Chemical Bonding

FIELDS OF STUDY: Organic Chemistry; Physical Chemistry; Molecular Biology

ABSTRACT

Chemical bonding, through bond formation and bond dissociation, is the fundamental process by which all chemical reactions occur and depends on a number of factors deriving from the electronic structure of atoms. The primary types of bonds are ionic and covalent.

Visualizing Chemical Bonding

Chemical bonding can be thought of as "hand shaking" between atoms. A good understanding of chemical bonding requires a sound appreciation of atomic structure, particularly the arrangement of electrons in atoms. The distribution of electrons in atoms is reflected in the arrangement of chemical bonds that can be formed about them. This arrangement determines the three-dimensional structure of molecules and the physical properties of the corresponding compounds. The basic mechanism of chemical bonding is the transfer of electrons between the atoms that have formed the bond to each other. This involves only the outermost or valence electrons in each atom.

Atomic Structure and Atomic Orbitals

The modern theory of atomic structure is based on the principles of quantum mechanics. According to this mathematical description of the atom, supported by physical observations, atoms contain a very small, dense nucleus composed of protons and neutrons. Surrounding the nucleus is a much larger, diffuse cloud of electrons, equal in number to the protons in the nucleus. The electrons are constrained to particular regions of space about the nucleus, called atomic orbitals, according to the specific energies they are allowed to have in those regions.

The rules of quantum mechanics dictate the number and distribution of electrons in the atomic orbitals of any particular atom, and this in turn determines the valence properties of the atom. The lowest level, or s, orbitals are spherical about the nucleus. The three p orbitals have a figure-eight shape and are oriented at right angles to each other, much like the x, y, and z axes of the Cartesian coordinate system. Four of the five d orbitals have a shape like two figure eights crossing each other, while the fifth is shaped like an oversized p orbital with a donut about its middle. The d orbitals are oriented at angles that place them between the p orbitals. These three sets of atomic orbitals are sufficient to describe the most common elements. A set of seven f orbitals, with more complex shapes and orientations, is available in higher elements.

The Chemical Bonding Continuum

The most common types of chemical bonds are ionic and covalent bonds, but it would be incorrect to think that all bonds fit neatly into one of these two categories. The nature of a chemical bond is affected by several factors, including the relative electronegativities of the bonded atoms and their size-related electron densities. These factors determine the extent to which charge separation occurs in the bond. In an ionic bond, one atom has lost electrons to become a positively charged ion, while the other has gained electrons to become a negatively charged ion. The charge separation between them is complete. In a purely covalent bond, the two electrons in the chemical bond between two atoms belong equally to both. When the atoms are dissimilar, however, one may exert a greater influence over the bonding electrons, resulting in an uneven distribution of charge between the two atoms, often referred to as a partial charge. This imparts a certain amount of ionic character to the covalent bond. Thus, chemical bonds actually exist on an ionic-covalent continuum.

Ionic Bonds

Ionic bonds are formed by the attraction between two oppositely charged ions. A simple common example is sodium chloride, each molecule of which consists of one positively charged sodium ion and one negatively charged chloride ion. The attraction between the ions is very strong, and ionic compounds are normally crystalline solids with high melting points. This also covers a broad range of values, as there are ionic compounds with low melting points and ionic compounds with very high melting points. Niobium carbide, for example, melts at 3,500 degrees Celsius, and nickel oxide melts at about 1,990 degrees Celsius. Magnesium chlorate hexahydrate, on the other hand, melts at a mere 35 degrees Celsius. There are higher and lower melting points of ionic solids, but the vast majority fall somewhere within this range.

The strength of an ionic bond is a reflection of the charge density of the ions. The charge density is a measure of the amount of electrical charge contained within a specific volume or region of space. Elements that form ions of small diameter have high charge density and are thus able to fit together in a crystal lattice (a regular array of ions, atoms, or molecules in the structure of a crystal) more compactly than ions with larger diameters. Such compounds tend to have higher melting points and hardness than compounds composed of larger or differently sized ions.

Ionic bonds can be strongly affected by solvents, particularly water. Most dissolve in water, in degrees ranging from sparingly soluble to completely soluble. Sodium chloride and similar salts are prime examples; they dissolve as the water molecules surround and stabilize both the positive and negative ions that make up the solid compound. As long as enough free solvent is available, the material will continue to dissolve, but when the amount of available solvent molecules is not sufficient to stabilize any more ions, dissolution ceases and the solution is said to be saturated. The energy state of the dissolved ions typically becomes lower as their entropy (degree of disorder) increases.

Covalent Bonds

Covalent bonds are the opposite of ionic bonds, as they represent the mutual sharing of electrons rather than the complete transfer of electrons that defines ions and ionic bonds. A covalent bond forms between two atoms by the overlap of atomic orbitals of corresponding energy. This allows two electrons, one from each atom, to occupy the resulting molecular orbital in a way that satisfies the valence-shell electron distribution of both atoms. The simplest example of this is the covalent bond between the two hydrogen atoms of the hydrogen molecule H₂. Each hydrogen atom has just one electron in an s atomic orbital, designated 1s because it has the lowest possible energy level, but an s orbital can contain up to two electrons. By combining their respective 1s orbitals and forming a directly overlapping molecular orbital, each hydrogen atom effectively contains two electrons in its 1s orbital, thus completely filling its valence shell. The same principle applies to the formation of all covalent bonds, regardless of the atoms involved.

