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A discussion about covalent bonding and molecular geometry

Molecular Geometry Photo by: Vladimir Fedorchuk Molecules, from simple diatomic ones to macromolecules consisting of hundreds of atoms or more, come in many shapes and sizes. The term "molecular geometry" is used to describe the shape of a molecule or polyatomic ion as it would appear to the eye if we could actually see one.

For this discussion, the terms "molecule" and "molecular geometry" pertain to polyatomic ions as well as molecules. Molecular Orbitals When two or more atoms approach each other closely enough, pairs of valence shell electrons frequently fall under the influence of two, and sometimes more, nuclei. Electrons move to occupy new regions of space new orbitals—molecular orbitals that allow them to "see" the nuclear charge of multiple nuclei.

When this activity results in a lower overall energy for all involved atoms, the atoms remain attached and a molecule has been formed. In such cases, we refer to the interatomic attractions holding the atoms together as covalent bonds.

  • In addition, with increased electron density in the spaces between the nuclei, nucleus-nucleus repulsions are minimized;
  • With oxygen atoms bonded to these sets of electrons, the oxygen—carbon—oxygen axis is a straight line, making the molecular geometry a straight line;
  • Though the approaches differ, they must ultimately converge because they describe the same physical reality;
  • However, let us start with some simple examples and your imagination will help you to extend this discussion to more complicated ones;
  • When applying VSEPR theory, attention is first focused on the electron pairs of the central atom , disregarding the distinction between bonding pairs and lone pairs.

These molecular orbitals may be classified according to strict mathematical probabilistic determinations of atomic behaviors. Though we may be oversimplifying a highly complex mathematics, it may help one to visualize sigma molecular orbitals as those that build up electron density along the internuclear axis connecting bonded nuclei, and pi molecular orbitals as those that build up electron density above and below the internuclear axis.

At their simplest levels, both approaches ignore nonvalence shell electrons, treating them as occupants of a discussion about covalent bonding and molecular geometry orbitals so similar to the original premolecular formation atomic orbitals that they are localized around the original nuclei and do not participate in bonding. The two approaches diverge mainly with respect to how they treat the electrons that are extensively influenced by two or more nuclei.

Though the approaches differ, they must ultimately converge because they describe the same physical reality: In MO theory, there are three types of molecular orbitals that electrons may occupy. Nonbonding molecular orbitals closely resemble atomic orbitals localized around a single nucleus.

They are called nonbonding because their occupation by electrons confers no net advantage toward keeping the atoms together. Bonding molecular orbitals correspond to regions where electron density builds up between two, sometimes more, nuclei. When these orbitals are occupied by electrons, the electrons "see" more positive nuclear charge than they would if the atoms had not come together. In addition, with increased electron density in the spaces between the nuclei, nucleus-nucleus repulsions are minimized.

Bonding orbitals allow for increased electron-nucleus attraction and decreased nucleus-nucleus repulsion, therefore electrons in such orbitals tend to draw atoms together and bond them to each other. One antibonding molecular orbital is formed for each bonding molecular orbital that is formed.

Antibonding orbitals tend to localize electrons outside the regions between nuclei, resulting in significant nucleus-nucleus repulsion—with little, if any, improvement in electron-nucleus attraction. Electrons in antibonding orbitals work against the formation of bonds, which is why they are called antibonding. However, when the number of bonding electrons is matched by the number of antibonding electrons, there is actually a dis advantage to having the atoms stay together, therefore no molecule forms.

Valence bond VB theory assumes that atoms form covalent bonds as they share pairs of electrons via overlapping valence shell orbitals. A single covalent bond forms when two atoms share a pair of electrons via the sigma overlap of two atomic orbitals—a valence orbital from each atom. A double bond forms when two atoms share two pairs of electrons, one pair via a sigma overlap of two atomic orbitals and one via a pi overlap. A triple bond forms by three sets of orbital overlap, one of the sigma type and two of the pi type, accompanied by the sharing of three pairs of electrons via those overlaps.

When a pair of valence shell electrons is localized at only one atom, that is, when the pair is not shared between atoms, it is called a lone or nonbonding pair. Let us apply this greatly simplified picture of VB theory to three diatomic molecules: VB theory says that an H 2 molecule forms when a 1 s orbital containing an electron that belongs to one atom overlaps a 1 s orbital with an electron of opposite spin belonging to the other, creating a sigma molecular orbital containing two electrons.

The two nuclei share the pair of electrons and draw together, giving both electrons access to the positive charge of both nuclei. Diatomic fluorine, F 2forms similarly, via the sigma overlap of singly occupied 2 a discussion about covalent bonding and molecular geometry orbitals.

Molecular Geometry

The HF molecule results from the sharing of a pair of electrons whereby an electron in a hydrogen 1 s orbital experiences sigma overlap with an electron in a fluorine 2 p orbital. Molecular Geometries This VB approach allows us to return to the focus of our discussion.

