"A chemical bond is a lasting attraction between atoms or ions that enables the formation of molecules, crystals, and other structures."
The study of the forces that hold atoms together to form molecules, including covalent, ionic, and hydrogen bonding.
Valence electrons: Electrons in an atom's outermost electron shell that are involved in chemical bonding.
Lewis symbols: Diagrams that represent the valence electrons of atoms and ions using dots and lines.
Octet rule: Atoms tend to gain, lose, or share electrons in order to attain a full valence shell of eight electrons.
Ionic bonding: The transfer of electrons from one atom to another to form ions that attract each other due to their opposite charges.
Covalent bonding: The sharing of electrons between two atoms to form a molecule.
Electronegativity: The ability of an atom to attract electrons towards itself in a covalent bond.
Polar covalent bonding: A type of covalent bonding where electrons are shared unevenly between atoms, resulting in a partially positive and partially negative charge.
Dipole moment: The measure of the separation of positive and negative charges in a polar covalent bond.
Lewis structures: Diagrams that show the arrangement of atoms in a molecule and the placement of individual electrons, based on the octet rule.
Resonance: The delocalization of electrons in a molecule, resulting in multiple Lewis structures that contribute to the overall structure and stability of the molecule.
Formal charge: The charge that would be on an atom if all electrons in a covalent bond were shared equally between atoms.
VSEPR theory: The theory that predicts the spatial arrangement of atoms in a molecule based on the repulsion between electron pairs.
Hybridization: The mixing of atomic orbitals to form hybrid orbitals that are used in covalent bonding.
Molecular orbitals: Orbitals formed by the combination of atomic orbitals in a molecule, which can be bonding or antibonding.
Bond order: The number of chemical bonds between two atoms, which can be calculated using molecular orbital theory.
Intermolecular forces: The attractions and repulsions between molecules, which affect their physical properties.
Hydrogen bonding: A special type of dipole-dipole attraction between hydrogen atoms and electronegative atoms, such as O, N, or F.
London dispersion forces: Weak attractions between molecules caused by the temporary dipole moments that arise due to the movement of electrons.
Ionic solids: Solids held together by ionic bonds, characterized by high melting and boiling points, and brittleness.
Covalent network solids: Solids held together by a three-dimensional network of covalent bonds, characterized by high melting and boiling points, and extreme hardness.
Ionic bonding: It is a type of chemical bonding in which atoms transfer electrons to each other to become stable. This type of bonding usually occurs between metals and nonmetals.
Covalent bonding: Covalent bonding occurs when atoms share electrons with one another to form molecules. This type of bonding is usually seen between nonmetals.
Metallic bonding: Metallic bonding is the type of bonding that occurs between positively charged metal ions and freely moving electrons. This type of bonding gives rise to the properties of metals, such as conductivity.
Hydrogen bonding: Hydrogen bonding is a type of dipole-dipole interaction between a hydrogen atom bonded to an electronegative atom (usually nitrogen, oxygen, or fluorine) and a lone pair of electrons on another electronegative atom in the same or another molecule.
Van der Waals forces: These are weak attractive forces between polar or nonpolar molecules due to fluctuations in electron density around the atoms of the molecules.
Dipole-dipole interactions: These are attractions between two polar molecules due to their permanent dipoles.
London dispersion forces: These are temporary dipole interactions between atoms or molecules that arise due to fluctuations in the electron clouds.
Pi bonding: Pi bonding is the bonding that occurs when two adjacent atoms share electrons in a pi bond, which is formed by the overlap of two p orbitals.
Sigma bonding: Sigma bonding occurs when two atoms share electrons in a sigma bond, which is formed by the overlap of two s orbitals.
Electrostatic Interactions: These are the interactions between ions or polar molecules, which arise due to their charged or partial-charged nature.
"The bond may result from the electrostatic force between oppositely charged ions as in ionic bonds."
"The bond may result from...the sharing of electrons as in covalent bonds."
"There are 'strong bonds' or 'primary bonds' such as covalent, ionic and metallic bonds."
"There are 'weak bonds' or 'secondary bonds' such as dipole-dipole interactions, the London dispersion force, and hydrogen bonding."
"The negatively charged electrons surrounding the nucleus and the positively charged protons within a nucleus attract each other."
"Electrons shared between two nuclei will be attracted to both of them."
"Constructive quantum mechanical wavefunction interference stabilizes the paired nuclei."
"Bonded nuclei maintain an optimal distance (the bond distance) balancing attractive and repulsive effects explained quantitatively by quantum theory."
"The atoms in molecules, crystals, metals, and other forms of matter are held together by chemical bonds."
"All bonds can be described by quantum theory, but, in practice, simplified rules and other theories allow chemists to predict the strength, directionality, and polarity of bonds."
"The octet rule and VSEPR theory are examples."
"More sophisticated theories are valence bond theory, which includes orbital hybridization and resonance."
"More sophisticated theories...include molecular orbital theory, which includes the linear combination of atomic orbitals and ligand field theory."
"Electrostatics are used to describe bond polarities and the effects they have on chemical substances."
"The strength of chemical bonds varies considerably."
"The octet rule is an example [of a theory] that allows chemists to predict the strength, directionality, and polarity of bonds."
"The VSEPR theory is an example [of a theory] that allows chemists to predict the strength, directionality, and polarity of bonds."
"Orbital hybridization is a component of valence bond theory."
"Ligand field theory is a component of molecular orbital theory."