9.1 - Lewis Dot Symbols

• Chemists utilize Lewis dot symbols, a system of dots invented by Lewis, to keep track of valence electrons in a chemical process and ensure that the total number of electrons does not change.

  • A Lewis dot symbol is made up of the element's symbol plus one dot for each valence electron in the element's atom.

9.2 - The Ionic Bond

• Alkali metals and alkaline earth metals, on the other hand, are more likely to create cations in ionic compounds.

  • Halogens and oxygen are most likely to produce anions.

  • The electrostatic force that holds ions together in an ionic molecule is called an ionic bond.

9.3 - Lattice Energy of Ionic Compounds

• The total stability of a solid ionic compound is determined by interactions between all of these ions, not only between a single cation and a single anion. T

  • The lattice energy of an ionic solid is a quantitative indicator of its stability.

  • The energy required to completely separate one mole of a solid ionic compound into gaseous ions is known as lattice energy.

• Using Coulomb's law, we can determine the lattice energy of an ionic compound if we know its structure and composition.

  • Coulomb's law states that the potential energy (E) between two ions is proportional to the product of their charges and inversely proportional to their separation.

• By assuming that the creation of an ionic compound occurs in a series of steps, we can indirectly derive lattice energy.

  • The Born-Haber cycle connects the lattice energies of ionic compounds with ionization energies, electron affinities, and other atomic and molecular characteristics.

9.4 - The Covalent Bond

• A covalent bond is a bond between two atoms in which two electrons are shared.

  • Compounds with purely covalent bonding are known as covalent compounds.

• In the production of F2, just two valence electrons are involved.

  • The nonbonding electrons are referred to as lone pairs.

  • Lone pairs are valence electron pairs that are not involved in the creation of covalent bonds.

• Shared electron pairs are displayed as lines or pairs of dots between two atoms in a Lewis structure

  • While lone pairs are shown as pairs of dots on individual atoms.

• Lewis' octet rule is as follows: When an atom other than hydrogen is surrounded by eight valence electrons, it tends to form bonds.

• Different forms of covalent bonds can be formed by atoms.

  • One electron pair holds two atoms together in a single bond. Multiple bonds, or bonds created when two atoms share two or more pairs of electrons, hold many compounds together.

  • The covalent bond between two atoms is called a double bond when they share two pairs of electrons.

• Single covalent bonds are shorter than many ones.

  • The distance between the nuclei of two covalently bound atoms in a molecule is called bond length.

9.5 - Electronegativity

• Because electrons spend more time in the proximity of one atom than the other, a polar covalent bond, or simply a polar bond, is formed.

• Other polar bonds can be considered as a bridge between a (nonpolar) covalent link, in which electrons are shared evenly, and an ionic bond, in which electron(s) are transferred nearly completely.

  • Electronegativity, or an atom's propensity to draw the electrons in a chemical bond toward itself, is a feature that helps us identify a nonpolar covalent connection from a polar covalent link.

9.6 - Writing Lewis Structures

• Writing the compound's skeletal structure using chemical symbols and bound atoms adjacent to one another. This is a basic task for simple compounds.

  • We must either be given the knowledge or make an educated estimate for increasingly complex chemicals.

  • The middle position is usually occupied by the least electronegative element. In the Lewis structure, hydrogen and fluorine commonly occupy the terminal (end) sites.

• Count the total number of valence electrons present.

  • Add the number of negative charges to the total for polyatomic anions.

  • We add two electrons to a 2 ion because the 2 charge suggests that there are two more electrons available.

  • We remove the number of positive charges from this total for polyatomic cations.

• Connect the core atom to each of the surrounding atoms with a single covalent link. Complete the atoms bonded to the center atom's octets.

  • If electrons from the central or surrounding atoms are not involved in bonding, they must be shown as lone pairs.

  • Step 2 determines the total number of electrons that will be used.

• If the central atom has fewer than eight electrons after completing steps 1–3, try adding double electron-pair

  • Or triple bonds between the surrounding atoms and the central atom, completing the octet of the central atom with lone pairs from the surrounding atoms.

9.7 - Formal Charge and Lewis Structure

• We can determine the distribution of electrons in the molecule and draw the most feasible Lewis structure.

  • By comparing the number of electrons in an isolated atom with the number of electrons associated with the same atom in a Lewis structure.

  • The formal charge of an atom is the difference in electrical charge between its valence electrons and the number of electrons assigned to it in a Lewis structure.

9.8 - The Concept of Resonance

• The structures illustrated are resonance structures, as shown by the double-headed arrow.

  • As a result, a resonance structure is one of two or more Lewis structures for a single molecule that cannot be accurately represented by just one Lewis structure.

  • The term resonance refers to the representation of a specific molecule by two or more Lewis structures.

9.9 - Exceptions to the Octet Rule

• The BN bond in the above-mentioned chemical differs from the covalent bonds previously outlined in that the N atom contributes both electrons.

  • A coordinate covalent bond is one in which one of the atoms provides both electrons

  • It’s defined as a covalent link in which one of the atoms donates both electrons.

9.10 - Bond Enthalpy

• The enthalpy change required to break a specific bond in 1 mole of gaseous molecules is a measure of a molecule's stability.