Structure of Atom Class 11 PPT – Class 10, Class 9

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Structure of Atom Class 11 PPT: Short Overview

The basic building block of all matter is called an atom. Atoms are a collection of various subatomic particles containing negatively charged electrons, positively charged protons and neutral particles called neutrons.

Each element has its own unique number of protons, neutrons and electrons. Both protons and neutrons have mass, whereas the mass of electrons is negligible.

Protons and neutrons exist at the centre of the atom in the nucleus.

Electrons move around the nucleus, and are arranged in shells at increasing distances from the nucleus. These shells represent different energy levels, the outermost shell being the highest energy level.

The number of protons that an atom has in its nucleus is called the atomic number. The total number of protons and neutrons in the nucleus of an atom is known as the mass number.

For example, a carbon atom containing six protons and six neutrons has a mass number of 12.

Elements are substances containing atoms of one type only, e.g. O2, N2 and Cl2. Compounds are substances formed when atoms of two or more elements join together, e.g. NaCl, H2O and HCl.

Although 109 elements exist naturally, some of them are extremely rare (check out the periodic table).

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Structure of Atom Class 11 [With PPT]: Orbitals and Electronic Configurations

It is important to understand the location of electrons, as it is the arrangement of the electrons that creates the bonds between the structure of atoms, and chemical reactions are just that to form new bonds.

Electrons are involved in the chemical bonding and reactions of an atom. Electrons are said to occupy orbitals in an atom. An orbital is a region of space that can hold two electrons.

The Role of Orbitals in the Structuring of the Atom

Electrons do not move freely in the space around the nucleus but are confined to regions of space called shells. Each shell can contain up to 2n 2 electrons, where n is the number of the shell. Each shell contains subshells known as atomic orbitals.

The first shell contains a single orbital known as the 1s orbital. The second shell contains one 2s and three 2p orbitals. These three 2p orbitals are designated as 2px, 2py and 2pz.

The third shell contains one 3s orbital, three 3p orbitals and five 3d orbitals. Thus, the first shell can hold only two electrons, the second shell eight electrons and the third shell up to 18 electrons, and so on.

As the number of electrons goes up, the shell numbers also increase. Therefore, electron shells are identified by the principal quantum number, n ¼ 1, 2, 3 and so on.

The Role of Electronic Configurations in the Structuring of the Atom

Electronic Configurations plays a mojor role in the Structuring of the Atom in Chemistry.

The electronic configuration of an atom describes the number of electrons that an atom possesses, and the orbitals in which these electrons are placed. The arrangements of electrons in orbitals, subshells and shells are called electronic configurations.

Electronic configurations can be represented by using noble gas symbols to show some of the inner electrons, or by using Lewis structures in which the valence electrons are represented by dots.

Valence is the number of electrons an atom must lose or gain to attain the nearest noble gas or inert gas electronic configuration. Electrons in the outer shells that are not filled are called valence electrons.

The ground-state electronic configuration is the lowest energy, and the excited-state electronic configuration is the highest energy orbital. If energy is applied to an atom in the ground state, one or more electrons can jump into a higher energy orbital.

Thus, it takes a greater energy to remove an electron from the first shell of an atom than from any other shells.

Examples: How Orbitals and Electronic Configurations Shape the Structure of the Atom

For example, after looking at the structure of the sodium atom has electronic configuration of two, eight and one.

Therefore, to attain the stable configuration, the Na atom must lose one electron from its outermost shell and become the nearest noble gas configuration, i.e. the configuration of neon, which has the electronic configuration of two and eight.

Thus, sodium has a valence of 1. Since all other elements of Group IA in the periodic table have one electron in their outer shells, it can be said that Group IA elements have a valence of 1.

At the far end on the right hand side of the periodic table, let us take another example, chlorine, which has the electronic configuration of two, eight and seven, and the nearest noble gas is argon, which has the electronic configuration of two, eight and eight.

To attain the argon electronic configuration chlorine must gain one electron. Therefore, chlorine has a valence of 1. Since all other elements of Group 7A in the periodic table have seven electrons in their outermost shells and they can gain one electron, we can say that the Group 7A elements have a valence of 1.

Each atom has an infinite number of possible electronic configurations. We are here only concerned with the ground-state electronic configuration, which has the lowest energy.

The ground-state electronic configuration of an atom plays an important role in determining the structure and properties of the atom.

