Types of Chemical Bonds

Various states, such as plasma, Bose-Einstein condensates, fermionic condensates, and quark-gluon plasma, are possible regardless of the common times of solid, liquid, and gas – for example, water exists as ice, liquid water, and vaporous steam – but various states, such as plasma, Bose-Einstein condensates, fermionic condensates, and quark-gluon plasma, are also possible. Apart from that, it is divided into pure substances and mixtures.

Chemical Bond

Chemical bonding is the process of uniting two or more atoms by the redistribution of electrons, resulting in each atom achieving a stable electronic state.

To achieve security, each of the atoms completes their duplet or octet by obtaining the nearest tolerable gas electronic arrangement. It’s usually performed by the formation of composite connections between particles.

In three ways, a particle can form chemical bonds:

1. At least one electron is lost to another atom.
2. Obtaining at least one electron from a different atom.
3. By transferring one electron to another atom.

Types of Chemical Bonds

Chemical bonds are classified into four groups:

• Ionic or electrovalent bond
• Covalent bond
• Hydrogen bond
• Polar covalent bond

Ionic Bond

The octet rule states that an atom is most stable when its valence shell has eight electrons. Atoms in a solid-state lose, acquire, or share electrons in their valence shell to achieve stability. A cation is formed when an atom loses one or more valence electrons to become a positively charged ion, whereas an anion is formed when an atom receives electrons and becomes a negatively charged ion. 

An electrical force of attraction known as the Ionic Bond emerges when electrons are completely transferred from the positively charged cation to the negatively charged anion.

The octet rule is satisfied when valence electrons are exchanged, allowing ions to reach electronic configurations of neighbouring noble gases. A positive superscript charge (+) to the right of the atom denotes a cation. A negative superscript charge (–) to the right of the atom represents an anion.

For example:

• When a sodium atom loses one electron, it gains one more proton than it loses an electron, giving it a positive (+1) charge overall. The chemical symbol for sodium-ion is Na⁺¹ or simply Na⁺.
• When a chlorine atom gets an additional electron, it forms the chloride ion Cl^–. Because of the octet rule, these ionic species are more stable than the atom.

Formation of an Ionic Bond

An ionic or electrovalent bond is a chemical bond produced between two atoms by the transfer of one or more electrons from an electropositive or metallic element to an electronegative or non-metallic element. The electrical arrangement of the sodium atom is 2,8,1, as we know. In its outermost shell, it just possesses one electron. It gets the inert gas electrical structure of Neon by donating this electron (2,8).

The electrical arrangement of the chlorine atom, on the other hand, is 2,8,7. It simply needs one more electron to complete its octet and achieve the inert configuration of Argon (2,8,8).

To remove an electron from its outermost shell and produce the positively charged sodium ion Na⁺, the sodium atom requires an amount of energy equal to its ionisation energy of roughly 496kJ/mol. This is referred to as an endergonic process since it consumes energy. The chlorine atom, on the other hand, which is missing one electron, takes an electron and releases energy equivalent to its electron affinity, i.e. To produce a chloride ion, 349kJ/mol is required. The Exergonic process gets its name from the fact that energy is released during the process.

As a result, the oppositely charged ions create electrostatic forces of attraction, which are powerful forces of attraction. These forces drive the ion closer together, resulting in the formation of an ionic bond. Electrostatic forces are hence the foundation of an ionic connection.

Electrovalency– The amount of atoms lost or acquired by one atom is referred to as electrovalency. In NaCl, the electrovalency of sodium and chlorine is one. As a result, they are monovalent. Calcium provides two valence electrons to generate calcium ions (2,8,8), and oxygen obtains two electrons to form oxide ions in the creation of Calcium oxide (CaO) (2,8). Calcium and oxygen have the same electrovalency, i.e. they are divalent. The oppositely charged ions are formed once the electrons are transferred. The development of an ionic link between sodium and fluorine atoms is seen here.