  1. As has already been pointed out, the result of this isotropy is that ions stack together in the locations necessary to achieve the lowest energy and in this way give rise to the common packing patterns characteristic of many ionic solids.
  2. When applying VSEPR theory, attention is first focused on the electron pairs of the central atom , disregarding the distinction between bonding pairs and lone pairs. In H 2 CO, the carbon atom's eight valence electrons are grouped into three sets, corresponding to the two single bonds and the one double bond.
  3. These molecular orbitals may be classified according to strict mathematical probabilistic determinations of atomic behaviors.

The geometry of a molecule or polyatomic ion is determined by the positions of individual atoms and their positions relative to one another. It can get very complicated. However, let us start with some simple examples and your imagination will help you to extend this discussion to more complicated ones. What happens when two atoms are bonded together in a diatomic molecule? The only possible geometry is a straight line. Hence, such a molecular geometry or shape is called "linear.

When four atoms bond together, they may form a straight or a zigzag line, a square or other two-dimensional shape in which all four atoms occupy the same flat plane, or they may take on one of several three-dimensional geometries such as a pyramid, with one atom sitting atop a base formed by the other three atoms.

With so many possibilities, it may come as a surprise that we can "predict" the shape of a molecule or polyatomic ion using some basic assumptions about electron-electron repulsions. We start by recognizing that, ultimately, the shape of a molecule is the equilibrium geometry that gives us the lowest possible energy for the system. Such a geometry comes about as the electrons and nuclei settle into positions that minimize nucleus-nucleus and electron-electron repulsions, and maximize electron-nucleus attractions.

Modern computer programs allow us to perform complex mathematical calculations for multiatomic systems with high predictive accuracy. However, without doing all the mathematics, we may "predict" molecular geometries quite well using VB theory. Valence shell electron pair repulsion approach.

In the valence shell electron pair repulsion VSEPR approach to molecular geometry, we begin by seeing the valence shell of a bonded atom as a spherical surface. Repulsions among pairs of valence electrons force the pairs to locate on this surface as far from each other as possible. Based on such considerations, somewhat simplified herein, we determine where all the electron pairs on the spherical surface of the atom "settle down," and identify which of those pairs correspond to bonds.

Once we know which pairs of electrons bond or glue atoms together, we can more easily picture the shape of the corresponding simple molecule.

  1. Once we know which pairs of electrons bond or glue atoms together, we can more easily picture the shape of the corresponding simple molecule.
  2. As has already been pointed out, the result of this isotropy is that ions stack together in the locations necessary to achieve the lowest energy and in this way give rise to the common packing patterns characteristic of many ionic solids. When four atoms bond together, they may form a straight or a zigzag line, a square or other two-dimensional shape in which all four atoms occupy the same flat plane, or they may take on one of several three-dimensional geometries such as a pyramid, with one atom sitting atop a base formed by the other three atoms.
  3. The two approaches diverge mainly with respect to how they treat the electrons that are extensively influenced by two or more nuclei. All four pairs are bonding , so the ion is predicted to be a regular tetrahedron, which it indeed is.
  4. Such a geometry comes about as the electrons and nuclei settle into positions that minimize nucleus-nucleus and electron-electron repulsions, and maximize electron-nucleus attractions.

However, in using VSEPR, we must realize that in a double or triple bond, the sigma and pi orbital overlaps, and the electrons contained therein, are located in the same basic region between the two atoms. Thus, the four electrons of a double bond or the six electrons of a triple bond are not independent of one another, but form coordinated "sets" of four or six electrons that try to get as far away from other sets of electrons as possible. In an atom's valence shell, a lone pair of electrons or, collectively, the two, four, or six electrons of a single, double, or triple bond each form a set of electrons.

It is repulsions among sets of valence shell electrons that determine the geometry around an atom. Their Lewis structures are and and In CO 2the double bonds group the carbon atom's eight valence electrons into two sets.

The two sets get as far as possible from each other by residing on opposite sides of the carbon atom, creating a straight line extending from one set of electrons through the carbon nucleus to the other. With oxygen atoms bonded to these sets of electrons, the oxygen—carbon—oxygen axis is a straight line, making the molecular geometry a straight line.

Carbon dioxide is a linear molecule.

  • When this activity results in a lower overall energy for all involved atoms, the atoms remain attached and a molecule has been formed;
  • Molecules with multiple bonds There are further rules in VSEPR theory that simplify the discussion of species with multiple bonds and of species in which resonance must be considered;
  • Note that the shape of the molecule is determined by the disposition of the atoms, not the disposition of the electron pairs;
  • When applying VSEPR theory, attention is first focused on the electron pairs of the central atom , disregarding the distinction between bonding pairs and lone pairs.

In H 2 CO, the carbon atom's eight valence electrons are grouped into three sets, corresponding to the two single bonds and the one double bond. Formaldehyde has the geometry of a trigonal or triangular planar molecule, "planar" emphasizing that the carbon occupies the same plane as the three peripheral atoms.