The ground-state electronic configuration of an atom can be determined by the following three principles.

Understanding The Aufbau Principle: A Key to Mastering the Structure of Atom in Class 11

The Aufbau principle states that the orbitals fill in order of increasing energy, from lowest to highest. Because a 1s orbital is closer to the nucleus it is lower in energy than a 2s orbital, which is lower in energy than a 3s orbital.

Understanding The Pauli Exclusion principle: A Key to Mastering the Structure of Atom in Class 11

The Pauli exclusion principle states that no more than two electrons can occupy each orbital, and if two electrons are present, their spins must be paired.

For example, the two electrons of a helium atom must occupy the 1s orbital in opposite spins.

Understanding Hund’s Rule: A Key to Mastering the Structure of Atom in Class 11

Hund’s rule explains that when degenerate orbitals (orbitals that have same energy) are present but not enough electrons are available to fill all the shell completely, then a single electron will occupy an empty orbital first before it will pair up with another electron.

This is understandable, as it takes energy to pair up electrons.

Therefore, the six electrons in the carbon atom are filled as follows: the first four electrons will go to the 1s and 2s orbitals, a fifth electron goes to the 2px, the sixth electron to the 2py orbital and the 2pz orbital will remain empty.

The ground-state electronic configurations for elements 1–18 are listed below (electrons are listed by symbol, atomic number and ground-state electronic configuration).

Let us see how we can write the ground-state electronic configurations for oxygen, chlorine, nitrogen, sulphur and carbon showing the occupancy of each p orbital.

Oxygen has the atomic number 8, and the ground-state electronic configuration for oxygen can be written as 1s ² 2s ² 2px ² 2py ¹ 2pz ¹ . Similarly, we can write the others as follows:

• Chlorine (atomic number 17): 1s ² 2s ² 2px ² 2py ² 2pz ² 3s ² 3px ² 3py ² 3pz ¹
• Nitrogen (atomic number 7): 1s ² 2s ² 2px ¹ 2py ¹ 2pz ¹
• Sulphur (atomic number 16): 1s ² 2s ² 2px ² 2py ² 2pz ² 3s ² 3px ² 3py ¹ 3pz ¹
• Carbon (atomic number 6): 1s ² 2s ² 2px ¹ 2py ¹ 2pz ⁰

Chemical Bonding Theories for Formation of Structure of Atom

Atoms form bonds in order to obtain a stable electronic configuration, i.e. the electronic configuration of the nearest noble gas.

All noble gases are inert, because their atoms have a stable electronic configuration in which they have eight electrons in the outer shell except helium (two electrons).

Therefore, they cannot donate or gain electrons. One of the driving forces behind the bonding in an atom is to obtain a stable valence electron configuration. A filled shell is also known as a noble gas configuration. Electrons in filled shells are called core electrons.

The core electrons do not participate in chemical bonding. Electrons in shells that are not completely filled are called valence electrons, also known as outer-shell electrons, and the energy level in which they are found is also known as the valence shell.

Carbon, for example, with the ground-state electronic configuration 1s 2 2s 2 2p 2 , has four outer-shell electrons. We generally use the Lewis structure to represent the outermost electrons of an atom.

Lewis Structure of Atom for Class 9, 10 and 11

Lets learn the role of Lewis Dot structure in structuring of an atom.

Lewis structures of atom provide information about what atoms are bonded to each other, and the total electron pairs involved.

According to the Lewis theory, an atom will give up, accept or share electrons in order to achieve a filled outer shell that contains eight electrons.

The Lewis structure of atom of a covalent molecule shows all the electrons in the valence shell of each atom; the bonds between atoms are shown as shared pairs of electrons.

Atoms are most stable if they have a filled valence shell of electrons. Atoms transfer or share electrons in such a way that they can attain a filled shell of electrons. This stable configuration of electrons is called an octet.

Except for hydrogen and helium, a filled valence shell contains eight electrons. Lewis structures of atom help us to track the valence electrons and predict the types of bond. The number of valence electrons present in each of the elements is to be considered first.

The number of valence electrons determines the number of electrons needed to complete the octet of eight electrons. Simple ions are atoms that have gained or lost electrons to satisfy the octet rule. However, not all compounds follow the octet rule.

Elements in organic compounds are joined by covalent bonds, a sharing of electrons, and each element contributes one electron to the bond.

