Glossary
Bonding Pairs
Pairs of valence electrons that are shared between two atoms, forming a covalent bond and holding the atoms together.
Example:
In a methane molecule (CH₄), there are four bonding pairs of electrons, each forming a single bond between carbon and a hydrogen atom.
Central Atom
The atom in a molecule or polyatomic ion to which other atoms are bonded, typically the least electronegative atom (excluding hydrogen).
Example:
In the molecule phosphorus trichloride (PCl₃), phosphorus is the central atom because it is bonded to all three chlorine atoms.
Covalent LDSs
Lewis diagrams for covalent compounds, depicting the sharing of electrons between nonmetal atoms to achieve stable electron configurations.
Example:
Drawing the covalent LDS for sulfur dioxide (SO₂) involves forming a double bond and lone pairs to satisfy the octets of sulfur and oxygen by sharing electrons.
Double Bonds
A covalent bond formed by the sharing of two pairs of electrons between two atoms, represented by two dashes.
Example:
Ethene (C₂H₄) contains a double bond between the two carbon atoms, which restricts rotation and influences its reactivity.
Empirical Formula
The simplest whole-number ratio of atoms in a compound, used to identify the elements and their relative quantities.
Example:
While the molecular formula for hydrogen peroxide is H₂O₂, its empirical formula is HO, representing the simplest ratio of hydrogen to oxygen atoms.
Expanded Octets
A situation where a central atom in a molecule has more than eight valence electrons, possible for elements in Period 3 or below due to the availability of d-orbitals.
Example:
Sulfur tetrafluoride (SF₄) features an expanded octet around the sulfur atom, which is bonded to four fluorine atoms and has one lone pair, resulting in 10 valence electrons.
Incomplete Octets
A situation where an atom in a molecule has fewer than eight valence electrons, typically observed in elements like hydrogen, beryllium, and boron.
Example:
Boron trichloride (BCl₃) is an example of a molecule with an incomplete octet on the central boron atom, which only has six valence electrons.
Ionic LDSs
Lewis diagrams specifically for ionic compounds, illustrating the transfer of electrons from a metal to a nonmetal to form stable ions.
Example:
The ionic LDS for potassium iodide (KI) shows potassium losing an electron to become K⁺ and iodine gaining one to become I⁻, both achieving noble gas configurations.
Lewis Diagrams
Visual tools that show how atoms bond together, displaying valence electrons and bonds to help understand molecular shapes and reactivity.
Example:
Drawing the Lewis diagram for water (H₂O) helps you visualize its two O-H bonds and two lone pairs on the oxygen atom.
Localized Electron Model
A model that assumes electrons in a molecule are localized, meaning they are either associated with a particular atom (lone pairs) or shared between two atoms (bonding pairs).
Example:
The localized electron model helps explain why electrons in a C-H bond stay between the carbon and hydrogen, rather than moving freely throughout the molecule.
Lone Pairs
Pairs of valence electrons that are not shared with another atom in a covalent bond and belong exclusively to one atom.
Example:
The nitrogen atom in ammonia (NH₃) has one lone pair of electrons, which significantly influences its trigonal pyramidal molecular geometry.
Octet Rule
A chemical rule stating that atoms tend to combine in such a way that each atom has eight electrons in its valence shell, achieving a stable electron configuration.
Example:
In carbon dioxide (CO₂), carbon forms double bonds with each oxygen to satisfy the octet rule, ensuring all atoms have eight valence electrons.
Odd Numbers of Electrons
Molecules that contain an odd total number of valence electrons, making it impossible for all electrons to be paired and satisfy the octet rule for every atom.
Example:
Nitrogen dioxide (NO₂) is a radical because it has an odd number of electrons (17 valence electrons), leading to an unpaired electron and high reactivity.
Single Bonds
A covalent bond formed by the sharing of one pair of electrons between two atoms, represented by a single dash.
Example:
In a molecule of hydrogen fluoride (HF), there is one single bond between the hydrogen and fluorine atoms.
Total Valence Electrons
The sum of all valence electrons from every atom in a molecule or polyatomic ion, which must be accounted for when drawing its Lewis structure.
Example:
To draw the Lewis structure for the carbonate ion (CO₃²⁻), you first calculate the total valence electrons by adding 4 (for C) + 3*6 (for O) + 2 (for the charge) = 24 electrons.
Triple Bonds
A covalent bond formed by the sharing of three pairs of electrons between two atoms, represented by three dashes.
Example:
Ethyne (C₂H₂) features a triple bond between its two carbon atoms, making it a very linear and reactive molecule.