Lesson Explainer: Types of Chemical Bonding Chemistry • Second Year of Secondary School

In this explainer, we will learn how to describe the different types of chemical bonding, understand the concept of valency, and represent chemical bonds using Lewis structures.

Atoms tend to be unstable when they are not bonded and significantly more stable when they are grouped together as compounds. Some atoms combine with each other and form simple covalently bonded molecules such as water () or oxygen (). Other atoms group together and form giant lattice structures. This includes both ionic and metallic lattice types. Simple covalently bonded molecules usually contain two or three atoms, and a giant lattice can contain an unimaginably high number of atoms or ions. Atoms can clearly form very different types of compounds, and this is a good reason to understand what causes some atoms to form covalent compounds and other atoms to form giant lattice structures.

The bonding properties of any one atom or ion are primarily determined by its number of valence electrons. Valence electrons are the electrons of an atom or ion that cannot be classed as core electrons. They are electrons that do not make up inner electron shells. They make up outer electron shells. We can better understand valence electrons if we examine the electron shells of some representative group 1 element atoms. The following figure shows the electron shells of three group 1 element atoms. It shows the electron shells of lithium, sodium, and potassium atoms, respectively, from left to right.

We can see that lithium has a total of just three electrons and that potassium has nineteen. Sodium has a total number of electrons that is between these two extremes. The figure shows that the group 1 element atoms have a different number of total electrons, but they all have the same number of valence electrons. It can be said that they all have the same number of valence electrons because they all have one outer shell electron.

Elements usually end up forming a giant metallically bonded lattice if they have a low number of valence electrons. Group 1 element atoms form a metallically bonded lattice because they have a single valence electron. The following figure shows how sodium atoms form a giant metallically bonded lattice. It represents the electrons as blue circles, and the symbol represents a large number.

The sodium atoms are shown to effectively transform into positively charged sodium ions when all of their valence electrons merge together and form one negatively charged sea of electrons. Valence electrons tend to be called delocalized electrons when they are decoupled from atoms in a metallic lattice. This is partly because valence electrons are highly mobile. It is also because valence electrons do not stay close to any one metal cation. Metallic lattice structures tend to be relatively stable because of the strong attractive electrostatic forces between their metal cations and their negatively charged sea of electrons.

Metallic bonding is just one type of chemical bonding. There are at least two other types of chemical bonding that you will have to learn about in this explainer. There is ionic bonding, which is responsible for making ionic compounds, and covalent bonding, which is responsible for making covalent compounds. Ionic compounds are made up of oppositely charged ions. They contain an incredibly high number of positively and negatively charged ions that are arranged as a giant three-dimensional lattice. Each one of the positively charged ions is surrounded by negatively charged ions, and each one of the negatively charged ions is surrounded by positively charged ions. Ionic bonds are the strong electrostatic interactions between positively and negatively charged ions in an ionic lattice. Ionic bonds tend to be strong and difficult to break. It usually takes a lot of energy to overcome the strong bridging electrostatic interactions between the oppositely charged ions of an ionic lattice.

Ionic bonds are usually formed when valence electrons are transferred from metal atoms to nonmetal atoms. The electron-transfer process produces oppositely charged ions that are attracted to each other. The metal atoms turn into positively charged ions as they lose electrons, and the nonmetal atoms turn into negatively charged ions as they gain electrons. These oppositely charged ions are drawn toward each other and they end up forming a three-dimensional lattice. Ionic bonds are the strong electrostatic forces between the oppositely charged ions of a three-dimensional ionic lattice. There is usually very little space between the positively and negatively charged ions of an ionic lattice.

Covalent bonding is another type of chemical bonding. Covalent bonds are formed when one nonmetal atom shares its valence electrons with another nonmetal atom. Covalent compounds are usually very small. They are usually smaller than a single nanometer, and they can contain no more than two or three individual atoms. Covalently bonded compounds are quite interesting structures because they do not contain any positively or negatively charged ions. They contain uncharged or partially charged atoms like hydrogen and oxygen. Covalent compounds are not formed when electrons are transferred between atoms. They are formed when valence electrons are shared between atoms.

The following figure represents three different types of chemical bonds. It is important to understand that the figure does not show any core electrons. It only shows valence electrons. The figure shows that the valence electrons of metals are essentially decoupled from any one metal atom. It also shows that the valence electrons of ionic and covalent compounds are coupled with individual ions or atoms.

