Mechanisms of Chemical Bonding

Characteristics of Chemical Bonding

The first step to understanding chemical bonding is to define it.

The characteristics of all types of chemical bonding include:

• Atoms interacting with one another to form molecules or compounds.

• Energy needed to form and break chemical bonds.

To be able to understand why chemical bonds form, we need to understand what electronegativity is. An atom's electronegativity refers to its ability to attract electrons to itself. Figure 1 shows the Pauling scale of electronegativity.

Figure 1. Pauling Scale of Electronegativity, Isadora Santos - Vaia Originals.

In a chemical bonding, the difference in electronegativity values between the atoms in a compound will determine the type of chemical bonding formed.

Types of Chemical Bonding

An atom's structure determines the type of bond it can form with another atom. There are three types of chemical bonding:

• Ionic bonding
• Covalent bonding
• Metallic metallic

Ionic Bonding

First, we have ionic bonding.

This type of ionic bonding usually occurs between metals and non-metals when there is a large difference in electronegativity between them. For example, potassium (K) has an electronegativity value of 0.8, whereas chlorine (Cl) has an electronegativity value of 3.0. Since the difference in electronegativity between sodium and chlorine is greater than 1.7, we say that the bonding between potassium and chlorine to form potassium chloride (KCl) is a type of ionic bonding.

To find the difference in electronegativity between two atoms, we can use the equation below:

$$ \Delta \text{EN = (Atom with larger EN value) - (Atom with smaller EN value) } $$

Ionic compounds formed from ionic bonding, tend to be crystalline solids with very high melting points. Ionic compounds also tend to be soluble in water (H₂O) and insoluble in nonpolar solvents. Ionic solids can only conduct electricity in aqueous solutions.

Let's look at the crystal lattice structure of potassium chloride (KCl), shown in figure 2.

Figure 2. Crystal lattice structure of the ionic solid KCl, Isadora Santos - Vaia Originals.

Now, let's look at the mechanism that forms an ionic bond between the sodium (Na) and chlorine (Cl) to form table salt, NaCl (figure 3).

In the periodic table, sodium is in group 1. Elements in this group tend to lose one electron to form a "+1" charged cation. On the other hand, halogens (group 17) such as chlorine tend to gain one electron, forming a "-1" charged anion.

Both elements in group 1 and group 17 want to have eight electrons in their outer shell, the same as a noble gas. For this to happen, sodium needs to lose one electron and go from having 11 electrons to having 10 electrons, just like the noble gas neon (which also has 10 total electrons)!

Similarly, chlorine wants to gain one electron, and have eight electrons in its outer shell, just like the noble gas argon, which contains a total of 18 electrons and 8 in its outer shell.

Therefore, in this chemical bonding between sodium and chlorine, sodium will donate one electron to chlorine, forming an ionic bond as a result of this transfer of electrons!

Another example of ionic bonding is the chemical bonding that occurs between magnesium (Mg) and oxygen (O). Being in group 2, magnesium wants to lose two electrons and form a +2 cation, in other to have the same number of electrons as the noble gas neon. Oxygen wants to gain two electrons to have the same number of electrons as neon and form a -2 cation.

Figure 3. Ionic bonding of NaCl and MgO, Daniela Lin - Vaia Originals.

Covalent Bonding

The second type of ionic bonding we will explore is covalent bonding. Covalent bonding usually occurs between two non-metals.

For covalent bonding to occur, the difference in electronegativity between two atoms needs to be less than 1.7. For instance, if we combine two hydrogen atoms to form a hydrogen molecule, each hydrogen atom will contribute one valence electron to the final molecule, and these two valence electrons will be shared by both atoms!

• The bond will be nonpolar (or pure) covalent if the difference in electronegativity is less than 0.5.

• If the difference in electronegative is greater than 0.5, but still less than 1.7, the bond form is considered a polar covalent bond. In polar covalent bonds, the electrons belong predominantly to one type of atom while being partially shared by the other type.

• If the difference in electronegative is equal to or greater than 1.7, the bond form is considered an ionic bond. In ionic bonds, the electrons belong to the more electronegative atom.

Compounds or molecules formed by covalent bonding tend to be gas, liquids or soft solids at room temperature. The melting and boiling point of covalent compounds are low. In water, they are insoluble, but in nonpolar solvents, they are soluble.

