma (σ) bond. - p orbitals overlap side-to-side to formation a pi (π) bond. - sp2-sp2 σ bond and 2p-2p π-bond result in sharing four electrons and formation of C-C double bond. - **sp Hybridization** > **sp Orbitals and the Structure of Acetylene, CHECH** - C-C a triple bond sharing six electrons. - Carbon 2s orbital hybridizes with a single p orbital giving two sp hybrids. - two p orbitals remain unchanged. > The sp orbitals are linear, 180° apart on the x-axis. > Two p orbitals are perpendicular on the y-axis and the z-axis. > Sharing of six electrons forms C≡C > The two sp orbital form o bonds with hydrogens. - **Differences Between the Different Kinds of Hybridization** | Bond Type | Bond Angle | No. Of Remaining p-orbitals | No. of Bonds | s character | Geometry | Example | | ---------- | ---------- | ---------- | ---------- | ---------- | ---------- | ---------- | | sp3 | 109.5° | 0 | 4 sigma bonds ( σ ) | 25% (1/4) | Tetrahedral | CH4 | | sp3 | 107.3° | 0 | 3 sigma bonds ( σ ) | 25% (1/4) | Trigonal Pyramidal | NH3 | | sp3 | 104.5° | 0 | 2 sigma bonds ( σ ) | 25% (1/4) | Bent | H2O | | sp2 | 120°. | 1 | 3 sigma bonds ( σ ) + 1 pi bond ( π ) | 33% (1/3) | Trigonal Planar | C2H4 | | sp2 | | 1 | 3 sigma bonds ( σ ) + 1 pi bond ( π ) | 33% (1/3) | Trigonal Planar | H2C=O | | sp2 | | 1 | 3 sigma bonds ( σ ) + 1 pi bond ( π ) | 33% (1/3) | Trigonal Planar | R2C=O | | sp | 180°. | 2 | 3 sigma bonds ( σ ) + 2 pi bonds ( π ) | 50% (1/2) | Linear | HCECH | | sp | | 2 | 3 sigma bonds ( σ ) + 2 pi bonds ( π ) | 50% (1/2) | Linear | O=C=O (CO2) | | sp | | 2 | 3 sigma bonds ( σ ) + 2 pi bonds ( π ) | 50% (1/2) | Linear | HCEN | - **Sigma ( σ ) and Pi Bonds ( π )** > Single bond: 1 σ -bond > Double bond: 1 σ-bond and 1 π-bond > Triple bond: 1 σ-bond and 2 π-bonds > How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? > σ- bonds = 6 + 1 = 7, π-bond = 1 # Skills > How to - Connect structure to Property? - Define and understand electronegativity? - Detect differences between Chemical Bonding? - Assign different types of hybridization? - Calculate No of bonds?