8.1: Chemical Bond Formation: An Introduction

Learning Objectives 

• Differentiate between ionic, covalent, and metallic bonding 

Types of Chemical Bonds

Chemical reactions result when the electrons of various atoms rearrange themselves so that there is a net attractive force between atoms called chemical bonds. These chemical bonds come in three types: ionic, covalent, and metallic. 

 

Ionic Bonding: Electron Transfer 

In  Chapter 2, we discussed how electrons are transferred from a metal to a nonmetal so that both atoms have an octet in their valance shell.  The cation and anion produced from this process are held together by electrostatic forces to form an ionic bond. For example, the transfer of two electrons to two bromine atoms results in the formation of the ionic compound, magnesium bromide. 

 

 

Covalent Bonding: Electron Sharing*

Covalent bond occurs when atoms(usually nonmetals) share  their valence electrons with each other so they each have a full octet (or duet in the case of hydrogen).  

This concept can be illustrated by using two hydrogen atoms, each of which has a single electron in its valence shell. (For small atoms such as hydrogen atoms, the valence shell will be the first shell, which holds only two electrons.) We can represent the two individual hydrogen atoms as follows:

In contrast, when two hydrogen atoms get close enough together to share their electrons, they can be represented as follows:

By sharing their valence electrons, both hydrogen atoms now have two electrons in their respective valence shells. Because each valence shell is now filled, this arrangement is more stable than when the two atoms are separate. The sharing of electrons between atoms is called a covalent bond, and the two electrons that join atoms in a covalent bond are called a bonding pair of electrons. A discrete group of atoms connected by covalent bonds is called a molecule—the smallest part of a compound that retains the chemical identity of that compound.

Chemists frequently use Lewis diagrams to represent covalent bonding in molecular substances. For example, the Lewis diagrams of two separate hydrogen atoms are as follows:

The Lewis diagram of two hydrogen atoms sharing electrons looks like this:

This depiction of molecules is simplified further by using a dash to represent a covalent bond. The hydrogen molecule is then represented as follows:

Remember that the dash, also referred to as a single bond, represents a pair of electrons.

 

Metallic Bonding 

 

A third major type of chemical bonding is metallic bonding. Whereas ionic bonds join metals to non-metals and covalent bonds join non-metals to each other, metallic bonding joins a bulk of metal atoms. A metallic substance may be a pure element (e.g. aluminum foil, copper wires), or it may be a mixture of two or more elements in an alloy (e.g. brass instruments, "white gold" jewelry). Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms. Even a soft metal like sodium (melting point 97.8°C) melts at a considerably higher temperature than the element (neon) which precedes it in the Periodic Table. However, unlike ionic compounds, metals are usually malleable rather than brittle, suggesting that they do not form a rigid lattice structure of oppositely charged ions; neither do metals form bonded molecules like covalent compounds, however. A different model of bonding is necessary to explain the properties of metallic substances. In the 1900's, Paul Drüde came up with the "sea of electrons" metallic bonding theory by modeling metals as a mixture of atomic cores (atomic cores = positive nuclei + inner shell of electrons) and valence electrons.

The Electron Sea Model**

Consider sodium metal as an example. Sodium has the electronic structure 1s²2s²2p⁶3s¹. When sodium atoms come together, the electron in the 3s atomic orbital of one sodium atom can share space with the corresponding electron on a neighboring atom to form a bond - in much the same sort of way that a covalent bond is formed. The difference, however, is that each sodium atom is being touched by eight other sodium atoms - and the sharing occurs between the central atom and the 3s orbitals on all of the eight other atoms. Each of these eight is in turn being touched by eight sodium atoms, which in turn are touched by eight atoms - and so on and so on, until you have taken in all the atoms in that lump of sodium. All of the 3s electrons on all of the atoms are shared in nondirectional bonds which extend over the whole piece of metal. The electrons can move freely within the lump of metal, and so each electron becomes detached from its parent atom. The electrons are said to be delocalized. The metal is held together by the strong forces of attraction between the positive nuclei and the delocalized electrons (Figure \(\PageIndex{1}\)).

Figure \(\PageIndex{1}\): Metallic Bonding: The Electron Sea Model: Positive atomic nuclei (orange circles) surrounded by a sea of delocalized electrons (yellow circles).

This is sometimes described as "an array of positive ions in a sea of electrons". If you are going to use this view, beware! Is a metal made up of atoms or ions? It is made of atoms. Each positive center in the diagram represents all the rest of the atom apart from the outer electron, but that electron has not been lost - it may no longer have an attachment to a particular atom, but it's still there in the structure. Sodium metal is therefore written as \(\ce{Na}\), not \(\ce{Na^+}\).

 

Determining the Type of Bond

 

As we delve deeper in bonding, you will soon discover that the difference between ionic, covalent, and metallic bonds is more of a spectrum rather than distinct borders clearly separating the three types of bonding. To distinguish these we need discuss the concept of electronegativity (\(\chi\)), which is a measure of how much an electron is attracted to an atom in a bond, and is covered in section 8.7 of this chapter. In that section we will see that Fluorine is the most electronegative element with a value of 4, while Francium and Cesium are the least electronegative with a value of 0.7.

A "rule of thumb" is that ionic bonds occur when \(\Delta \chi)\), the difference in electronegativity,  is greater than 2,  and either covalent or metallic bonding occurs when it is less than 0.5, with Metallic bonding occurring when the average electronegativity is less that 2 and covalent when the average is greater than two. This can be explained by the van Arkel-Ketelaar diagram.

Fig.8.1.1: van Arkel-Ketelaar diagram showing how electronegativity can be used to distinguish the three types of chemical bonds. The three corners of the diagram represent the extremes

• Bottom Left: Metallic bonds occur between atoms of low electronegativity
• Bottom right: Covalent bonds occur between atoms of high electronegativity
• Top middle: Ionic bonds occur between an atom of high and an atom of low electronegativity.

It needs to be understood that real bonds can be a mixture of these, and for example, many covalent bonds are polar covalent, and we will cover this in section 8.7.

Video 8.1.1 (3:53 Youtube uploaded by Richard Thornley) goes over the van Arkel-Ketelaar diagram in more detail.

Most of this chapter deals with covalent bonds, but it is important to realize that there are also ionic and metallic. In all of these cases, we consider the bonds to occur through interactions of the valence shell electrons, and that core electrons are held too tightly to the parent nucleus to be involved with bonding

 

Contributors 

• *Original Source 
• **Original Source 
• Modified by Ronia Kattoum (UA of Little Rock)