Let's take a look at the mechanism of formation of covalent bonds between two chlorine (Cl) atoms (figure 4). Since chlorine is an atom with an atomic number of 17, it contains 17 protons and 17 electrons.

Since both atoms want to gain one electron to have a full outer shell, they agree to share two valence electrons! In a similar way, two oxygen atoms can share electrons to form an O₂ molecule.

Figure 4. Example of covalent bonding in Cl₂ and O₂, Daniela Lin - Vaia Originals.

Metallic Bonding

Now that we know what ionic and covalent bonding are, let's talk about metallic bonding. This chemical bonding only happens between metals atoms, specifically when the positively charged metal nuclei are attracted to the delocalized valence electrons of another metal atom.

Metallic solids usually have high melting and boiling points, and are insoluble in both water (H₂O) and nonpolar solvents. Metallic solids are lustrous and good conductors of heat and electricity.

The electrons in metallic bonding can move around the positive metal ions, which are arranged in a giant lattice structure. It is thanks to this ability of electrons to freely move around that makes metallic solids so good at conducting electricity!

Figure 5. Example of metallic bonding, Daniela Lin - Vaia Originals.

The Strongest Chemical Bond

Now that we have learned the different types of chemical bonds, we can consider the question of which chemical bond is the strongest. Is it metallic, covalent, or ionic? Well, this depends on bond length.

When dealing with bond length, the longer the bond length, the easier it is to break the bond, whereas the shorter the bond length, the stronger the bond, and the harder it is to break it. Therefore, covalent bonds should be the strongest chemical bond because atoms are closer together and are pulling on one another with more force than ionic bonds!

The figure below shows the bond length and strength of some covalent compounds (figure 6).

Chemical Bond Examples

Let's take a look at some other examples involving chemical bonding. Figure 7 shows the dot-and-cross diagram for calcium chloride (CaCl₂). The bond formed between calcium and the two chlorine atoms is an ionic bond.

In this case, Calcium (Ca) wants to lose two electrons and become a +2 ion, while each chlorine atom wishes to gain one electron and become a "-1" ion. So, calcium donates one electron to each chlorine atom!

Now, in the case of ammonia (NH₃), the nitrogen atom (N) has 5 electrons in its outermost shells. So, it wants to gain 3 electrons to fill its outer shells. So, it will share electrons with the three hydrogen atoms so that they can all have full outer shells. Since the electrons are being shared, this is an example of covalent bonding.

The Chemical Bond Energy Transfer Mechanism

Lastly, let's explore bond energy and the mechanism of energy transfer in chemical bonds The exchange of energy between a chemical reaction and its surroundings is described by a change in enthalpy (ΔH).

The formula to calculate enthalpy change is as follows:

$$ \Delta \text{H = H}_{products} \text{ - H}_{reactants} $$

Where,

• \( \Delta \text { H}_{products} \) is the enthalpy of products
• \( \Delta \text { H}_{reactants} \) is the enthalpy of reactants

The mechanism of bond breaking is considered endothermic. In exothermic reactions, the enthalpy change (ΔH) is positive.

On the other hand, bond making is said to be exothermic. In exothermic processes, the products contain less energy than the reactants as heat is given out to the surroundings. Therefore, exothermic reactions have a negative enthalpy change (ΔH).

Figure 9. Enthalpy profile diagrams for Exothermic vs. Endothermic reactions, Daniela Lin - Vaia Originals.

The formation of ionic bonds is a highly exothermic process because when ions combine to form an ionic solid, large amounts of energy are released to the surroundings! When dealing with ionic bonding, the energy given out when ions of opposite charges come together and form a crystal lattice is called lattice energy (\( \Delta \text{H }_{lattice}^{\Theta} \)). For example, the lattice energy for the formation of the ionic solid NaCl is -787 kJ mol^-1.

• In ionic bonding, the more exothermic the lattice energy, the stronger the ionic bonding in the lattice!

Metallic and covalent bonds are also exothermic since they all involve the formation of bonds! For example, the covalent bond formed between hydrogen and oxygen, has a bond enthalpy (ΔH) of -463 kJ/mol.

Now, I hope that understand the general characteristics of chemical bonding, types of chemical bonding, and the mechanisms involved in chemical bonding!