The previous paragraphs have shown that there are some important similarities between the different chemical bonding types. It was explained that both metallic and ionic compounds contain ions that are packed very tightly together in the shape of a three-dimensional lattice. The structure of both lattice types is maintained with strong electrostatic forces of attraction between cations and a sea of delocalized electrons or oppositely charged ions.

The previous paragraphs have also shown that there are many important differences between chemical bonding types. It was stated that covalent compounds contain neutrally or partially charged atoms and that ionic compounds contain a combination of positively and negatively charged ions. It was also stated that covalent compounds tend to be incredibly small and that a metallic or ionic lattice tends to be much larger. One of the most important differences between metallic and nonmetallic bonding types is the presence or absence of delocalized electrons. Metals contain a sea of delocalized electrons, but there are no comparable delocalized electrons in covalently or ionically bonded compounds. The following table recapitulates most of the information that has been described in the preceding paragraphs.

The octet rule is an incredibly simple scientific hypothesis that can be used to explain the formation of ionic and covalent compounds. It states that atoms tend to be more stable if they have eight valence electrons and the same electron configuration as a noble gas atom. The following figure shows the electron configuration of the three noble gas atoms that are the basis for the octet rule. The octet rule predicts that atoms usually react so that they end up with an electron configuration that matches one of the noble gas atoms.

Let us consider the reaction of metal sodium with nonmetal chlorine. The reaction is highly exothermic, and it produces a highly stable sodium chloride product. Sodium is a group 1 element, and it has a single valence electron. It can effectively gain the same electron configuration as a neon atom if it loses a single valence electron. It can effectively end up with the same electron configuration as a noble gas atom if its single valence electron is transferred to a nonmetal atom.

Chlorine is a group 17 element, and it has seven valence electrons. It can effectively gain the same electron configuration as an argon atom if it gains a single valence electron. It can effectively end up with the same electron configuration as a noble gas atom if it gains a single valence electron from a metal atom.

Sodium and chlorine atoms end up having the same electron configuration as a noble gas atom if they react with each other. The sodium metal atoms end up transforming into positively charged ions as they lose electrons, and the chlorine atoms end up transforming into negatively charged ions as they gain these electrons. The product sodium and chloride ions have a stable electron configuration because they both have eight valence electrons and the same electron configuration as a noble gas atom. This is shown in the following figure.

Let us now consider representative examples of atoms sharing electrons to effectively gain the same electron configuration as a noble gas atom.

The following figure shows how two chlorine atoms can combine together when they share a single valence electron. Each one of the chlorine atoms effectively gains a valence electron when the chlorine atoms combine together. They both go from having seven valence electrons to having eight valence electrons and the same electron configuration as an argon atom. A single shared pair of electrons is usually called a single covalent bond. We can state that the diatomic chlorine molecule contains two chlorine atoms that are linked together with a single chlorine–chlorine () covalent bond.

The next figure shows how two different types of covalent compounds can be formed from oxygen atoms. It shows how oxygen atoms can combine with each other and form one oxygen molecule. It also shows how two oxygen atoms can combine with one carbon atom and form one carbon dioxide molecule. The oxygen atoms go from having six valence electrons to having eight valence electrons when they either bond with each other or when each two oxygen atoms bond with one carbon atom. You will notice here that the oxygen atoms share two valence electrons when they produce either diatomic oxygen molecules or triatomic carbon dioxide molecules. Two pairs of shared electrons are usually called a double covalent bond. We can state that a diatomic oxygen molecule contains two oxygen atoms that are linked together with a double oxygen–oxygen () covalent bond. Similarly, each carbon dioxide molecule contains two oxygen atoms that are linked to a carbon atom with two double covalent bonds ().

Some atoms tend to share more than two valence electrons because they have five valence electrons or less. Nitrogen is a group 15 nonmetal atom, and it has five valence electrons. It can effectively have eight valence electrons and the same electron configuration as the neon atom if it shares three valence electrons and makes a so-called triple covalent bond. The following diagram shows how two nitrogen atoms can combine together and make one diatomic nitrogen molecule. The molecule is made up of two atoms that each have the same electron configuration as the neon atom. They both have eight valence electrons because they each contribute three valence electrons to the triple nitrogen–nitrogen () covalent bond.