Conditions for the formation of Ionic Bond

• The number of valence electrons present in the atoms involved in bonding-

Atoms with one, two, or three valence electrons in groups I(A), II(A), or 13(A) of the modern periodic table prefer to lose electrons and form positively charged species known as cations. Atoms with 5,6, or 7 valence electrons in groups 15,16, and 17 of the current periodic table, on the other hand, are more likely to take electrons and form negatively charged species known as anions. Thus, potassium, which is in group I(A) of the current periodic table and has one electron in its outermost shell, is suitable for creating an ionic connection with chlorine, which is in group 17 and has seven electrons in its outermost shell. As a result, potassium chloride is classified as an ionic substance.

• The low ionization energy of the metal –

The ionisation energy is the minimal amount of energy necessary to remove an electron from a neutral gaseous atom’s outermost shell. Consider how sodium ion (Na⁺) is formed from sodium atom. Sodium’s ionisation energy is around 500 kJ/mol, which is relatively low. It’s quite easy for it to shed electrons and become a sodium ion. This ion may also form ionic bonds with other anions like Cl^– and Br^–. As a result, the creation of ionic bonds is favoured when the metal’s ionisation energy is low.

• The electron affinity of a non-metal – 

The energy generated when an electron is introduced to a neutral isolated gaseous atom is defined as electron affinity. Consider a fluorine atom, which contains seven electrons in its outermost orbit. It receives one electron without hesitation to complete its octet structure. Per mol, this reaction releases roughly 328kJ of energy. The resulting fluoride ion has lower energy than the fluorine atom. We all know that reduced energy equals more stability. As a result, the fluoride ion is more stable than fluorine. As a result, a non-greater metal’s electron affinity favours the creation of an anion, resulting in a stable ionic combination.

• The lattice energy of the ionic compound –

The lattice energy of the ionic compounds is a key feature that influences their stability. The energy generated when one gram mole of a crystal is created from its gaseous ions is known as lattice energy. Electrostatic forces of attraction hold the component ions of an ionic compound together in any crystal.

•  The difference in electronegativity between two atoms –

The capacity of an atom to pull electrons towards itself during bond formation is measured by its electronegativity. When the electronegativity values of the atoms differ by 1.7 or more, the creation of an ionic connection is more likely. Consider the sodium chloride (NaCl) and hydrogen chloride (HCl) molecules.

Electronegativity differential =3.0(Cl)–0.9(Na)=2.1 for NaCl.

In sodium chloride, this aids in the establishment of a stable ionic connection.

Electronegativity differential =3.0(Cl)–2.1(H)=0.9 for HCl.

The link between hydrogen and chlorine is covalent due to the lesser electronegativity difference.

Writing Formula of an Ionic Compound

The following two requirements must be met in order to derive the chemical formulae of ionic compounds:

1. For optimum stability, the cation and anion should obey the octet rule.
2. Ions should unite in such a way that their charges balance out and the ionic composition as a whole is neutral.

The number of electrons provided or received is represented by the charges on the anion and cation. In ionic bonding, the compound’s net charge must be zero.

Properties of an Ionic Bond

1. The cation is always formed by a metal, while the anion is always formed by a non-metal.
2. At room temperature, most ionic compounds are crystalline solids – the component ions of the ionic compound have a strong attraction to one another and are organised in a three-dimensional structure. This arrangement gives the crystal its distinctive geometrical form.
3. Ionic compounds have high melting points because the ions in them are kept together by extremely strong attraction forces. Ionic compounds have relatively high melting points, which means that breaking the link between them requires a lot of heat energy.
4. Electrical conductivity – Ionic compounds are typically non-conductors of electricity in the solid-state. The electrical force of attraction between the ions is broken when heated over their melting point, and the ions are free to travel. Electricity can now travel across these free ions.
5. When an ionic chemical such as sodium chloride is introduced to water, the water molecule’s negative end attracts the cations and pulls them out of the crystal. Similarly, the positive end of the water molecule attracts the anions, causing the chemical to dissolve in water. Ionic chemicals are therefore soluble in polar solvents such as water.
6. Ionic compounds are fragile, they fracture into fragments when an external force is applied to their crystals. This occurs because the Na⁺ and Cl^– ions in sodium chloride crystals are lined up against one other in a lattice with a strong electrostatic attraction.

When an external force is applied, the ions’ alignment changes in such a manner that similar charges are brought closer together. The ions migrate apart as a result of the high electrostatic repulsion. The form of the crystals fractures when the ions smash into bits.

