Chemical Bonding and Molecular Structure

The concept or idea that matter is made of particles is referred to as the particulate nature of matter. We have seen that matter is made up of particles from the diagrammatic representation of the three states of matter. The fact that particles of matter are in constant motion can be seen from the fact that when a reagent bottle containing ammonia is opened in the chemistry laboratory, the odour will be perceived by somebody standing outside the laboratory after some seconds. Again if hydrogen sulphide gas is generated inside a chemistry laboratory, the odour of the gas can easily be perceived at different distances within the environment, or when a gas cylinder is opened in the laboratory without lighting the bunsen burner, the odour of the gas can also be perceived by people at different distances within the environment. If matter were not made of particles that were in constant motion, the odour of the gas will not be perceived at different distances from the source. Scientists discovered that matter is made up of particles which can be atoms, molecules or ions.

 

Concept of the Atom

An atom is the smallest particle of an element, which can take part in a chemical reaction. If a piece of solid element like copper or zinc is ground into very tiny pieces, the smallest part of it which can take part in a chemical reaction is called an atom.

The word "atom" comes from the Greek word "atomos", meaning "indivisible" and early philosophers like Democritus (around 400 BC) suggested that matter is made of tiny particles that can’t be divided further. Later, scientists like John Dalton, J.J. Thomson, Ernest Rutherford, Niels Bohr, and modern quantum physicists developed our current understanding of atomic structure.

 

Dalton’s Atomic Theory

Dalton’s Atomic Theory was proposed by John Dalton in 1803. It was the first scientific theory to describe the nature of matter in terms of atoms.

 

Postulates of Dalton’s Atomic Theory

1. All matter is made up of tiny, indivisible particles called atoms.
   Dalton believed atoms could not be broken down into smaller parts.
2. Atoms of the same element are identical in mass and properties.
   For example, all hydrogen atoms were thought to be exactly alike.
3. Atoms of different elements have different masses and properties.
   For instance, hydrogen and oxygen atoms are different in weight and behavior.
4. Atoms cannot be created, divided, or destroyed in a chemical reaction.
   Instead, they are rearranged during a reaction.
5. Atoms combine in simple whole-number ratios to form compounds.
   For example, water is formed when 2 hydrogen atoms combine with 1 oxygen atom (H₂O).
6. In chemical reactions, atoms are rearranged but not changed into other types of atoms.

 

Limitations of Dalton’s Atomic Theory

Over time, scientists discovered several facts that showed Dalton’s ideas were not entirely accurate. These are some of the main limitations:

1. Atoms are not indivisible.
   Today, we know atoms are made of subatomic particles – protons, neutrons, and electrons. An atom can be divided into these parts.
2. Atoms of the same element are not always identical.
   The discovery of isotopes showed that atoms of the same element can have different masses (due to different numbers of neutrons).
3. Atoms can be converted into other atoms.
   In nuclear reactions, atoms can change from one type into another (e.g., radioactive decay), which violates Dalton’s fourth postulate.
4. Atoms do not always combine in simple whole-number ratios.
   In some compounds, atoms combine in complex ratios (e.g., in non-stoichiometric compounds found in inorganic chemistry).
5. The theory couldn’t explain the existence of charged particles.
   Dalton’s model didn’t account for ions (atoms that gain or lose electrons to become charged).

 

Modern Corrections and Improvements

Scientific advancements in the 19th and 20th centuries helped correct and expand Dalton's ideas:

1. Discovery of Electrons – J.J. Thomson discovered electrons in 1897, proving atoms are divisible.
2. Discovery of the Nucleus – Ernest Rutherford’s gold foil experiment in 1911 showed that atoms have a dense central nucleus.
3. Isotopes Explained – Frederick Soddy explained the concept of isotopes, which corrected the idea that all atoms of an element are identical.
4. Discovery of Neutrons – James Chadwick discovered neutrons in 1932, further confirming that atoms have internal structure.
5. Quantum Mechanics – Modern atomic theory includes quantum models which describe how electrons move in orbitals, not fixed paths.

 

 

 

The three main subatomic particles of atom:

Particle	Charge	Location in Atom	Mass
Proton	+1 (Positive)	In the nucleus	1 atomic mass unit
Neutron	0 (Neutral)	In the nucleus	1 atomic mass unit
Electron	–1 (Negative)	Orbits around nucleus	Very tiny (1/1836)

 

Structure of an Atom
Dalton claimed atoms are indivisible and indestructible, but starting around 1850, people started to identify sub-atomic particles. Nowadays we know atoms are composed of subatomic particles (electrons, protons and neutrons). Faraday in 1834 passed electricity through an aqueous solution, and brought about chemical changes (which implies that chemicals are related to electricity).

There are two types of electrical charge, positive (+) (+ve) and negative (-) (-ve).

Law of Electrostatic Attraction (1784): like charges repel one another, unlike (opposite) charges attract.

Thomson is credited with discovering the electron in 1897.
By passing current through a gas at low pressure (in a gas discharge tube), he generated species (cathode rays) that had very low mass (much less than atoms) and were negatively charged.

Nucleus is the center of the atom, where protons and neutrons are located It is dense and positively charged. Electrons on the other hand move in regions called shells or energy levels around the nucleus. They are involved in chemical bonding and reactions.

