Chapter 9 - Chemical Bonding I
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.
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.