The number of electrons necessary to complete the octet determines the number of electrons that must be contributed and shared by a different element in a bond. This analysis finally determines the number of bonds that each element may enter into with other elements.

Lewis Structure of Atom: Bond Structure and Valance Electrons

In a single bond atomic structure, two atoms share one pair of electrons and form a s bond. Double bond structure of atom, shares two pairs of electrons and form a s bond and a p bond. In a triple bond structure of atom, two atoms share three pairs of electrons and form a s bond and two p bonds.

Sodium (Na) loses a single electron from its 3s orbital to attain a more stable neon gas configuration (1s 2 2s 2 2p6 ) with no electron in the outer shell.

An atom having a filled valence shell is said to have a closed shell configuration.

The total number of electrons in the valence shell of each atom can be determined from its group number in the periodic table. The shared electrons are called the bonding electrons and may be represented by a line or lines between two atoms.

Role of Non Electrons in Structuring of Atom and Its Role: A Class 11 Guide

The valence electrons that are not being shared are the nonbonding electrons or lone pair electrons, and they are shown in the Lewis structure of atom by dots around the symbol of the atom. A species that has an unpaired electron are called radicals.

Usually they are very reactive, and are believed to play significant roles in aging, cancer and many other ailments. In neutral organic compounds, C forms four bonds, N forms three bonds (and a lone pair), O forms two bonds (and two lone pairs) and H forms one bond.

The number of bonds an atom normally forms is called the valence. Lewis structure of atom shows the connectivity between atoms in a molecule by a number of dots equal to the number of electrons in the outer shell of an atom of that molecule.

How Lewis Dot Structures Helps Understanding Atomic Structure

A pair of electrons is represented by two dots, or a dash. When drawing Lewis structures, it is essential to keep track of the number of electrons available to form bonds and the location of the electrons.

The number of valence electrons of an atom can be obtained from the periodic table because it is equal to the group number of the atom.

Examples: Deriving from Lewis Structure of Atom

For example, hydrogen (H) in Group 1A has one valence electron, carbon (C) in Group 4A has four valence electrons, and fluorine (F) in Group 7A has seven valence electrons.

To write the Lewis formula of CH₃F, first of all, we have to find the total number of valence electrons of all the atoms involved in this structure of atom, i.e. C, H and F, having four, one and seven valence electrons, respectively.

The carbon atom bonds with three hydrogen atoms and one fluorine atom, and it requires four pairs of electrons.

The remaining six valence electrons are with the fluorine atom in the three nonbonding pairs.

In the periodic table, the period 2 elements C, N, O, and F have valence electrons that belong to the second shell (2s and three 2p).

The shell can be completely filled with eight electrons. In period 3, elements Si, P, S and Cl have the valence electrons that belong to the third shell (3s, three 3p and five 3d ).

The shell is only partially filled with eight electrons in 3s and three 3p, and the five 3d orbitals can accommodate an additional ten electrons.

For these differences in valence shell orbitals available to elements of the second and third periods, we see significant differences in the covalent bonding of oxygen and sulphur, and of nitrogen and phosphorus.

Although oxygen and nitrogen can accommodate no more than eight electrons in their valence shells, many phosphorus-containing compounds have 10 electrons in the valence shell of phosphorus, and many sulphur-containing compounds have 10 and even 12 electrons in the valence shell of sulphur.

Deriving the Structure of Atom from Lewis Structure: Class 11

• (a) Draw a tentative structure for atom. The element with the least number of atoms is usually the central element.
• (b) Calculate the number of valence electrons for all atoms in the compound.
• (c) Put a pair of electrons between each symbol.
• (d) Place pairs of electrons around atoms beginning with the outer atom until each has eight electrons, except for hydrogen. If an atom other than hydrogen has fewer than eight electrons then move unshared pairs to form multiple bonds.

If the structure is an ion, electrons are added or subtracted to give the proper charge. Lewis structures of atom are useful as they show what atoms are bonded together, and whether any atoms possess lone pairs of electrons or have a formal charge.

A formal charge is the difference between the number of valence electrons an atom actually has when it is not bonded to any other atoms, and the number of nonbonding electrons and half of its bonding electrons.

Thus, a positive or negative charge assigned to an atom is called a formal charge. The decision as to where to put the charge is made by calculating the formal charge for each atom in an ion or a molecule.

For example, the hydronium ion (H3O⁺) is positively charged and the oxygen atom has a formal charge of +1.