Covalent Bond

• In simple terms, a covalent bond is the exchanging of electrons between particles to achieve the honourable gas configuration of individual iotas.
• The electrical strength of curiosity holds the iotas in a covalent bond together. This power is located in the midst of the reinforced particles’ strongly charged cores and the oppositely charged electrons they share.
• The holding pair of electrons are the electrons that unite iotas in a covalent bond. The organisation of a distinct grouping of particles termed an atom—the smallest component of a compound that holds the synthetic character of that compound—is made possible by these holding pairs of electrons.
• In the occasional table, this type of holding occurs between two particles of a similar component or between components that are close to one another. This holding occurs mostly between non-metals, although it may also be observed between non-metals and metals.

Types of Covalent Bonds

The covalent bond can be categorised into the following categories based on the number of shared electron pairs:

1. Single Covalent Bond
2. Double Covalent Bond
3. Triple Covalent Bond

• Single Bonds

When just one pair of electrons is shared between the two participating atoms, a single bond is established. One dash is used to symbolize it (-). Although it has a lower density and is weaker than double and triple bonds, this type of covalent bond is the most stable.

For Example, One Hydrogen atom has one valence electron and one Chlorine atom has seven valence electrons in the HCL molecule. By sharing one electron, a single bond is established between hydrogen and chlorine in this situation.

• Double Bonds

When two pairs of electrons are shared between the two participating atoms, a double bond is established. Two dashes (=) are used to indicate it. Double covalent bonds are substantially more powerful than single covalent bonds, but they are also less stable.

Example: One carbon atom has six valence electrons and two oxygen atoms have four valence electrons in a carbon dioxide molecule.

Carbon shares two of its valence electrons with one oxygen atom and two with another oxygen atom to complete its octet. CO2 has two double bonds because each oxygen atom shares its two electrons with carbon.

• Triple Bond

When the two participating atoms share three pairs of electrons, a triple bond is established. The least stable forms of covalent bonds are triple covalent bonds, which are indicated by three dashes (≡).

For Example: Each nitrogen atom with five valence electrons contributes three electrons to form three electron pairs for partaking in the construction of a nitrogen molecule. The two nitrogen atoms create a triple bond as a result.

• Polar Covalent Bond: Because the electronegativity of the combining atoms differs, this sort of covalent bond develops when electrons are shared unequally. The attraction for electrons will be larger for more electronegative atoms. Between the atoms, the electronegative difference is more than zero but less than 2.0. As a result, that atom’s shared pair of electrons will be closer.

Example, An uneven electric potential causes molecules to form hydrogen bonds. The hydrogen atom interacts with the electronegative fluorine, hydrogen, or oxygen in this situation.

• Nonpolar Covalent Bond: When atoms share an equal number of electrons, this sort of covalent connection is produced. The difference in electronegativity between two atoms is zero. It happens when the atoms joining have a comparable electron affinity (diatomic elements).

Example, Gas molecules such as hydrogen and nitrogen include nonpolar covalent bonds.

Polarization of Covalent Bonds: The electron cloud is always closer to the more electronegative of the two atoms participating in the sigma bond in sigma bonds between two distinct atoms. As a result, the bond develops a permanent dipole, and the covalent bond is said to be polarised.

