"A chemical bond is a lasting attraction between atoms or ions that enables the formation of molecules, crystals, and other structures."
The way in which atoms combine to form molecules, including covalent bonding and ionic bonding.
Atomic structure: The basic structure of an atom, including the subatomic particles and their arrangement within the nucleus and electron shells.
Electronegativity: A measure of how strongly an atom attracts electrons towards itself in a chemical bond.
Bonding types: Covalent, ionic, metallic, and hydrogen bonding, each with its unique characteristics.
Lewis structures: A notation system that depicts the bonding between atoms and their electron distribution.
Molecular geometry: The three-dimensional arrangement of atoms in a molecule.
Hybridization: The combination of atomic orbitals to form hybrid orbitals that account for the molecular geometry and the bonding properties of a compound.
Valence bond theory: A model that describes covalent bonding as the overlap of atomic orbitals.
Molecular orbital theory: A model that describes the bonding in molecules as the result of the interaction of atomic orbitals to form molecular orbitals.
Intermolecular forces: The attractive and repulsive forces between molecules that influence their physical and chemical properties.
Van der Waals forces: The attractive forces between molecules that result from fluctuations in electron density.
Dipole-dipole interactions: The attractive forces between polar molecules due to the partial charges on their atoms.
Hydrogen bonding: The strong dipole-dipole interaction between molecules that contain hydrogen bonded to nitrogen, oxygen, or fluorine.
London dispersion forces: The weak electrostatic attraction between nonpolar molecules that results from the instantaneous fluctuations in electron density.
Polarizability: The ability of a molecule to deform its electron cloud in response to an external electric field.
Bond strengths: The energies required to break a bond, including bond dissociation energy and bond enthalpy.
Spectroscopic methods: Tools used to study the electronic and vibrational energies of molecules, including infrared and UV-visible spectrometry.
Molecular dynamics simulations: Computational methods that model molecular behavior and allow for the prediction of molecular properties.
Molecular symmetry: The ways in which the arrangement of atoms in a molecule can be mirrored or rotated.
Chemical reactions: The ways in which molecules can combine or break apart to form different compounds.
Covalent bonding: A type of molecular bonding where two atoms share electrons to form a stable and neutral molecule.
Ionic bonding: A type of molecular bonding where two or more atoms exchange electrons to form an ionically bonded molecule.
Hydrogen bonding: A type of molecular bonding where hydrogen atoms form a weak bond with another atom due to the presence of a partial negative charge on the other atom.
Van der Waals bonding: A type of molecular bonding where a weak attraction occurs between neutral molecules due to the random fluctuations in electron density.
Metallic bonding: A type of molecular bonding seen in metals where the valence electrons are shared between many atoms, resulting in a highly conductive structure.
Polar covalent bonding: A type of molecular bonding where electrons are shared unequally between two atoms with different electronegativity values.
Nonpolar covalent bonding: A type of molecular bonding where electrons are shared equally between two atoms with similar electronegativity values.
Coordinate covalent bonding: A type of molecular bonding where one atom contributes both electrons to the bond, rather than each atom contributing one.
Pi bonding: A type of molecular bonding where electrons are shared in a region parallel to the axis between two atoms; commonly found in double and triple bonds between carbon atoms.
Sigma bonding: A type of molecular bonding where electrons are shared in a region on the line directly between two atoms; commonly found in single bonds between atoms.
"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."