 

Atomic Number and Mass Number

Every atom has an atomic number, which refers to the number of protons found in the nucleus of that atom. This number is incredibly important because it defines the identity of the element. For example, an atom that has one proton is always hydrogen, while an atom with six protons is always carbon. No matter what form it takes or how many electrons or neutrons it has, an atom with six protons will always be a carbon atom.

In addition to the atomic number, atoms also have a mass number. The mass number is the total number of protons and neutrons in the atom’s nucleus. Since electrons are extremely small and have very little mass compared to protons and neutrons, they are not included in the mass number.

In nutshell,

Mass Number (A) = Number of Protons (Z) + Number of Neutrons (N)

Example:

Carbon has 6 protons and 6 neutrons

Hence;

➤ Atomic number = 6

➤ Mass number = 6 + 6 = 12

An element X with mass number 14 and atomic number of 7 has the symbol X. Thus, an element X has a mass number of 14 and an atomic number of 7

Isotopy

Isotopy refers to the occurrence of atoms of the same element that have the same atomic number but different mass numbers. These atoms of the same element having the same atomic number but different mass numbers are called isotopes. The difference in the mass number of these isotopes are due to the difference in the number of neutrons in their nuclei. The occurrence and natural abundances of isotopes can be experimentally determined using an instrument called a mass spectrometer. Mass spectrometry (MS) is widely used in chemistry, forensics, medicine, environmental science, and many other fields to analyze and help identify the substances in a sample of material.

Example of some elements and their isotopes: Hydrogen has three isotopes (1, 2, and 3), Chlorine has two isotopes (35 and 37), Copper has two (63 and 65) etc.

Play Question

2. Magnesium has three naturally occurring isotopes: Mg-24 (78.99%), Mg-25 (10.00%), Mg-26 (11.01%). Calculate the average atomic mass of magnesium.

3. Boron has two isotopes: B-10 (19.9%), B-11 (80.1%). Calculate the average atomic mass of boron.

 

 

 

 

Molecules

A molecule of a compound or element is the smallest particle of a compound or element which is

capable of independent existence. A molecule may be composed of atoms of the same element or

atoms of different elements. For example, a nitrogen molecule (N₂) contains two atoms of nitrogen, an oxygen molecule (O₂) contains two atoms of oxygen but a molecule of water (H₂O)

contains two atoms of hydrogen and one atom of oxygen.

 

Homogeneous and Heterogeneous Molecules

These terms are often used to describe the types of atoms present in a molecule.

1. Homogeneous Molecules: is a molecule made up of atoms of the same element.

Key Features:

• All atoms in the molecule are identical.
• They are typically elemental molecules.
• Often involve covalent bonds.

Examples:

Molecule	Composition	Explanation
O₂	Oxygen + Oxygen	Both atoms are oxygen — same element
N₂	Nitrogen + Nitrogen	Same type of atom
Cl₂	Chlorine + Chlorine	Identical atoms bonded together

 

 

 

2. Heterogeneous Molecules

Definition:

A heterogeneous molecule is a molecule made up of atoms of different elements.

Key Features:

• Atoms in the molecule are not identical.
• They are always compound molecules.
• Can have covalent or ionic bonds.

Example:

Molecule	Composition	Explanation
H₂O	Hydrogen + Oxygen	Two different elements
CO₂	Carbon + Oxygen	Different atoms chemically bonded
NH₃	Nitrogen + Hydrogen	Mixed types of atoms
NaCl (ionic)	Sodium + Chloride	Different atoms, ionic compound

 

Atomicity:

Atomicity of an element is the number of atoms in one molecule of the element. Elements whose molecule contains one atom is said to be mono-atomic e.g copper (Cu), sodium (Na).

Element whose molecule contains two atoms are said to be diatomic e.g O₂. Element whose

molecule contains three atoms are tri-atomic e.g O₃. Element whose molecule contains four atoms

are said to be tetra-atomic while a molecule containing more than four atoms is said to be poly-atomic e.g S₈, CH₄, C₆H₁₂O₆

 

IONS

An ion is an electrically charged atom or group of atoms. It is formed as a result of the loss or

gain of electrons. The electrons lost or gained are equal to the valence of the ions.

Formation of Ions: Examples

a. Sodium atom (Na) → Sodium ion (Na⁺)

Electron configuration:
Na: 2,8,1 → loses 1 electron → Na⁺: 2,8
Stable like neon (Ne)

b. Chlorine atom (Cl) → Chloride ion (Cl⁻)

Cl: 2,8,7 → gains 1 electron → Cl⁻: 2,8,8
Stable like argon (Ar)

 

There are two types of ions. These are:

i. Positively charged ion or cation

ii. Negatively charged ion or anion.

A cation is formed when an atom or group loses electrons. For example

Cations include: ammonium ion (NH⁴⁺), hydrogen ion (H⁺) and metallic ions such as Na⁺, Ca²⁺  etc.

An anion is formed when an atom or group gains electrons. For example;

Anions include hydroxide ion and acid radicals. Examples of acid radical are:

SO₄^2- = Tetraoxosulphate (vi) ion

NO^3- = Trioxonitrate (v) ion