Properties of Covalent Bond

• Because of the modest intermolecular forces of attraction, covalent compounds have low boiling and melting points. At ambient temperature, these chemicals exist in all three physical states. While covalent interactions between atoms are fairly strong, intermolecular forces or attractions between molecules/compounds are comparatively moderate. When a smaller quantity of energy is applied to covalent bonds, molecules can split from one another. As a result, these chemicals are extremely volatile.
• The enthalpy of fusion is the amount of energy required to melt one mole of a solid material under constant pressure. The enthalpy of vaporisation is the amount of energy necessary to vaporise one mole of a liquid at constant pressure. It takes just 1 to 10 times the amount of heat to alter the phase of a molecular covalent compound as it does to change the phase of an ionic molecule.
• Weak intermolecular forces of attraction characterise covalent compounds, causing them to adopt the forms of gases, liquids, and soft solids. There are exceptions, as with many features, particularly when molecular compounds take on crystalline forms.
• The majority of combustible compounds are made up of hydrogen and carbon atoms. In the presence of oxygen, these chemicals may easily undergo combustion processes, producing carbon dioxide and water. Because carbon and hydrogen have similar electronegativities, they are found in many molecular compounds together.
• In an aqueous solution, ions are required for the passage of electricity. When molecular chemicals are mixed with water, they dissolve into molecules rather than ions. Because there are no free mobile ions to conduct electricity, when dissolved in water, they normally do not conduct electricity very effectively.
• A polar solvent, such as water, dissolves polar covalent compounds effectively. Sugar and ethanol are two examples of molecular molecules that dissolve easily in water. Nonpolar covalent compounds, on the other hand, do not dissolve readily in water, such as water and oil. These molecules cannot be hydrated by water.
• Weak intermolecular forces of attraction hold both covalent bonds and organic molecules in organic solvents together. The covalent molecules in covalent compounds are easily miscible with the organic molecules in organic solvents because they have the same sort of weak intermolecular forces of attraction. As a result, most covalent compounds are soluble in organic solvents.
• The bond length is also taken into consideration by Lewis’ theory; the stronger the connection and the more electrons shared, the shorter the bond length.

Difference between Covalent and Ionic Bond

Parameters	Covalent Bond	Ionic Bond
What is it?	It is a kind of chemical bonding between two non-metallic atoms characterised by the sharing of electron pairs between atoms as well as other covalent connections.	It is a sort of connection generated by the strong electrostatic force of attraction between oppositely charged ions in a chemical molecule, also known as an electrovalent bond.
Occurs between	Two non-metals or a non-metal and a metalloid	One metal and one non-metal
Nature of component particles	It is made up of discrete electrically neutral molecules.	Cations and anions are oppositely charged particles that make up ionic compounds.
State at room temperature	Gases, liquids, or low melting solids	Crystalline solids
Polarity	Low	High
Solubility	In polar fluids, polar covalent molecules dissolve. HCl in water, for example. Only organic liquids are soluble in non-polar covalent compounds.	Ionic chemicals are only soluble in polar solvents like water because they are polar. Inorganic liquids are insoluble.
Formation	They are created when two non-metals with identical electronegativities come together. Neither atom has enough strength to pull electrons from the other. For stability, they share electrons from the outermost orbital with their neighbours.	They are created when a metal(+ion) and a non-metal interact (-ve ion). Metals are stronger than non-metals, and non-metals may easily get electrons from metals. These two ions are attracted to one another and create an ionic connection.
Conductivity	Ions do not dissociate from non-polar covalent molecules. As a result, don’t conduct electricity. In water, polar covalent compounds quickly dissociate and act as good electrical conductors. HCl in water, for example.	Non-polar covalent compounds do not break down into ions. As a result, you will not be able to conduct electricity. Polar covalent compounds break down readily in water and behave as good electrical conductors. For instance, HCl in water.
Melting Point	Low	High
Boiling Point	Low	High
Example	Methane, Hydrochloric acid	Sodium Chloride, Sulphuric acid

Sample Questions

Question 1: What is Chemical Bond?

Answer:

Chemical bonding is the process of redistributing electrons between two or more atoms so that each atom achieves a stable electronic state.

Question 2: What are the types of chemical bonds?

Answer:

Covalent, hydrogen bonds, van der Waals contacts, and ionic or electrovalent bonds are the four chemical bonds found in chemistry.

Question 3: Which chemical bonds are the strongest?

Answer:

The covalent link is the most powerful chemical bond. Due to the mutual sharing of electrons, they developed between two atoms. Water is a classic example of a covalent bond because both hydrogen and oxygen atoms exchange electrons.

Question 4: What causes a diamond’s melting point to be so high?

Answer:

Diamond is a carbon allotrope. Each of the carbon atoms in a diamond is covalently bound to four other carbon atoms. As a result, a massive covalent structure is formed. As a result, diamond is very hard and has a very high melting point.

Question 5: What are polar covalent bonds?

Answer:

When electrons are shared unequally, this type of covalent bond forms because the electronegativity of the joining atoms vary. For more electronegative atoms, the pull for electrons will be stronger. The electronegative difference between the atoms is more than zero but less than 2.0. As a result, the shared pair of electrons of that atom will be closer together.