An uncharged oxygen atom must have six electrons in its valence shell. In the hydronium ion, oxygen bonds with three hydrogen atoms.

So, only five electrons effectively belong to oxygen, which is one less than the valence electrons. Thus, oxygen bears a formal charge of +1.

Elements of the second period, including carbon, nitrogen, oxygen and fluorine, cannot accommodate more than eight electrons as they have only four orbitals (2s, 2px, 2py and 2pz) in their valence shells.

Various types of Chemical Bonding in Structuring of Atom for Class 11

A chemical bond is the attractive force that holds two atoms together. Let us learn the role of Chemical Bonding in Structuring of Atom for Class 11.

Valence electrons take part in bonding. An atom that gains electrons becomes an anion, a negatively charged ion, and an atom that loses electrons becomes a cation, a positively charged ion.

Metals tend to lose electrons and nonmetals tend to gain electrons. While cations are smaller than atoms, anions are larger. Atoms decrease in size as they go across a period, and increase in size as they go down a group and increase the number of shells to hold electrons.

The energy required for removing an electron from an atom or ion in the gas phase is called ionization energy.

Atoms can have a series of ionization energies, since more than one electron can always be removed, except for hydrogen.

In general, the first ionization energies increase across a period and decrease down the group. Adding more electrons is easier than removing electrons. It requires a vast amount of energy to remove electrons.

Ionic Bonds in the Structure of Atoms

Ionic bonds result from the transfer of one or more electrons between atoms. The more electronegative atom gains one or more valence electrons and hence becomes an anion. The less electronegative atom loses one or more valence electrons and becomes a cation.

A single headed arrow indicates a single electron transfer from the less electro negative element to the more electronegative atom. Ionic compounds are held together by the attraction of opposite charges.

Thus, ionic bonds consist of the electrostatic attraction between positively and negatively charged ions. Ionic bonds are commonly formed between reactive metals, electropositive elements (on the left hand side of the periodic table), and nonmetals, electronegative elements (on the right hand side of the periodic table).

For example, Na (electronegativity 0.9) easily gives up an electron, and Cl (electronegativity 3.0) readily accepts an electron to form an ionic bond.

In the formation of ionic compound Na⁺Cl^–, the single 3s valence electron of Na is transferred to the partially filled valence shell of chlorine.

Na([Ne]3s¹) + Cl([Ne]3s²3p⁵) → Na⁺([Ne]3s⁰) + Cl^–([Ne]3s²3p⁶) → Na⁺Cl^–

Covalent Bonds in the Structure of Atoms

Covalent bonds result from the sharing of electrons between atoms. In this case, instead of giving up or acquiring electrons, an atom can obtain a filled valence shell by sharing electrons.

For example, two chlorine atoms can achieve a filled valence shell of 18 electrons by sharing their unpaired valence electrons.

Similarly, hydrogen and fluorine can form a covalent bond by sharing electrons. By doing this, hydrogen fills its only shell and fluorine achieves its valence shell of eight electrons.

Understanding Nonpolar and Polar Covalent Bonds in the Structure of Atoms: A Class 11 Guide

In general, most bonds within organic molecules, including various drug molecules, are covalent. The exceptions are compounds that possess metal atoms, where the metal atoms should be treated as ions.

If a bond is covalent, it is possible to identify whether it is a polar or nonpolar bond. In a nonpolar covalent bond, the electrons are shared equally between two atoms.

As for example, H-H and F-F. Bonds between different atoms usually result in the electrons being attracted to one atom more strongly than the other. Such an unequal sharing of the pair of bonding electrons results in a polar covalent bond.

How Polar Covalent Bonds Effects the Structure of Atoms

In a polar covalent bond, one atom has a greater attraction for the electrons than the other atom, e.g. chloromethane (CH₃Cl). When chlorine is bonded to carbon, the bonding electrons are attracted more strongly to chlorine.

In other words, in a polar covalent bond, the electron pair is not shared equally. This results in a small partial positive charge on the carbon, and an equal but opposite partial negative charge on the chlorine.

Bond polarity is measured by dipole moment (m, which for chloromethane is 1.87). The dipole moment is measured in a unit called the debye (D). Generally, the C-H bond is considered nonpolar.

3D Structures of Atom of Covalent Compounds

Chemists use two parameters, bond lengths and bond angles, to describe the 3D structures of atom of covalent compounds. A bond length is the average distance between the nuclei of the atoms that are covalently bonded together.

