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Types and Properties of Chemical Bonds

1. Ionic Bond

As the name suggests, ionic bonds are a result of the attraction between ions. Ions are formed when an atom loses or gains an electron. These types of bonds are commonly formed between a metal and a nonmetal .

Examples

• Sodium(Na) and chlorine (Cl) combine to form stable crystals of sodium chloride (NaCl), also known as common salt.
• Magnesium(Mg) and oxygen (O) combine to form magnesium oxide (MgO).
• Potassium(K) and chlorine (Cl) combine to form potassium chloride (KCl)
• Calcium(Ca) and fluorine (F) combine to form calcium fluoride (CaF₂

Ionic Bond Properties

Due to the presence of a strong force of attraction between cations and anions in ionic bonded molecules, the following properties are observed:

1. The ionic bonds are the strongest of all the bonds.
2. The ionic bond has charge separation, and so they are the most reactive of all the bonds in the proper medium.
3. The ionic bonded molecules have high melting and boiling point.
4. The ionic bonded molecules in their aqueous solutions or in the molten state are good conductors of electricity. This is due to the presence of ions which acts as charge carriers

2. Covalent Bond

In the case of a covalent bond, an atom shares one or more pairs of electrons with another atom and forms a bond. This sharing of electrons happens because the atoms must satisfy the octet (noble gas configuration) rule while bonding. Such a type of bonding is common between two nonmetals. The covalent bond is the strongest and most common form of chemical bond in living organisms. Together with the ionic bond, they form the two most important chemical bonds ^].

A covalent bond can be divided into a nonpolar covalent bond and a polar covalent bond. In the case of a nonpolar covalent bond, the electrons are equally shared between the two atoms. On the contrary, in polar covalent bonds, the electrons are unequally distributed between the atoms.

Examples

• Two atoms of iodine(I) combine to form iodine (I₂) gas.
• One atom of carbon(C) combines with two atoms of oxygen (O) to form a double covalent bond in carbon dioxide (CO₂).
• Two atoms of hydrogen(H) combine with one atom of oxygen (O) to form a polar molecule of water (H₂O).
• Boron(B) and three hydrogens (H) combine to form the polar borane (BH₃).

Properties of Covalent Bond

If the normal valence of an atom is not satisfied by sharing a single electron pair between atoms, the atoms may share more than one electron pair between them. Some of the properties of covalent bonds are:

• Covalent bonding does not result in the formation of new electrons. The bond only pairs them.
• They are very powerful chemical bonds that exist between atoms.
• A covalent bond normally contains the energy of about ~80 kilocalories per mole (kcal/mol).
• Covalent bonds rarely break spontaneously after it is formed.
• Covalent bonds are directional where the atoms that are bonded showcase specific orientations relative to one another.
• Most compounds having covalent bonds exhibit relatively low melting points and boiling points.
• Compounds with covalent bonds usually have lower enthalpies of vaporization and fusion.
• Compounds formed by covalent bonding don’t conduct electricity due to the lack of free electrons.
• Covalent compounds are not soluble in water.

Ionic Bond vs Covalent Bond

3. Metallic Bonds

A metallic bond is a chemical bond in which a cloud of free moving valence electrons bonds to positively charged ions in a metal. It is defined as the sharing of free electrons among positively charged metal ions in a lattice.

Metallic bonds have a completely different structure than ionic and covalent bonds. Metallic bonds are formed only between metal atoms. Ionic bonds connect metals to nonmetals and metallic bonds connect a large number of metal atoms.

Metallic bonds can be found in pure metals and alloys, as well as certain metalloids. For example, graphene (a carbon allotrope) has two-dimensional metallic bonding. Other sorts of chemical bonds can be formed between the atoms of metals, even if they are pure. For example- the mercurous ion (Hg2+) can create metal-metal covalent bonds. Pure gallium creates covalent bonds between pairs of atoms that are connected to surrounding pairs via metallic bonds.

Electrons are liberated from the atoms and delocalized throughout the metal, allowing them to travel freely. Interactions between ions and electrons are still there.These interactions produce a binding force that keeps the metallic crystal together. 