Nitrogen can similarly gain the same electron configuration as neon if it makes multiple single covalent bonds with other atoms. Nitrogen can effectively gain three valence electrons if it makes three single covalent bonds with adjacent hydrogen atoms. The following figure shows how a single nitrogen atom can effectively get eight valence electrons if it makes an ammonia molecule. We can state that an ammonia molecule contains one nitrogen atom that is linked to three hydrogen atoms through three single covalent bonds.

The following table shows how many electrons metals need to lose to have the same electron configuration as a noble gas. It also shows how many electrons nonmetals need to obtain or share to have the same electron configuration as a noble gas.

The table shows that there is a relationship between the group number, the number of valence electrons, and the number of electrons needed to be lost or obtained. Group 2 metals have two outer shell electrons, and they need to lose two valence electrons to have the same electron configuration as a noble gas. Group 16 nonmetals have six valence electrons, and they tend to gain or share two electrons to have eight valence electrons and the same electron configuration as a noble gas.

Lewis structures are simple schematic illustrations that show how valence shell electrons are shared or transferred between atoms when they make covalently or ionically bonded compounds. The following image shows the Lewis structure for a diatomic chlorine () molecule, which has one covalent bond. Each valence shell electron is shown as a single small dot. The red dots represent the valence electrons of one chlorine atom, and the black dots represent the valence shell electrons of the other chlorine atom. The Lewis structure shows how two single chlorine atoms can gain the same electron configuration as a noble gas if they share a single pair of electrons.

The following Lewis structure shows how chlorine atoms get the same electron configuration as argon when they react with sodium metal and produce sodium chloride ().

The image shows that the sodium atoms transfer single valence shell electrons to chlorine atoms, and this generates oppositely charged ions that both have eight valence electrons. We should notice here that the sodium and chloride ions have equal but opposite electric charges. This means that a sodium chloride will not have an overall electric charge. It will be neutrally charged. Ionic compounds always have an overall neutral electric charge.

Some ionic compounds have to contain an unequal number of positively and negatively charged ions to have an overall neutral electric charge. This point can be better understood if we focus on magnesium fluoride. The compound has the chemical formula . It contains two negatively charged fluoride ions for each magnesium ion. The compound can only have zero overall electric charge if there are two fluoride ions for each magnesium ion. The Lewis structure for the formation of magnesium fluoride is shown below. It is clear that the reaction of magnesium and fluorine atoms generates two fluoride ions for each magnesium ion.

The following table shows how positively and negatively charged ions of different charges can combine and make different types of ionic compounds. It is clear that ionic compounds always contain the ratio of positively and negatively charged ions that gives them zero overall electric charge. The insights gained here can be used to predict the chemical formula of any one type of ionic compound.

Example 3: Understanding How to Draw Lewis Structures for Fluoride Ions

Which of the following is the correct Lewis structure for a fluoride ion?

A.

B.

C.

D.

E.

Answer

Fluorine atoms have seven valence shell electrons, and they form fluoride ions when they gain a single outer shell electron. This means that fluoride ions must have eight valence shell electrons and a negative electrostatic charge. This can be seen in diagram E, our correct answer.

It is usually relatively simple to predict how many covalent bonds are formed by an atom if we use the octet rule. The octet rule states that atoms have a more stable electron configuration if they obtain eight valence electrons and the same electron configuration as a noble gas. Atoms tend to form a number of covalent bonds that gives them eight valence electrons and the same electron configuration as a noble gas. Group 17 elements tend to form a single covalent bond because they have seven valence electrons and need one more to have eight. Group 16 elements tend to form two covalent bonds because they have six valence electrons and need two more to have eight. This information is summarized in the following table. It is important to appreciate that a double covalent bond is made up of two covalent bonds and that a triple bond is made up of three. Oxygen atoms can make two single covalent bonds or one double covalent bond to have eight valence electrons.

This explainer has explored different types of chemical bonding. It is important to stress here, however, that it did not comprehensively cover the different ways atoms can transfer or share electrons or the various electrostatic interactions that can arise between adjacent atoms. Other explainers will explore concepts like coordinate covalent bonding and hydrogen bonding to provide a more holistic understanding of chemical bonding and the different electrostatic interactions that can arise between adjacent atoms. Coordinate covalent bonds are a special type of covalent bond. Hydrogen bonds are unusually strong intermolecular interactions that exist between some types of molecules.