A bond angle is the angle formed by the interaction of two covalent bonds at the atom common to both. Covalent bonds are formed when atomic orbitals overlap.

The overlap of atomic orbitals is called hybridization, and the resulting atomic orbitals are called hybrid orbitals. There are two types of orbital overlap, which form sigma (s) and pi (p) bonds. Pi bonds never occur alone without the bonded atoms also being joined by a s bond.

Therefore, a double bond consists of a s bond and a p bond, whereas a triple bond consists of a s bond and two p bonds. A sigma overlap occurs when there is one bonding interaction that results from the overlap of two s orbitals or an s orbital overlaps a p orbital or two p orbitals overlap head to head.

A p overlap occurs only when two bonding interactions result from the sideways overlap of two parallel p orbitals. The s orbital is spherical in shape and p orbitals are in dumbbell shapes.

Let us consider the formation of s overlap in the hydrogen molecule (H₂), from two hydrogen atoms. Each hydrogen atom has one electron, which occupies the 1s orbital. The overlap of two s orbitals, one from each of two hydrogen atoms, forms a s bond.

The electron density of a s bond is greatest along the axis of the bond. Since s orbitals are spherical in shape, two hydrogen atoms can approach one another from any direction resulting in a strong s bond.

Electronegativity and Chemical Bonding for Structure of Atom Class 11

Electronegativity play an crucial role in structuring of an atom. It is the ability of an atom that is bonded to another atom or atoms to attract electrons strongly towards it.

This competition for electron density is scaled by electronegativity values. Elements with higher electro negativity values have greater attraction for bonding electrons.

Thus, the electronegativity of an atom is related to bond polarity. The difference in electronegativity between two atoms can be used to measure the polarity of the bonding between them.

How the Difference in Electronegativity Affects Structure of An Atom: A Class 11 Perspective

The greater the difference in electronegativity between the bonded atoms, the greater is the polarity of the bond. If the difference is great enough, electrons are transferred from the less electronegative atom to the more electronegative one, hence an ionic bond is formed.

Only if the two atoms have exactly the same electronegativity is a nonpolar bond formed. Electronegativity increases from left to right and bottom to top in the periodic table as shown below (electronegativity is shown in parentheses).

In general, if the electronegativity difference is equal to or less than 0.5 the bond is nonpolar covalent, and if the electronegativity difference between bonded atoms is 0.5–1.9 the bond is polar covalent.

If the difference in electronegativities between the two atoms is 2.0 or greater, the bond is ionic. Some examples are shown below.

Electrons in a polar covalent bond are unequally shared between the two bonded atoms, which results in partial positive and negative charges.

The separation of the partial charges creates a dipole. The word dipole means two poles, the separated partial positive and negative charges. A polar molecule results when a molecule contains polar bonds in an unsymmetrical arrangement.

Nonpolar molecules whose atoms have equal or nearly equal electronegativities have zero or very small dipole moments, as do molecules that have polar bonds but the molecular geometry is symmetrical, allowing the bond dipoles to cancel each other.

Bond polarity and intermolecular forces in Structuring an Atom

Bond polarity is a useful concept for describing the sharing of electrons between atoms. The shared electron pairs between two atoms are not necessarily shared equally and this leads to a bond polarity.

Atoms, such as nitrogen, oxygen and halogens, that are more electronegative than carbon have a tendency to have partial negative charges.

Atoms such as carbon and hydrogen have a tendency to be more neutral or have partial positive charges. Thus, bond polarity arises from the difference in electronegativities of two atoms participating in the bond formation.

This also depends on the attraction forces between molecules, and these interactions are called intermolecular interactions or forces. The physical properties, e.g. boiling points, melting points and solubilities of the molecules are determined, to a large extent, by intermolecular nonbonding interactions.

There are three types of nonbonding intermolecular interaction:

• dipole– dipole interactions,
• van der Waals forces, and
• hydrogen bonding.

These interactions increase significantly as the molecular weights increase, and also increase with increasing polarity of the molecules.

How Dipole-Dipole Interactions Shape the Structure of Atoms: A Class 11 Overview

The interactions between the positive end of one dipole and the negative end of another dipole are called dipole–dipole interactions.

As a result of dipole–dipole interactions, polar molecules are held together more strongly than nonpolar molecules.

Dipole–dipole interactions arise when electrons are not equally shared in the covalent bonds because of the difference in electronegativity.

