lone pair

Examples of lone pair in the following topics:

• Lone Electron Pairs

  • So far, we have only discussed geometries without any lone pairs of electrons.
  • As you likely noticed in the table of geometries and the AXE method, adding lone pairs changes a molecule's shape.
  • The geometries of molecules with lone pairs will differ from those without lone pairs, because the lone pair looks like empty space in a molecule.
  • This means that there are three bonded atoms and one lone pair for a coordination number of four around the nitrogen, the same as occurs in H2O.
  • The lone pair attached to the central nitrogen creates bond angles that differ from the tetrahedral 109.5 °.

• Amines

  • Amines are compounds characterized by the presence of a nitrogen atom, a lone pair of electrons, and three substituents.
  • The amine functional group contains a basic nitrogen atom with a lone pair of electrons.
  • The general structure of an amine contains a nitrogen atom, a lone pair of electrons, and three substituents.
  • In general, the effect of alkyl groups raises the energy of the lone pair of electrons, thus elevating the basicity.
  • Additionally, the effect of the aromatic ring delocalizes the lone pair of electrons on nitrogen into the ring, resulting in decreased basicity.

• sp3 Hybridization

  • If lone electron pairs are present on the central atom, thet can occupy one or more of the sp3 orbitals.
  • For example, in the ammonia molecule, the fourth of the sp3 hybrid orbitals on the nitrogen contains the two remaining outer-shell electrons, which form a non-bonding lone pair.
  • Two of these are occupied by the two lone pairs on the oxygen atom, while the other two are used for bonding.

• Organic Reactions Overview

  • Note that the number of bonds plus lone pairs (there are 0 lone pairs) is conserved in this process.

• Introduction to Lewis Structures for Covalent Molecules

  • These non-bonding valence electrons are called 'lone pairs' of electrons and should always be indicated in Lewis diagrams.
  • These are 'lone pairs' of electrons.
  • Covalent bonds are indicated as dashes and lone pairs of electrons are shown as pairs of dots. in carbon dioxide, each oxygen atom has two lone pairs of electrons remaining; the covalent bonds between the oxygen and carbon atoms each use two electrons from the oxygen atom and two from the carbon.
  • Two H atoms, each contributing an electron, share a pair of electrons.
  • Notice the lone pairs of electrons on the oxygen atoms are still shown.

• Molecular Geometries

  • .— count as one X); and E represents the number of lone electron pairs surrounding the central atom.
  • Tetra- signifies four, and -hedral relates to a face of a solid; "tetrahedral" literally means "having four faces. " This shape is found when there are four bonds all on one central atom, with no lone electron pairs.
  • The bond angles are all 90°, and just as four electron pairs experience minimum repulsion when they are directed toward the corners of a tetrahedron, six electron pairs try to point toward the corners of an octahedron.
  • The A represents the central atom; the X represents the number of sigma bonds between the central atoms and outside atoms; and the E represents the number of lone electron pairs surrounding the central atom.
  • Apply the VSEPR model to determine the geometry of molecules where the central atom contains one or more lone pairs of electrons.

• Applying the VSEPR Model

  • The valence shell electron pair repulsion (VSEPR) model predicts the shape of individual molecules based on the extent of electron-pair electrostatic repulsion.
  • In 5-coordinated molecules containing lone pairs, these non-bonding orbitals will preferentially reside in the equatorial plane at 90° angles, with respect to no more than two axially-oriented bonding orbitals.
  • There are well known examples of 6-coordinate central atoms with one, two, and three lone pairs.
  • Notice the two lone pairs of electrons on the oxygen atom.
  • Use the molecule's formula, along with the number of bonded atoms and lone pair electrons, to determine its geometry.

• Complex Ion Equilibria and Solubility

  • A ligand is a species which can use its lone pair of electrons to form a dative covalent bond with a transition metal.
  • Empty orbitals of low energy, enabling them to accept the lone pair of electrons from ligands.
  • They form complex ions readily when their partially filled d subshell accepts donated electron pairs from other ions or molecules.
  • The number of lone pairs of electrons a cation can accept is known as the coordination number of the cation.

• Table of Geometries

  • The valence shell electron pair repulsion (VSEPR) model focuses on the bonding and nonbonding electron pairs present in the outermost (valence) shell of an atom that connects with two or more other atoms.
  • If the central atom also contains one or more pairs of non-bonding electrons, these additional regions of negative charge will behave much like those associated with the bonded atoms.
  • The orbitals containing the various bonding and non-bonding pairs in the valence shell will extend out from the central atom in directions that minimize their mutual repulsions.
  • If the central atom possesses partially occupied d-orbitals, it may be able to accommodate five or six electron pairs, forming what is sometimes called an "expanded octet."
  • Apply the VSEPR model to determine the geometry of a molecule that contains no lone pairs of electrons on the central atom.

• Reactions of Coordination Compounds

  • These are generally bound to the central atom by a coordinate covalent bond (donating electrons from a lone electron pair into an empty metal orbital).
  • Most substrates have a singlet ground-state; that is, they have lone electron pairs (e.g., water, amines, ethers).