KSIU Organic Chemistry Lecture 1 PDF

CHE 111: Organic Chemistry For level 1 : Biotechnology and Petroleum Programs (Faculty of Basic Sciences) Level 1 Lecture 1: Introduction: Atomic structure and chemical bonds A.Prof/ Omayma Fawzy Abdel Gawad Date : 07/ 10 / 2024 Welcome to the Chemistry world ...

CHE 111: Organic Chemistry For level 1 : Biotechnology and Petroleum Programs (Faculty of Basic Sciences) Level 1 Lecture 1: Introduction: Atomic structure and chemical bonds A.Prof/ Omayma Fawzy Abdel Gawad Date : 07/ 10 / 2024 Welcome to the Chemistry world  1. Atomic structure and chemical bonds Nucleus - Electrons - electron orbitals Chemical bonds What is chemical bone? Types of chemical bonds Electronegativity Formal charge 3 Atomic structure Refers to the composition and arrangement of atoms, which are the basic building blocks of matter 1. Atoms The smallest unit of an element, consisting of a nucleus surrounded by electrons. 2. Nucleus Located at the center of the atom, it contains: Protons: Positively charged particles. Neutrons: Neutral particles (no charge). 3. Electrons Negatively charged particles that orbit the nucleus in various energy levels or shells. 4. Atomic Number The number of protons in the nucleus, which determines the element's identity (e.g., hydrogen has an atomic number of 1). 5. Mass Number The total number of protons and neutrons in the nucleus. It gives an approximation of the atom's mass. 8. Ions Atoms that have gained or lost electrons, resulting in a positive (cation) or negative (anion) charge. 6. Isotopes Variants of the same element that have the same number of protons but different numbers of neutrons (e.g., carbon-12 and carbon-14). Isotopes are atoms of the same element with the same atomic number (number of protons) but different mass numbers (total number of protons and neutrons). Examples: 1. Carbon Isotopes: Carbon-12 (\text{^{12}C}): 6 protons, 6 neutrons; stable. Carbon-14 (\text{^{14}C}): 6 protons, 8 neutrons; radioactive, used in radiocarbon applications. 2. Hydrogen Isotopes: Protium (H): 1 proton, 0 neutrons; stable. Deuterium (D or \text{^{2}H}): 1 proton, 1 neutron; stable. Tritium (T or \text{^{3}H}): 1 proton, 2 neutrons; radioactive. 7. Electron Configuration The distribution of electrons among the various energy levels and orbitals influences an atom's chemical properties. Electron orbitals: The electrons are distributed around the nucleus in orbits, and these orbitals are sometimes called shells or energy levels.  Electrons occupy the lowest energy orbitals available before filling higher energy ones. Order of Filling: Orbitals are filled in the following order based on increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. Valence Electrons The electrons in the outermost shell are involved in chemical bonding. Significance: The distribution of valence electrons is crucial for understanding an element's reactivity and bonding behavior.  Chemical bonds Attraction force between atoms, molecules, or ions that enables the formation of chemical compounds Types of chemical bonds: 1. Ionic bond 2. Covalent bond 3. Metallic bond 4. Hydrogen bond 5. Coordinate bond 11 1. Ionic Bonds Formed when one atom donates an electron to another, resulting in the creation of positively and negatively charged ions. (formed when an atom (usually metals) lose one or more electrons forming a positive ion and another atom (usually non-metals) gaining one or more electrons forming a negative ion) Characteristics: 1. Typically occur between metals and nonmetals. 2. The electrostatic attraction between oppositely charged ions holds them together. Examples: The octet rule The octet rule is a fundamental concept in chemistry that explains how atoms tend to bond by achieving a stable electron configuration, typically resembling that of the nearest noble gas. The octet rule: In forming compounds, they gain, lose, or share electrons to achieve a stable electron configuration characterized by complete valence electrons. 2. Covalent bond A type of chemical bond where two atoms share one or more pairs of electrons, forming molecules. In a covalent bond, each atom contributes at least one electron to form a shared pair. This sharing allows each atom to attain a full valence shell, which increases stability. Types of Covalent Bonds: Single Bond: Involves one pair of shared electrons (e.g., H₂O). Double Bond: Involves two pairs of shared electrons (e.g., CO₂). Triple Bond: Involves three pairs of shared electrons (e.g., N₂).  Notes  Ionic bonds are very common in inorganic compounds, but rare in organic ones.  The ionization energy of carbon is too large and the electron affinity too small for carbon to realistically form a C4+ or C4− ion.  