The Covalent Bond

Comparison between Covalent and Ionic Compounds

Covalent and ionic compounds have distinct physical properties.

Two Classes of Compounds

Compounds are defined as substances containing two or more different chemical elements. They have distinct chemical structures characterized by a fixed ratio of atoms held together by chemical bonds. Here, we discuss two classes of compounds based on the bond type that holds the atoms together: ionic and covalent.

Covalent Compounds

Covalent bonds are characterized by the sharing of electrons between two or more atoms. These bonds mostly occur between nonmetals or between two of the same (or similar) elements.Two atoms with similar electronegativity will not exchange an electron from their outermost shell; the atoms instead share electrons so that their valence electron shell is filled.

Examples of compounds that contain only covalent bonds are methane (CH₄), carbon monoxide (CO), and iodine monobromide (IBr).

Covalent bonding between hydrogen atoms: Since each hydrogen atom has one electron, they are able to fill their outermost shells by sharing a pair of electrons through a covalent bond.

Ionic Compounds

Ionic bonding occurs when there is a large difference in electronegativity between two atoms. This large difference leads to the loss of an electron from the less electronegative atom and the gain of that electron by the more electronegative atom, resulting in two ions. These oppositely charged ions feel an attraction to each other, and this electrostatic attraction constitutes an ionic bond.

Ionic bonding occurs between a nonmetal, which acts as an electron acceptor, and a metal, which acts as an electron donor. Metals have few valence electrons, whereas nonmetals have closer to eight valence electrons; to easily satisfy the octet rule, the nonmetal will accept an electron donated by the metal. More than one electron can be donated and received in an ionic bond.

Some examples of compounds with ionic bonding include NaCl, KI, MgCl₂.

Formation of sodium fluoride (NaF): The transfer of an electron from a neutral sodium atom to a neutral fluorine atom creates two oppositely charge ions: Na⁺ and F^–. Attraction of the oppositely charged ions is the ionic bond between Na and F.

Effect on Physical Properties

Covalent and ionic compounds can be differentiated easily because of their different physical properties based on the nature of their bonding. Here are some differences:

1. At room temperature and normal atmospheric pressure, covalent compounds may exist as a solid, a liquid, or a gas, whereas ionic compounds exist only as solids.
2. Although solid ionic compounds do not conduct electricity because there are no free mobile ions or electrons, ionic compounds dissolved in water make an electrically conductive solution. In contrast, covalent compounds do not exhibit any electrical conductivity, either in pure form or when dissolved in water.
3. Ionic compounds exist in stable crystalline structures. Therefore, they have higher melting and boiling points compared to covalent compounds.

Single Covalent Bonds

Single covalent bonds are sigma bonds, which occur when one pair of electrons is shared between atoms.

Hierarchical Structure of the Atom

There are four hierarchical levels that describe the position and energy of the electrons an atom has. Here they are listed along with some of the possible values (or letters) they can have:

1. Principal energy levels (1, 2, 3, etc.)
2. Sublevels (s, p, d, f)
3. Orbitals
4. Electrons

Principal energy levels are made out of sublevels, which are in turn made out of orbitals, in which electrons are found.

Atomic Orbitals

An atomic orbital is defined as the probability of finding an electron in an area around an atom’s nucleus. Generally, orbital shapes are drawn to describe the region in space in which electrons are likely to be found. This is referred to as “electron density.”

Atomic orbitals: The shapes of the first five atomic orbitals are shown in order: 1s, 2s, and the three 2p orbitals. Both blue and orange-shaded regions represent regions in space where electrons can be found ‘belonging’ to these orbitals.

Sigma Bonds

Covalent bonding occurs when two atomic orbitals come together in close proximity and their electron densities overlap. The strongest type of covalent bonds are sigma bonds, which are formed by the direct overlap of orbitals from each of the two bonded atoms. Regardless of the atomic orbital type, sigma bonds can occur as long as the orbitals directly overlap between the nuclei of the atoms.

Orbital overlaps and sigma bonds: These are all possible overlaps between different types of atomic orbitals that result in the formation of a sigma bond between two atoms. Notice that the area of overlap always occurs between the nuclei of the two bonded atoms.

Single covalent bonds occur when one pair of electrons is shared between atoms as part of a molecule or compound. A single covalent bond can be represented by a single line between the two atoms. For instance, the diatomic hydrogen molecule, H₂, can be written as H—H to indicate the single covalent bond between the two hydrogen atoms.

Sigma bond in the hydrogen molecule: Higher intensity of the red color indicates a greater probability of the bonding electrons being localized between the nuclei.

Double and Triple Covalent Bonds

Double and triple bonds, comprised of sigma and pi bonds, increase the stability and restrict the geometry of a compound.

Double and Triple Covalent Bonds

Covalent bonding occurs when electrons are shared between atoms. Double and triple covalent bonds occur when four or six electrons are shared between two atoms, and they are indicated in Lewis structures by drawing two or three lines connecting one atom to another. It is important to note that only atoms with the need to gain or lose at least two valence electrons through sharing can participate in multiple bonds.

Bonding Concepts

Hybridization

Double and triple bonds can be explained by orbital hybridization, or the ‘mixing’ of atomic orbitals to form new hybrid orbitals. Hybridization describes the bonding situation from a specific atom’s point of view. A combination of s and p orbitals results in the formation of hybrid orbitals. The newly formed hybrid orbitals all have the same energy and have a specific geometrical arrangement in space that agrees with the observed bonding geometry in molecules. Hybrid orbitals are denoted as sp^x, where s and p denote the orbitals used for the mixing process, and the value of the superscript x ranges from 1-3, depending on how many p orbitals are required to explain the observed bonding.