Covalent bonds can display pronounced bond polarity when the bond is between dissimilar atoms. The surrounding environment can also induce polarity in an otherwise nonpolar covalent bond. The water molecule, H₂O, is the quintessential example of this effect. The oxygen atom of the water molecule forms a covalent bond to two separate hydrogen atoms, thus completely filling the valence shell of all three atoms (the valence shell of an oxygen atom contains six electrons but is capable of containing a maximum of eight). However, because the oxygen atom is very electronegative, meaning that it has a strong tendency to attract electrons, and the hydrogen atoms are very electropositive, the electrons in the covalent bonds between them are pulled toward the oxygen atom, leaving the hydrogen atoms with a positive partial charge and causing the positive end of the O−H bond to be attracted to the negative end of an O−H bond on another molecule. This forms what is known as a hydrogen bond, which is an example of dipole-dipole attraction. A dipole is an object in which the positive and negative charges have separated, such as a molecule that is more positively charged at one end and more negatively charged at the other.

Another example of a hydrogen bond is the organic compound fluoroethane, C₂H₅F. In this compound, the two carbon atoms are bonded to each other, with one also bonded to three hydrogen atoms and the other bonded to two hydrogen atoms and one fluorine atom. The highly electronegative fluorine atom induces a dipole in the covalent bond it forms with the carbon atom, which in turn increases the polarity of the covalent bonds between that carbon atom and the two hydrogen atoms. Not only does the entire fluoroethane molecule have an overall dipole due to the presence of the fluorine atom but each covalent bond within the molecule also has a stronger dipole character in comparison to those of the corresponding covalent bonds in the parent molecule, ethane (C₂H₆).

Covalent sigma bonds are formed directly between two atoms. A second type of covalent bond, the pi bond, forms when orbitals on adjacent atoms do not overlap directly but do so in a side-by-side manner instead. The overlap of p orbitals on adjacent carbon atoms, for example, forms very stable pi bonds alongside the sigma bond between the two carbon atoms. The higher-energy d orbitals can also take part in pi bonding, but this is not as commonly observed due to the energy differences between p and d orbitals.

Another type of covalent bond is the coordination bond, also known as a dative bond or a dipolar bond. A coordination bond is formed when one atom provides both of the electrons in a covalent bond. When ions or molecules bond to a central metal atom in this manner, the compound formed is known as a coordination complex, and the ions and molecules surrounding the central atom are called ligands. If the compound has a net electrical charge, it is a complex ion. Like all other types of bonds, coordination bonds range in strength from very weak to very strong.

Ionic or Covalent?

A great many compounds exist in which all atoms are bonded covalently but readily dissociate into ions. Organic acids and bases are typical of this type of compound, as interaction with water has the ability to dissociate the molecule by extracting a hydrogen ion (H⁺), leaving the remainder as a negatively charged ion. The ease with which this occurs varies greatly from compound to compound and is highly dependent on the compound’s molecular structure. The transmission of charge between adjacent atoms, known as the inductive effect, can greatly enhance the ability of the compound to convert the covalent bond of a particular hydrogen atom to an ionic bond. Acetic acid (CH₃COOH) is a weak acid that dissolves primarily as the whole molecule, with only about one in ten thousand molecules dissociating into H⁺ and an acetate ion (C₂H₃O₂⁻). By contrast, chloroacetic acid (ClCH₂COOH), though still considered a weak acid, dissociates one hundred times more readily than acetic acid.

Induction also allows the formation of chemical bonds to atoms that are normally nonreactive. Xenon is one of the so-called inert gases, which means that the valence electron shells of xenon atoms are already completely filled. However, when xenon reacts with the extremely electronegative element fluorine, a stable crystalline compound identified as xenon hexafluoride (XeF₆) can be produced, as can other compounds of xenon and fluorine. The high electronegativity of the fluorine atoms and the relatively large size of the xenon atom allow the fluorine atoms to draw valence electrons from the xenon atom into bond formation. The formation of such compounds, like the dissociation of organic acids and bases, blurs the line between ionic and covalent bonds and reinforces the fact that chemical bonds exist on a continuum between covalent and ionic.

Reactions of metal atoms with relatively electron-rich organic compounds form compounds through coordination bonds. Like all other types of bonds, coordination bonds range in strength from very weak to very strong. Coordination bonding does not necessarily involve the transfer of electrons. In many cases, coordination compounds form just by the overlap of orbitals of appropriate energy configuration. A well-known example of such a compound is bis(benzene)chromium, Cr(C₆H₆)₂. This compound is formed in the gas phase by heating a sample of chromium metal in an atmosphere of benzene vapor. Bis(benzene)chromium deposits on the inside of the reaction vessel as a dark brown-black solid. The compound forms as the p orbitals of the benzene molecules coordinate to the vacant valence orbitals of the chromium atom. Coordination bonds are highly reversible and are an extremely important aspect of many catalytic processes.

PRINCIPAL TERMS

• covalent bond: a type of chemical bond in which electrons are shared between two adjacent atoms.
• dipole-dipole: describes a type of interaction, either attraction or repulsion, between two molecular dipoles, which are molecules that are polarized due to the greater concentration of electrons in one area.
• hydrogen bond: a weak type of chemical bond formed by the attraction of a hydrogen atom to an electronegative atom—an atom with a strong tendency to attract electrons—in the same or another molecule.
• ionic bond: a type of chemical bond formed by mutual attraction between two ions of opposite charges.
• partial charge: a term used to indicate a degree of charge separation in bonded atoms due to the different electronegativities of the atoms at either end of the bond.

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