What kinds of bonds that link carbon to other elements in millions of organic compounds?????  Lewis structures – Lewis dot formulas - Lewis dot structures - electron dot structures - Lewis electron dot structures (LEDs) These diagrams show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. Is a diagrammatic representation of the valence electrons in a molecule. It illustrates how atoms are bonded together and show any lone pairs of electrons. How to draw Lewis formula or Lewis structure? Step 1: Find the Total Number of Valence Electrons. Step 2: Find the Number of Electrons Needed to Make the Atoms "Happy“. Step 3: Determine the Number of Bonds in the Molecule. Step 4: Choose a Central Atom. Step 5: Draw a Skeletal Structure. Step 6: Place Electrons Around Outside Atoms. Examples 1. H2O 2. SO3 Electronegativity Is a measure of an atom's ability to attract electrons when it forms a chemical bond. It plays a crucial role in determining the nature of bonds between atoms and influences molecular properties. Bond polarity The distribution of electrical charge across a bond between two atoms. It arises from differences in electronegativity between the bonded atoms, influencing the molecule's overall polarity and its chemical properties.  According to bond polarity covalent bond classified into: Nonpolar Covalent Bonds Polar Covalent Bonds Occur between atoms with similar or identical Form between atoms with a moderate difference in electronegativities (difference of 0 to 0.4). electronegativities (difference of 0.4 to 1.7). Example: The bond in molecular nitrogen (N₂) or Example: In water (H₂O), oxygen is more oxygen (O₂) is nonpolar since both atoms attract electronegative than hydrogen, leading to a partial electrons equally. negative charge (δ-) on oxygen and a partial positive charge (δ+) on hydrogen, creating a polar molecule. Dipole moment Is a vector quantity that measures the separation of positive and negative charges in a molecule. It is an important concept in chemistry and physics, particularly in understanding molecular polarity, intermolecular forces, and the behavior of molecules in electric fields. The dipole moment (𝜇) is defined as the product of the charge (𝑞) and the distance (r) between the centers of positive and negative charge: 𝜇=𝑞×𝑟 Examples: Hydrogen Chloride (HCl): - HCl has a polar covalent bond with H (electronegativity = 2.1) and Cl (electronegativity = 3.0). - The dipole moment points toward chlorine, indicating that it has a partial negative charge.3 NH3 is not a planar molecule while BF3 is a planar molecule Due to the presence of a lone pair of electron on nitrogen atom of NH3 its structure is non- planar. Hence it possesses a net dipole moment. Assignment 1: Why CO2 has net dipole moment zero while SO2 has net dipole moment ? 3. Metallic Bond Is a type of chemical bonding that occurs between metal atoms. It involves the sharing of free electrons among a lattice of metal cations, leading to a unique set of properties characteristic of metals. Electron Sea Model: In metallic bonding, metal atoms release some of their electrons, which then become delocalized and free to move throughout the metal lattice. This "sea of electrons" allows for the conduction of electricity and heat, as these electrons can flow freely. Cation Formation: Metal atoms lose their valence electrons to form positively charged ions (cations). The remaining metal cations are held together by the electrostatic attraction to the surrounding electron sea. Propertise of metallic bonds Electrical Conductivity: Metals conduct electricity due to the mobility of the delocalized electrons. Thermal Conductivity: The free electrons also facilitate the transfer of heat. Malleability and Ductility: Metals can be hammered or stretched into thin sheets (malleability) and drawn into wires (ductility) without breaking. This is because the layers of atoms can slide over one another while maintaining the metallic bond through the electron sea. Luster: Metals have a shiny appearance due to the ability of delocalized electrons to absorb and re-emit light. 4. hydrogen bond Is a type of attractive interaction that occurs between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom. Hydrogen bonding is crucial in many biological and chemical processes. Formation: A hydrogen bond forms when a hydrogen atom covalently bonded to a highly electronegative atom (such as oxygen, nitrogen, or fluorine) experiences an attraction to another electronegative atom. This interaction occurs due to the partial positive charge on the hydrogen atom and the partial negative charge on the electronegative atom. Types of hydrogen bond 1. Intramolecular Hydrogen Bonds Definition: Occur within a single molecule, where a hydrogen atom forms a bond with an electronegative atom in the same molecule. Example: In certain organic compounds like salicylic acid, the –OH group can form a hydrogen bond with a carbonyl (C=O) group within the same molecule, influencing its structure and properties. 