Chemical Bonds PDF

Types of Chemical Bonds  Chemical bonds can be classified into three types, depending on the types of atoms involved in the bonding:  Ionic bond  Covalent bond Intramolecular Force  Metallic bond...

Types of Chemical Bonds  Chemical bonds can be classified into three types, depending on the types of atoms involved in the bonding:  Ionic bond  Covalent bond Intramolecular Force  Metallic bond 120 Types of Chemical Bonds: The Ionic Bond Ionic bond: results when electrons have been transferred between atoms, resulting in oppositely charged ions that attract each other.  Generally formed when metal atoms bond to nonmetal atoms.  Method: electron transfer. 121 Types of Chemical Bonds: The Covalent Bond Covalent bond: results when two atoms share some of their electrons:  Generally formed when nonmetal atoms bond together  Shared electrons hold the atoms together by attracting nuclei of both atoms.  Method: electron sharing  Multiple Covalent Bonds:  Single covalent bond: A covalent bond formed by sharing one electron pair (2eˉ). Represented by a single line: H−H  Double covalent bond: formed by sharing two electron pairs (4eˉ). Represented by a double line: O=O  Triple covalent bond: formed by sharing three electron pairs (6eˉ). Represented by a triple line: N≡N 122 Types of Chemical Bonds: The Coordinate Bond Coordinate 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.  Coordinate bond is a sharing of lone pair of electrons from one atom called donor (Lewis base) to another atom called acceptor (Lewis acid).  Lewis acid: electron pair acceptor e.g. H+, AlCl3, FeBr3, BF3.  Lewis base: electron pair donor e.g. compounds containing heteroatoms (O, S, N) e.g. NH3, H2O. 123 Representing Valence Electrons with Dots (Lewis Structures) Lewis Structures: simple diagrams to visualize the number of valence electrons in atoms of main-group elements by dots. The dots are placed around the element’s symbol with a maximum of two dots per side. Each dot represents one valence electron.  Remember: the number of valence electrons for main group element is equal to the group number of the element (except for helium, which is in group 8A but has only two valence electrons).  Note: While the exact location of dots is not critical, here we first place dots singly before pairing (except for helium which always has two paired dots) 124 Representing Valence Electrons with Dots (Lewis Structures)  The electron configuration of Oxygen is as follows:  Its Lewis structure is as follows: 125 Representing Valence Electrons with Dots (Lewis Structures)  Lewis structure for all period 2 elements: Practice: Draw the Lewis dot structure of a phosphorus atom. Solution: Since phosphorus is in Group 5A in the periodic table, it has 5 valence electrons. Represent these as five dots surrounding the symbol for phosphorus: 126 Lewis Structures: For Covalent Bonding Hydrogen and oxygen have the following Lewis structures: In water, hydrogen and oxygen share their valence electrons so that each hydrogen atom gets a duet and the oxygen atom gets an octet. 127 Lewis Theory Predicts That Hydrogen Should Exist as H2 The individual Hydrogen atoms has the following Lewis structure: When two hydrogen atoms share their valence electrons, they each get a duet, a stable configuration for hydrogen. Lewis theory predicts that elemental hydrogen exists as a diatomic molecule (H2). 128 Lewis Structures: Double and Triple Covalent Bonds Oxygen exists as a diatomic molecule (O2): Nitrogen exists as a diatomic molecule (N2): 129 Assessment 1. Write the Lewis structure for each atom or ion: a. Al b. sodium ion c. magnesium ion d. chloride ion 2. Use Lewis structures to explain why each element occurs as diatomic molecules: a. hydrogen b. bromine c. oxygen d. nitrogen 3. Write the Lewis structure for each compound: a. PH3 b. SCl2 c. HI d. CH4 e. NaF f. CaO g. SrBr2 h. K2O 4. Determine whether a bond between each pair of atoms would be nonpolar covalent, polar covalent, or ionic. a. Br & Br b. C & Cl c. Mg & I d. Sr & O 5. Order these compounds in order of increasing carbon–carbon bond strength and in order of decreasing carbon–carbon bond length: HC≡CH , H2C═CH2 , H3C─CH3 130  B- Molecules, Compounds and Chemical Bonds 4- Intermolecular Forces and Bond Polarity Intermolecular Forces Intermolecular forces: are the attractive forces that exist between molecules.  In contrast to intermolecular forces, intramolecular forces hold atoms together in a molecule.  Generally, intermolecular forces are much weaker than intramolecular forces. It usually requires much less energy to evaporate a liquid than to break the bonds in the molecules of the liquid. From the chemistry point of view:  The strength of the intermolecular forces (IMF) determines whether a compound has a high or low melting point and boiling point, and thus if the compound is a solid, liquid, or gas at a given temperature.  IMF influences the solubility of substances in various solvents.  IMF affects the rate and outcome of chemical reactions by influencing how reactant molecules come together and interact. 132 Intermolecular Forces in Covalent Molecules  There are three different types of intermolecular forces in covalent molecules, presented in order of increasing strength:  London dispersion forces (also called van der Waals forces)  Dipole–dipole interactions (also called van der Waals forces)  Hydrogen bonding Intermolecular Forces: London Dispersion Forces London dispersion forces: are very weak interactions due to the momentary changes in electron density in a molecule. Temporary dipole All covalent compounds exhibit London dispersion forces. These intermolecular forces are the only intermolecular forces present in nonpolar compounds. The strength of these forces is related to the size of the molecule.  The larger the molecule, the larger the attractive force between two molecules, and the stronger the intermolecular forces. 134 Intermolecular Forces: Dipole–dipole interactions Dipole–dipole interactions: are the attractive forces between the permanent dipoles of two polar molecules.  