Why covalent bonds are formed

To satisfy the Octet Rule, Carbon needs 4 more valence electrons. Since each Oxygen atom has 3 lone pairs of electrons, they can each share 1 pair of electrons with Carbon; as a result, filling Carbon's outer valence shell Satisfying the Octet Rule.

A Triple bond is when three pairs of electrons are shared between two atoms in a molecule. It is the least stable out of the three general types of covalent bonds. It is very vulnerable to electron thieves!

Below is a Lewis dot structure of Acetylene demonstrating a triple bond. As you can see from the picture below, Acetylene has a total of 2 Carbon atoms and 2 Hydrogen atoms. Each Hydrogen atom has 1 valence electron whereas each Carbon atom has 4 valence electrons. Each Carbon needs 4 more electrons and each Hydrogen needs 1 more electron. Hydrogen shares its only electron with Carbon to get a full valence shell.

Now Carbon has 5 electrons. Because each Carbon atom has 5 electrons single bond and 3 unpaired electrons--the two Carbons can share their unpaired electrons, forming a triple bond. Now all the atoms are happy with their full outer valence shell.

A Polar Covalent Bond is created when the shared electrons between atoms are not equally shared. This occurs when one atom has a higher electronegativity than the atom it is sharing with. The atom with the higher electronegativity will have a stronger pull for electrons Similiar to a Tug-O-War game, whoever is stronger usually wins. As a result, the shared electrons will be closer to the atom with the higher electronegativity, making it unequally shared.

A polar covalent bond will result in the molecule having a slightly positive side the side containing the atom with a lower electronegativity and a slightly negative side containing the atom with the higher electronegativity because the shared electrons will be displaced toward the atom with the higher electronegativity. As a result of polar covalent bonds, the covalent compound that forms will have an electrostatic potential. This potential will make the resulting molecule slightly polar, allowing it to form weak bonds with other polar molecules.

One example of molecules forming weak bonds with each other as a result of an unbalanced electrostatic potential is hydrogen bonding , where a hydrogen atom will interact with an electronegative hydrogen, fluorine, or oxygen atom from another molecule or chemical group. As you can see from the picture above, Oxygen is the big buff creature with the tattoo of "O" on its arm. The little bunny represents a Hydrogen atom. The blue and red bow tied in the middle of the rope, pulled by the two creatures represents--the shared pair of electrons--a single bond.

Because the Hydrogen atom is weaker, the shared pair of electrons will be pulled closer to the Oxygen atom. This table is just a general guide, however, with many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1.

Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2. The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic.

Some compounds contain both covalent and ionic bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. Bond polarities play an important role in determining the structure of proteins.

Using the electronegativity values in Figure 7. The polarity of these bonds increases as the absolute value of the electronegativity difference increases. Table 1 shows these bonds in order of increasing polarity. Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other.

The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent.

When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized as ionic. NaCl consists of discrete ions arranged in a crystal lattice, not covalently bonded molecules.

Skip to main content. Module 7: Chemical Bonding and Molecular Geometry. Search for:. Covalent Bonding Learning Outcomes Describe the formation of covalent bonds Define electronegativity and assess the polarity of covalent bonds. Portrait of a Chemist: Linus Pauling Linus Pauling, shown in Figure 4, is the only person to have received two unshared individual Nobel Prizes: one for chemistry in for his work on the nature of chemical bonds and one for peace in for his opposition to weapons of mass destruction.

Example 1: Electronegativity and Bond Polarity Bond polarities play an important role in determining the structure of proteins. Table 1. Key Concepts and Summary Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. After bonding, the chlorine atom is now in contact with eight electrons in its outer shell, so it is stable.

The hydrogen atom is now in contact with two electrons in its outer shell, so it is also stable. Both nuclei are strongly attracted to the shared pair of electrons in the covalent bond, so covalent bonds are very strong and require a lot of energy to break.

Atoms may form multiple covalent bonds — they share not only one pair of electrons, but two or more pairs. Atoms of different elements will form either one, two, three or four covalent bonds with other atoms. Generally, orbital shapes are drawn to describe the region in space in which electrons are likely to be found.

Atomic orbitals : The shapes of the first five atomic orbitals are shown in order: 1s, 2s, and the three 2p orbitals. 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, H 2 , 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 bonds, comprised of sigma and pi bonds, increase the stability and restrict the geometry of a compound. 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.

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 , depending on how many p orbitals are required to explain the observed bonding.

Hybridized orbitals : A schematic of the resulting orientation in space of sp 3 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. 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.

The simplest example of an organic compound with a double bond is ethylene, or ethene, C 2 H 4. 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.