Chemical Bonding Concepts and Lewis Symbols

The electrons in the outermost shell of an atom that participate in bonding.

A pictorial representation of an atom's valence electrons, where each dot represents an electron.

Atoms tend to gain, lose, or share electrons to achieve a stable configuration with eight valence electrons, like the noble gases.

The electrostatic force that holds oppositely charged ions together in a compound, formed by the transfer of electrons from a metal to a nonmetal.

A bond formed by the sharing of electrons between two nonmetal atoms.

A bond found in metals, where metal nuclei are surrounded by a sea of delocalized electrons, leading to high conductivity and malleability.

The energy released when oppositely charged ions come together to form an ionic compound. It helps to explain the stability of ionic compounds.

Transition metals form cations by losing electrons, but they often don't reach a full noble gas configuration. They can lose up to three electrons from the d orbitals.

Transition metals tend to lose their valence shell electrons first, followed by d electrons to achieve the desired charge on the ion.

A group of atoms held together by covalent bonds with an overall charge.

A diagram that represents the formation of covalent bonds using Lewis symbols, showing shared electron pairs as lines and unshared electron pairs as dots.

The sharing of more than one pair of electrons between two atoms.

The distance between the nuclei of two atoms in a bond.

The unequal sharing of electrons in a covalent bond.

A covalent bond where electrons are shared equally between atoms.

A covalent bond where one atom exerts a stronger attraction for bonding electrons than the other, leading to an unequal sharing of electrons.

The ability of an atom in a molecule to attract electrons to itself.

A scale that quantifies electronegativity values, ranging from 0.7 (Cs) to 4.0 (F).

Electronegativity increases across a period and decreases down a group.

The separation of positive and negative charges in a molecule, resulting in a dipole moment.

The quantitative measure of the magnitude of a dipole.

A visual representation of a molecule showing its bonding and nonbonding electrons, following specific rules to ensure octets are complete for most atoms.

The process of combining atomic orbitals to create new hybrid orbitals with different shapes and energies. This is done to explain bonding in polyatomic molecules and to explain the geometry of molecules.

A hybrid orbital formed by mixing one s and one p atomic orbital. They have two lobes that are 180 degrees apart, resulting in linear geometry.

A hybrid orbital formed by mixing one s and two p atomic orbitals. They have three lobes that lie in a trigonal plane, resulting in a trigonal planar geometry.

A hybrid orbital formed by mixing one s and three p atomic orbitals. They have four lobes that point towards the vertices of a tetrahedron, resulting in a tetrahedral geometry.

The process of using d orbitals in addition to s and p orbitals to create hybrid orbitals. This is necessary for molecules with more than four electron pairs around the central atom.

A hybrid orbital formed by mixing one s, three p, and one d atomic orbital. These orbitals have five lobes arranged in a trigonal bipyramidal geometry.

A hybrid orbital formed by mixing one s, three p, and two d atomic orbitals. These orbitals have six lobes arranged in an octahedral geometry.

To determine the hybridization of a molecule, first determine the electron-domain geometry using VSEPR theory, then determine the hybridization needed to accommodate the electron pairs based on the geometry. Finally, name the geometry based on the positions of the atoms.

The repulsion between lone pairs of electrons is greater than the repulsion between bonding pairs of electrons, causing a decrease in bond angle as the number of lone pairs increases.

The repulsion between electrons in multiple bonds is greater than the repulsion between electrons in single bonds, leading to larger bond angles in molecules with multiple bonds.

A molecule with an expanded octet has more than eight electrons surrounding the central atom, forming five or six electron domains. This results in specific electron-domain geometries.

The five electron domains in a trigonal bipyramidal structure are arranged in a three-dimensional shape where three electron pairs lie in a plane, and two are above and below this plane.

The six electron domains in an octahedral structure are arranged in a three-dimensional shape where four electron pairs lie in a plane, and two are above and below this plane.

In a trigonal bipyramidal structure, the three electron pairs in the plane are called equatorial, and the two electron pairs above and below the plane are called axial. Equatorial positions experience less repulsion compared to axial positions.

To minimize electron repulsion, lone pairs in a trigonal bipyramidal structure preferentially occupy the equatorial positions, while bonding pairs can occupy both axial and equatorial positions.

To minimize electron repulsion, lone pairs in an octahedral structure are located in opposite positions on the top and bottom of the structure.

A molecule with a bent molecular geometry has a net dipole moment because the bond dipoles do not cancel each other out.

A linear molecule with identical atoms has no overall dipole moment because the bond dipoles cancel each other out.

A linear molecule with different atoms has an overall dipole moment because the bond dipoles do not cancel each other out.

A trigonal pyramidal molecule has an overall dipole moment because the bond dipoles do not cancel each other out.

A trigonal planar molecule with identical atoms has no overall dipole moment because the bond dipoles cancel each other out.

A trigonal planar molecule with different atoms has an overall dipole moment because the bond dipoles do not cancel each other out.

Valence bond theory explains the formation of covalent bonds through the overlap of atomic orbitals on different atoms, leading to the sharing of electrons and the formation of a bonding region.

Lewis structures that differ only in the placement of electrons, representing different possible arrangements of electrons in a molecule.

The true structure of a molecule that is a blend or hybrid of all its possible resonance structures.

An important category of organic molecules that have a cyclic structure with delocalized electrons, resulting in a stable, planar structure.

A molecule with an odd number of electrons where electrons cannot be paired completely.

A molecule where the central atom has less than eight valence electrons, typically found in compounds of boron or beryllium, which have a smaller tendency to achieve an octet.

A scenario where the central atom in a molecule has more than eight valence electrons, usually found in elements from the third period or below with available 'd' orbitals.

The spatial arrangement of atoms in a molecule, which describes the bond angles and overall shape.

The arrangement of electron domains around the central atom in a molecule, based on minimizing electron-electron repulsion.

The theory that predicts the shape of a molecule by considering the repulsion between electron pairs around the central atom.

A region around an atom where electrons are localized, either as bonding pairs or lone pairs.

A pair of electrons that is shared between two atoms, forming a covalent bond.

A pair of electrons that is not involved in bonding, but rather resides on a single atom.

The angles formed by the lines joining the nuclei of atoms in a molecule.

A three-dimensional shape where four points are located at the vertices of a regular tetrahedron, with the central atom at the center.

A three-dimensional shape with a central atom surrounded by six atoms at equal distances, forming eight vertices of a regular octahedron.