For example, hydrogen fluoride has a dipole moment of 1.98 D, which lies along the H-F bond.

As the fluorine atom has greater electronegativity than the hydrogen atom, the electrons are pulled towards fluorine, as shown below.

The arrow indicates the electrons are towards the more electronegative atom fluorine. The δ⁺ and δ^– symbols indicate partial positive and negative charges.

Understanding Vander Waal’s Forces in the Structure of Atoms: A Class 11 Guide

Relatively weak forces of attraction that exist between nonpolar molecules are called van der Waals forces or London dispersion forces. Dispersion forces between molecules are much weaker than the covalent bonds within molecules.

Electrons move continuously within bonds and molecules, so at any time one side of the molecule can have more electron density than the other side, which gives rise to a temporary dipole.

Because the dipoles in the molecules are induced, the interactions between the molecules are also called induced dipole–induced dipole interactions. van der Waals forces are the weakest of all the intermolecular interactions.

Alkenes are nonpolar molecules, because the electronegativities of carbon and hydrogen are similar.

Consequently, there are no significant partial charges on any of the atoms in an alkane. Therefore, the size of the van der Waals forces that hold alkane molecules together depends on the area of contact between the molecules.

The greater the area of contact, the stronger are the van der Waals forces, and the greater is the amount of energy required to overcome these forces.

For example, isobutane (b.p. -10.2 °C) and butane (b.p. -0.6 °C), both with the molecular formula C₄H₁₀, have different boiling points.

Isobutane is a more compact molecule than butane. Thus, butane molecules have a greater surface area for interaction with each other than isobutane.

The stronger interactions that are possible for n-butane are reflected in its boiling point, which is higher than the boiling point of isobutane.

Understanding Hydrogen Bonding in the Structure of Atoms: A Class 11 Guide with PPT

Hydrogen bonding is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of the same (intramolecular) or a different molecule (intermolecular).

It is an unusually strong force of attraction between highly polar molecules in which hydrogen is covalently bonded to nitrogen, oxygen or fluorine.

Therefore, a hydrogen bond is a special type of interaction between atoms. A hydrogen bond is formed whenever a polar covalent bond involving a hydrogen atom is in close proximity to an electronegative atom such as O or N.

The attractive forces of hydrogen bonding are usually indicated by a dashed line rather than the solid line used for a covalent bond.

For example, water molecules form intermolecular hydrogen bonding.

The above diagram shows a cluster of water molecules in the liquid state. Water is a polar molecule due to the electronegativity difference between hydrogen and oxygen atoms.

The polarity of the water molecule with the attraction of the positive and negative partial charges is the basis for the hydrogen bonding.

The Crucial Role of Hydrogen Bonds: How Atomic Structure Determines Their Strength

Hydrogen bonding is responsible for certain characteristics of water, e.g. surface tension, viscosity and vapour pressure.

Hydrogen bonding occurs with hydrogen atoms covalently bonded to oxygen, fluorine or nitrogen, but not with chlorine, which has larger atom size.

The strength of a hydrogen bond involving an oxygen, a fluorine or a nitrogen atom ranges from 3 to 10 kcal/mol, making hydrogen bonds the strongest known type of intermolecular interaction. The intermolecular hydrogen bonding in water is responsible for the unexpectedly high boiling point of water (b.p. 100 °C).

Hydrogen bonds are interactions between molecules and should not be confused with covalent bonds to hydrogen within a molecule. Hydrogen bonding is usually stronger than normal dipole forces between molecules, but not as strong as normal ionic or covalent bonds.

The hydrogen bond is of fundamental importance in biology. The hydrogen bond is said to be the ‘bond of life’. The double helix structure of DNA is formed and held together with hydrogen bonds.

The nature of the hydrogen bonds in proteins dictates their properties and behaviour. Intramolecular hydrogen bonds (within the molecule) in proteins result in the formation of globular proteins, e.g. enzymes or hormones.

On the other hand, intermolecular hydrogen bonds (between different molecules) tend to give insoluble proteins such as fibrous protein.

Cellulose, a polysaccharide, molecules are held together through hydrogen bonding, which provides plants with rigidity and protection. In drug–receptor binding, hydrogen bonding often plays an important role.

Significance of Chemical Bonding in Atomic Structure in Drug–Receptor Interactions

Now, Let us learn the significance of Chemical Bonding in formation of Atomic Structure in Drug–Receptor Interactions

Most drugs interact with receptor sites localized in macromolecules that have protein-like properties and specific three-dimensional shapes.