A metallic bond can be found in the binding force. The function of metallic bonding is explained in detail below:

1. Outer energy levels (the s and p orbitals) of metal atoms overlap. At least one of the valence electrons in a metallic bond is not shared with a neighbouring atom, nor is it lost in the formation of an ion. Instead, the electrons form an ‘electron sea’ in which valence electrons can freely flow from one atom to the other.
2. The electron sea model explains metallic bonding. Metallic bonding can be seen as a result of a material having more delocalized energy levels than delocalized electrons (electron deficit), which further causes localised unpaired electrons to become delocalized and mobile. Electrons may change energy levels and travel in any direction across a lattice.
3. Metallic cluster formation, in which delocalized electrons flow around localised cores, is another kind of bonding. The development of bonds is highly influenced by environmental factors. For example, under high pressure, hydrogenacts as a metal. As pressure decreases, the transition of bonding takes place from metallic to nonpolar covalent bond.

Properties of Metallic bonds

Metals have various physical and chemical properties. These properties include the capacity to carry electricity and heat, a low ionisation energy, and a low electronegativity. Their physical characteristics include a glossy look, malleability and ductility. Metals have a crystalline structure but can easily be deformed. 

Some of the basic properties of metals are: 

1. Conductivity

Electrons in the electron sea are free to travel and transport charge, most metals are efficient electrical conductors. Conductive non-metals (such as graphite), molten ionic compounds, and aqueous ionic compounds all conduct electricity. Since electrons are free, if electrons from an external source are pushed into a metal wire at one end, the electrons would flow through the wire at the same pace and emerge at the other end.

2. Thermal Conductivity

Metals transmit heat because free electrons in the metals transport energy away from the heat source. The reason behind this is that atom vibrations (phonons) travel like a wave through a solid metal.

3. Malleability

Metals are frequently malleable, or capable of being shaped or hammered into a shape. As the binding force between metals is asymmetric, drawing or sculpting a metal can not fracture it. The electrons in a metal are free to move away from one another and do not push like-charged ions together.

4. Ductility

Metals are ductile, or capable of being pulled into thin wires, because local connections between atoms are quickly broken and rebuilt. Single atoms or whole sheets of atoms can glide past each other and rebuild bonds.

5. Lustre

Metals are usually lustrous or have a metallic sheen. Once a certain minimum thickness is reached, they become opaque. Photons are reflected off the flat surface of the electron sea. The amount of light that can be reflected has an upper frequency limit.

Examples

• Sodium metal
• Aluminumfoil
• Copperwire

5. Hydrogen Bonds

Ionic and covalent bonds are strong bonds that require considerable energy to break. However, not all bonds between elements are ionic or covalent bonds. Weaker bonds can also form.

These are attractions that occur between positive and negative charges that do not require much energy to break. Two weak bonds that occur frequently are hydrogen bonds and van der Waals interactions. These bonds give rise to the unique properties of water and the unique structures of DNA and proteins.

When polar covalent bonds containing a hydrogen atom form, the hydrogen atom in that bond has a slightly positive charge. This is because the shared electron is pulled more strongly toward the other element and away from the hydrogen nucleus.

 Because the hydrogen atom is slightly positive (δ+), it will be attracted to neighboring negative partial charges (δ–). When this happens, a weak interaction occurs between the δ+ charge of the hydrogen atom of one molecule and the δ– charge of the other molecule. This interaction is called a hydrogen bond.

This type of bond is common; for example, the liquid nature of water is caused by the hydrogen bonds between water molecules (Figure 4). Hydrogen bonds give water the unique properties that sustain life. If it were not for hydrogen bonding, water would be a gas rather than a liquid at room temperature.

 Hydrogen bonds form between slightly positive (δ+) and slightly negative (δ–) charges of polar covalent molecules, such as water.

Hydrogen bonds can form between different molecules and they do not always have to include a water molecule. Hydrogen atoms in polar bonds within any molecule can form bonds with other adjacent molecules.

For example, hydrogen bonds hold together two long strands of DNA to give the DNA molecule its characteristic double-stranded structure. Hydrogen bonds are also responsible for some of the three-dimensional structure of proteins.

Van der Waals interactions

Like hydrogen bonds, van der Waals interactions are weak attractions or interactions between molecules. They occur between polar, covalently bound, atoms in different molecules. Some of these weak attractions are caused by temporary partial charges formed when electrons move around a nucleus. These weak interactions between molecules are important in biological systems.

Peptide Bond

Within a protein, multiple amino acids are linked together by peptide bonds, thereby forming a long chain. Peptide bonds are formed by a biochemical reaction that extracts a water molecule as it joins the amino group of one amino acid to the carboxyl group of neighboring amino acids. Aside from peptide bonds, hydrogen bonds, ionic bonds, and disulfide bonds are also common in proteins . Examples include polypeptides like insulin and growth hormone.