Hybridized orbitals: A schematic of the resulting orientation in space of sp³ hybrid orbitals. Notice that the sum of the superscripts (1 for s, and 3 for p) gives the total number of formed hybrid orbitals. In this case, four orbitals are produced which point along the direction of the vertices of a tetrahedron.

Pi Bonds

Pi, or [latex]\pi[/latex], bonds occur when there is overlap between unhybridized p orbitals of two adjacent atoms. The overlap does not occur between the nuclei of the atoms, and this is the key difference between sigma and pi bonds. For the bond to form efficiently, there has to be a proper geometrical relationship between the unhybridized p orbitals: they must be on the same plane.

Pi bond formation: Overlap between adjacent unhybridized p orbitals produces a pi bond. The electron density corresponding to the shared electrons is not concentrated along the internuclear axis (i.e., between the two atoms), unlike in sigma bonds.

Multiple bonds between atoms always consist of a sigma bond, with any additional bonds being of the π type.

Examples of Pi Bonds

The simplest example of an organic compound with a double bond is ethylene, or ethene, C₂H₄. The double bond between the two carbon atoms consists of a sigma bond and a π bond.

Ethylene bonding: An example of a simple molecule with a double bond between carbon atoms. The bond lengths and angles (indicative of the molecular geometry) are indicated.

From the perspective of the carbon atoms, each has three sp² hybrid orbitals and one unhybridized p orbital. The three sp² orbitals lie in a single plane at 120-degree angles. As the carbon atoms approach each other, their orbitals overlap and form a bond. Simultaneously, the p orbitals approach each other and form a bond. To maintain this bond, the p orbitals must stay parallel to each other; therefore, rotation is not possible.

A triple bond involves the sharing of six electrons, with a sigma bond and two [latex]\pi[/latex] bonds. The simplest triple-bonded organic compound is acetylene, C₂H₂. Triple bonds are stronger than double bonds due to the the presence of two [latex]\pi[/latex] bonds rather than one. Each carbon has two sp hybrid orbitals, and one of them overlaps with its corresponding one from the other carbon atom to form an sp-sp sigma bond. The remaining four unhybridized p orbitals overlap with each other and form two [latex]\pi[/latex] bonds. Similar to double bonds, no rotation around the triple bond axis is possible.

Observable Consequences of Multiple Bonds

Bond Strength

Covalent bonds can be classified in terms of the amount of energy that is required to break them. Based on the experimental observation that more energy is needed to break a bond between two oxygen atoms in O₂ than two hydrogen atoms in H₂, we infer that the oxygen atoms are more tightly bound together. We say that the bond between the two oxygen atoms is stronger than the bond between two hydrogen atoms.

Experiments have shown that double bonds are stronger than single bonds, and triple bonds are stronger than double bonds. Therefore, it would take more energy to break the triple bond in N₂ compared to the double bond in O₂. Indeed, it takes 497 kcal/mol to break the O₂ molecule, while it takes 945 kJ/mol to do the same to the N₂ molecule.

Bond Length

Another consequence of the presence of multiple bonds between atoms is the difference in the distance between the nuclei of the bonded atoms. Double bonds have shorter distances than single bonds, and triple bonds are shorter than double bonds.

Physical Properties of Covalent Molecules

The covalent bonding model helps predict many of the physical properties of compounds.

First described by Gilbert Lewis, a covalent bond occurs when electrons of different atoms are shared between the two atoms. These cases of electron sharing can be predicted by the octet rule. The octet rule is a chemical rule that generalizes that atoms of low atomic number (< 20) will combine in a way that results in their having 8 electrons in their valence shells. Having 8 valence electrons is favorable for stability and is similar to the electron configuration of the inert noble gases. In a covalent bond, the shared electrons contribute to each atom’s octet and thus enhance the stability of the compound.

The Lewis bonding theory can explain many properties of compounds. For example, the theory predicts the existence of diatomic molecules such as hydrogen, H₂, and the halogens (F₂, Cl₂, Br₂, I₂). A H atom needs one additional electron to fill its valence level, and the halogens need one more electron to fill the octet in their valence levels. Lewis bonding theory states that these atoms will share their valence electrons, effectively allowing each atom to create its own octet.

Several physical properties of molecules/compounds are related to the presence of covalent bonds:

• Covalent bonds between atoms are quite strong, but attractions between molecules/compounds, or intermolecular forces, can be relatively weak. Covalent compounds generally have low boiling and melting points, and are found in all three physical states at room temperature.
• Covalent compounds do not conduct electricity; this is because covalent compounds do not have charged particles capable of transporting electrons.
• Lewis theory also accounts for bond length; the stronger the bond and the more electrons shared, the shorter the bond length is.

However, the Lewis theory of covalent bonding does not account for some observations of compounds in nature. The theory predicts that with more shared electrons, the bond between the two atoms should be stronger. According to the theory, triple bonds are stronger than double bonds, and double bonds are stronger than single bonds. This is true. However, the theory implies that the bond strength of double bonds is twice that of single bonds, which is not true. Therefore, while the covalent bonding model accounts for many physical observations, it does have its limitations.