valence shell electron pair repulsion theory

Examples of valence shell electron pair repulsion theory in the following topics:

• The Shape of Molecules

  • The following examples make use of this notation, and also illustrate the importance of including non-bonding valence shell electron pairs (colored blue) when viewing such configurations.
  • Bonding configurations are readily predicted by valence-shell electron-pair repulsion theory, commonly referred to as VSEPR in most introductory chemistry texts.
  • This simple model is based on the fact that electrons repel each other, and that it is reasonable to expect that the bonds and non-bonding valence electron pairs associated with a given atom will prefer to be as far apart as possible.
  • In each case there are four regions of electron density associated with the valence shell so that a tetrahedral bond angle is expected.
  • The compound boron trifluoride, BF3, does not have non-bonding valence electrons and the configuration of its atoms is trigonal.

• 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.

• 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.
  • According to VSEPR, the valence electron pairs surrounding an atom mutually repel each other; they adopt an arrangement that minimizes this repulsion, thus determining the molecular geometry.
  • The number of atoms bonded to a central atom combined with the number of pairs of its nonbonding valence electrons is called its steric number.
  • The bonding geometry will not be tetrahedral when the valence shell of the central atom contains nonbonding electrons.
  • Lots and lots of practice problems for VSEPR theory.

• Multielectron Atoms

  • However, when more electrons are involved, each electron (in the $n$-shell) feels not only the electromagnetic attraction from the positive nucleus, but also repulsion forces from other electrons in shells from '1' to '$n$'.
  • This causes the net force on electrons in the outer electron shells to be significantly smaller in magnitude.
  • Therefore, these electrons are not as strongly bonded to the nucleus as electrons closer to the nucleus.
  • The shielding theory also explains why valence shell electrons are more easily removed from the atom.
  • Each has 10 electrons, and the number of nonvalence electrons is two (10 total electrons minus eight valence electrons), but the effective nuclear charge varies because each has a different number of protons:

• Hund's Rule

  • Hund's Rule defines the behavior of unpaired valence shell electrons, providing insight into an atom's reactivity and stability.
  • Electrons tend to minimize repulsion by occupying their own orbital, rather than sharing an orbital with another electron.
  • In the latter case, the repulsive force increases, which separates electrons.
  • When atoms come into contact with one another, it is the outermost electrons of these atoms, or valence shell, that will interact first.
  • An atom is least stable (and therefore most reactive) when its valence shell is not full.

• Explanation of Valence Bond Theory

  • Each hydrogen atom needs one more electron to complete its valence energy shell.
  • The nitrogen atom needs three more electrons to complete its valence energy shell.
  • The nitrogen atom will share three of its electrons so that each of the hydrogen atoms now has a complete valence shell.
  • Each of the hydrogen atoms will share its electron with the nitrogen atom to complete its valence shell.
  • To complete their valence shells, they bond and share one electron with each other.

• Lone Electron Pairs

  • So far, we have only discussed geometries without any lone pairs of electrons.
  • The orbitals containing the various bonding and nonbonding pairs in the valence shell will extend out from the central atom in directions that minimize their mutual repulsions.
  • In the water molecule (AX2E2), the central atom is O, and the Lewis electron dot formula predicts that there will be two pairs of nonbonding electrons.
  • Two of the coordination positions are occupied by the shared electron-pairs that constitute the O–H bonds, and the other two by the non-bonding pairs.
  • The electron-dot structure of NH3 places one pair of nonbonding electrons in the valence shell of the nitrogen atom.

• The Shielding Effect and Effective Nuclear Charge

  • However, when more electrons are involved, each electron (in the n-shell) feels not only the electromagnetic attraction from the positive nucleus but also repulsion forces from other electrons in shells from 1 to n-1.
  • The shielding effect explains why valence shell electrons are more easily removed from the atom.
  • The more shielding that occurs, the further the valence shell can spread out.
  • The valence shell is shell 2 and contains 8 valence electrons.
  • Once again, the electron configuration is the same as in the previous examples and the number of nonvalence electrons is 2 (by losing one electron, the valence shell becomes the n=2 shell).

• sp3 Hybridization

  • To form four bonds, the atom must have four unpaired electrons; this requires that carbon's valence 2s and 2p orbitals each contain an electron for bonding.
  • 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.
  • The observed H-O-H bond angle in water (104.5°) is less than the tetrahedral angle (109.5°); one explanation for this is that the non-bonding electrons tend to remain closer to the central atom and thus exert greater repulsion on the other orbitals, pushing the two bonding orbitals closer together.

• Writing Lewis Symbols for Atoms

  • Helium is one of the noble gases and contains a full valence shell.
  • In the Lewis symbol, the electrons are depicted as two lone pair dots.
  • Electrons can inhabit a number of energy shells.
  • Different shells are different distances from the nucleus.
  • The electrons in the outermost electron shell are called valence electrons, and are responsible for many of the chemical properties of an atom.