For example, the carbon–oxygen bond in formaldehyde, H2C=O, is polar because oxygen is more electronegative than carbon. This polar bond gives formaldehyde a permanent dipole, making it a polar molecule.  The dipoles in adjacent formaldehyde molecules can align so that the partial positive and partial negative charges are close to each other. These attractive forces due to permanent dipoles are much stronger than London dispersion forces. 135 Intermolecular Forces: Hydrogen Bonding Hydrogen bonding is a strong type of dipole-dipole interaction which occurs when a hydrogen atom bonded to O, N, or F, is electrostatically attracted to an O, N, or F atom in another molecule.  Hydrogen bonding is only possible between two molecules that contain a hydrogen atom bonded to a very electronegative atom—that is, oxygen, nitrogen, or fluorine  These forces are weaker than intramolecular bonds, but are much stronger than other intermolecular forces, causing these compounds to have high boiling points.  Hydrogen bonds are the strongest of the three types of intermolecular forces 136 Intermolecular Forces: Types of Hydrogen Bonding Intermolecular hydrogen bond: Refers to reaction between two same or different molecules.  Occurs when hydrogen locates between two electronegative groups (N, S, O). Intramolecular hydrogen bond: Refers to hydrogen bonds within the same molecule.  Stronger than intermolecular hydrogen bonds.  Higher boiling and melting points. Intramolecular hydrogen bonding of salicylaldehyde 137 Intermolecular Forces: Types of Hydrogen Bonding Examples of Intra- and Intermolecular Hydrogen Bonds Intermolecular Forces: Ion-dipole Forces Ion–dipole forces: are the electrostatic attractions between a charged ion and a dipole.  They are common in solutions and play an important role in the dissolution of ionic compounds, like NaCl, in polar solvent such as water.  The strength of ion–dipole interactions is directly proportional to: 1. the charge on the ion 2. the magnitude of the dipole of polar molecules. 139 Why are intermolecular forces important in pharmacy? IMF forces are of great importance in the field of pharmacy, as they play a crucial role in various aspects of pharmaceutical science and drug development. Here are some key ways in which IMF are important in pharmacy: 1- Drug solubility and formulation:  The solubility of a drug molecule in certain solvent is largely determined by the IMF between the drug and the solvent. Understanding these forces is crucial for designing effective drug formulations. 2- Drug absorption and bioavailability:  IMF influence the ability of a drug to be absorbed from the gastrointestinal tract into the bloodstream. For example, hydrogen bonding and ionic interactions can affect the permeability of drug molecules across biological membranes.  The bioavailability of a drug, which is the fraction of the administered dose that reaches the systemic circulation, is also affected by intermolecular forces during absorption, distribution, and metabolism. Why are intermolecular forces important in pharmacy? 3- Drug stability and shelf life: IMF such as hydrogen bonding and van der Waals interactions, can stabilize the structure of drug molecules and influence their resistance to degradation. Understanding the IMF involved in drug-excipient and drug-drug interactions is crucial for developing stable pharmaceutical formulations with an adequate shelf life. 4- Drug delivery and targeting: IMF play a role in the design of drug delivery systems, such as liposomes, nanoparticles, and polymeric carriers, which aim to improve the targeting and controlled release of drugs. 5- Protein-drug interactions: Many drugs exert their therapeutic effects by interacting with specific proteins, such as enzymes, receptors, or transport proteins. The strength and nature of these protein-drug interactions are determined by IMF. Intermolecular Forces: Problems 1- What types of intermolecular forces are present in each compound: (a) HCl; (b) C2H6 (ethane); (c) NH3? N.B. London dispersion forces are present in all covalent compounds. Dipole–dipole interactions are present only in polar compounds with a permanent dipole. Hydrogen bonding occurs only in compounds that contain an O – H, N – H, or H – F bond. a. HCl has London forces like all covalent compounds. HCl has a polar bond, so it exhibits dipole–dipole interactions. HCl has no H atom on an O, N, or F, so it has no intermolecular hydrogen bonding. b. C2H6 is a nonpolar molecule since it has only nonpolar C – C and C – H bonds. Thus, it exhibits only London forces. c. NH3 has London forces like all covalent compounds. NH3 has a net dipole from its three polar bonds, so it exhibits dipole–dipole interactions. NH3 has a H atom bonded to N, so it exhibits intermolecular hydrogen bonding.142 Intermolecular Forces: Problems 2. What types of intermolecular forces are present in each molecule? a. Cl2 b. HCN c. HF d. CH3Cl e. H2 3. Which of the compounds in each pair has stronger intermolecular forces? a. CO2 or H2O b. CO2 or HBr c. HBr or H2O d. CH4 or C2H6B 4. Determine which compound can form inter or intramolecular hydrogen bonding: a. 7.30 e e 143 Electronegativity and Bond Polarity  Electronegativity (EN): is the ability of an atom (in a molecule) to attract the bond electrons to itself.  is higher for nonmetals; and lower for metals  The greater the difference in electronegativity (ΔEN), the more polar the bond. δ+ and δ- in polar molecular compounds represent the partial positive and negative charges, to differentiate them from the full charge (+ or -) on ions in ionic compounds. 144 Electronegativity Values for Elements (Unitless) 145 Electronegativity and Bond Types 146 Electronegativity and Bond Types: Examples - Example: Based on the of values of electronegativity (EN) of elements, which bond is more polar: (B-Cl) or (C-Cl)? Answer: - The ΔEN of Cl and B = 3.0 - 2.0 = 1.0 - The ΔEN of Cl and C = 3.0 - 2.5 = 0.5  Hence, the B-Cl bond is more polar. - Exercise 1: Which of the following bonds is the most polar? (a) HF (b) SeF (c) NP (d) GaCl - Exercise 2: Predict the type of each bond (use the table of EN values): (a) HBr (b) OO (c) HO (d) SO 147