2. Intermolecular Hydrogen Bonds Definition: Occur between different molecules. This type is more common and plays a critical role in determining the properties of substances. Types: Strong Intermolecular Hydrogen Bonds: Typically occur between very electronegative atoms (like O, N, and F) and can significantly influence boiling and melting points. Example: Water (H₂O) exhibits strong intermolecular hydrogen bonds, contributing to its high boiling point. Weak Intermolecular Hydrogen Bonds: Form between less electronegative atoms or in less polar environments. Example: Hydrogen bonds in hydrocarbons or certain organic compounds that contain polar functional groups. 5. Coordinate bond (also known as a dative bond) Is a type of covalent bond where both electrons in the bond originate from the same atom. This bond forms when a lone pair of electrons from one atom is shared with another atom that has an empty orbital. Formation: A coordinate bond is formed when one atom donates a pair of electrons to another atom that is electron-deficient (lacking a complete valence shell). The atom donating the electrons is called the donor, while the atom accepting the electrons is referred to as the acceptor. Metallic complexes  Formal charge It is a concept used in chemistry to determine electrical charge distribution in a molecule or ion. It helps in understanding the stability and reactivity of chemical species. It is the hypothetical charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. Formal Charge (FC)=Valence Electrons − Non-bonding Electrons− ½ × Bonding Electrons Where: Valence Electrons: The number of electrons in the outer shell of the atom (from the periodic table). Non-bonding Electrons: The number of lone pair electrons on the atom. Bonding Electrons: The total number of electrons in bonds (counted for both atoms involved in the bond). Steps to Calculate Formal Charge 1. Draw the Lewis Structure: Start by drawing the Lewis structure of the molecule or ion. 2. Count Valence Electrons: Determine the number of valence electrons for each atom. 3. Identify Non-bonding and Bonding Electrons: Count the lone pair (non-bonding) electrons and half the bonding electrons for each atom. 4. Apply the Formula: Use the formula to calculate the formal charge for each atom. Example 1: Formal Charge in Water (H₂O) 1. Lewis Structure: Oxygen has 6 valence electrons, and each hydrogen has 1. Water has two H-O bonds and two lone pairs on oxygen. 2. Count Electrons: Oxygen: Valence Electrons = 6 Non-bonding Electrons = 4 (two lone pairs) Bonding Electrons = 4 (two bonds with hydrogen) Hydrogen: Valence Electrons = 1 Non-bonding Electrons = 0 Bonding Electrons = 2 (one bond with oxygen) 3. Calculate Formal Charges: Oxygen: 𝐹𝐶= 6 − 4 − ½ × 4 = 6 − 4 − 2 = 0 Each Hydrogen: 𝐹𝐶= 1 − 0 − ½ × 2 = 1 − 0 − 1 = 0  O=C=O Example 2: Carbon Dioxide (CO₂) 1. Lewis Structure: Carbon (C) has 4 valence electrons. Each Oxygen (O) has 6 valence electrons. The structure consists of one carbon atom double-bonded to two oxygen atoms. 2. Count Valence Electrons: Total for CO₂: 4 + 6 + 6 = 16 valence electrons. 3. Identify Non-bonding and Bonding Electrons: For Carbon: Valence Electrons = 4 Non-bonding Electrons = 0 Bonding Electrons = 8 (4 from each double bond) For Each Oxygen: Valence Electrons = 6 Non-bonding Electrons = 4 (2 lone pairs) Bonding Electrons = 4 (2 from the double bond) 3. Calculate Formal Charges: Carbon: 𝐹𝐶= 4 − 0 − ½ × 8 = 4 − 0 − 4 = 0 Each Oxygen: 𝐹𝐶= 6 − 4 − ½ × 4 = 6 − 4 − 2 = 0 Assignment 2: Calculate the formal charge for: Ammonium Ion (NH₄⁺)  Importance of formal charge 1. Assessing Stability of Molecules The distribution of formal charges can indicate the stability of a molecule. Generally, structures with formal charges close to zero are more stable. 2. Identifying Resonance Structures Formal charge helps identify and evaluate resonance structures, which are alternative Lewis structures that contribute to the overall electronic structure of a molecule. 3. Predicting Reactivity and Functional Groups Formal charge can help predict where a molecule is likely to react, as regions with higher formal charges tend to be more electrophilic or nucleophilic. 4. Determining the Best Lewis Structure When multiple Lewis structures can be drawn, formal charge helps identify the most favorable one by minimizing formal charges.