A receptor is the specific chemical constituents of the cell with which a drug interacts to produce its pharmacological effects. One may consider that every protein that acts as the molecular target for a certain drug should be called a receptor.

However, this term mainly incorporates those proteins that play an important role in the intercellular communication via chemical messengers. As such, enzymes, ion channels and carriers are usually not classified as receptors.

The term receptor is mostly reserved for those protein structures that serve as intracellular antennas for chemical messengers.

Upon recognition of the appropriate chemical signal (known as the ligand), the receptor proteins transmit the signal into a biochemical change in the target cell via a wide variety of possible pathways.

A minimum three-point attachment of a drug to a receptor site is essential for the desired effect. In most cases, a specific chemical structure is required for the receptor site and a complementary drug structure. Slight changes in the molecular structure of the drug may drastically change specificity, and thus the efficacy.

However, there are some drugs that act exclusively by physical means outside cells, and do not involve any binding to the receptors.

These sites include external surfaces of skin and gastrointestinal tract. Drugs also act outside cell membranes by chemical interactions, e.g. neutralization of stomach acid by antacids.

Unveiling Drug–Receptor Interaction: The Crucial Role of Atomic Structure

How drug-receptor interaction plays an important role in structuring the atoms of a drug molecule.

The drug–receptor interaction, i.e. the binding of a drug molecule to its receptor, is governed by various types of chemical bonding that have been discussed earlier. A variety of chemical forces may result in a temporary binding of the drug to its receptor.

Interaction takes place by utilizing the same bonding forces as involved when simple molecules interact.

As for e.g. covalent (40– 140 kcal/mol), ionic (10 kcal/mol), ion–dipole (1–7 kcal/mol), dipole–dipole (1–7 kcal/mol), van der Waals (0.5–1 kcal/mol), hydrogen bonding (1–7 kcal/ mol) and hydrophobic interactions (1 kcal/mol). However, most useful drugs bind through the use of multiple weak bonds (ionic and weaker).

Covalent bonds are strong, and practically irreversible. Since the drug– receptor interaction is a reversible process, covalent bond formation is rather rare except in a few situations.

Some drugs that interfere with DNA function by chemically modifying specific nucleotides are mitomycin C, cisplatin and anthramycin.

How Atomic Structure Shapes the Impact of Drugs That Interfere with DNA Function

Mitomycin C is a well characterized antitumour agent, which forms a covalent interaction with DNA after reductive activation, forming a cross-linking structure between guanine bases on adjacent strands of DNA, thereby inhibiting single strand formation.

Similarly, anthramycin is another antitumour drug, which binds covalently to N-2 of guanine located in the minor groove of DNA. Anthramycin has a preference for purine–G–purine sequences (purines are adenine and guanine) with bonding to the middle G.

Cisplatin, an anticancer drug, is a transition metal complex, cis-diamine-dichloro-platinum. The effect of the drug is due to the ability to platinate the N-7 of guanine on the major groove site of the DNA double helix.

This chemical modification of the platinum atom cross-links two adjacent guanines on the same DNA strand, interfering with the mobility of DNA polymerases.

Many drugs are acids or amines, easily ionized at physiological pH, and able to form ionic bonds by the attraction of opposite charges in the receptor site.

Exploring Ionic Interaction: How Atomic Structure Influences Chemical Bonds

For example the ionic interaction between the protonated amino group on salbutamol or the quaternary ammonium on acetylcholine and the dissociated carboxylic acid group of its receptor site.

Similarly, the dissociated carboxylic group on the drug can bind with amino groups on the receptor. Ion–dipole and dipole–dipole bonds have similar interactions, but are more complicated and are weaker than ionic bonds.

Exploring Polar–polar Interaction: How Atomic Structure Influences Chemical Bonds

Polar–polar interaction, e.g. hydrogen bonding, is also an important binding force in drug–receptor interaction, because the drug–receptor interaction is basically an exchange of the hydrogen bond between a drug molecule, surrounding water and the receptor site.

Formation of hydrophobic bonds between nonpolar hydrocarbon groups on the drug and those in the receptor site is also common. Although these bonds are not very specific, the interactions take place to exclude water molecules.

Repulsive forces that decrease the stability of the drug–receptor interaction include repulsion of like charges and steric hindrance.

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