Matter, the substance that occupies physical space and possesses mass, is composed of discrete particles, each playing a unique role in its composition and behavior. These particles include atoms, molecules, and ions, which form the fundamental building blocks of matter.

An atom, considered the fundamental unit of an element, is the smallest indivisible particle that retains the characteristic properties of that element. It consists of a nucleus, composed of protons and neutrons, surrounded by orbiting electrons. Atoms combine and interact with one another through chemical reactions to form compounds and molecules.

A molecule, on the other hand, represents the smallest particle of a substance that can exist independently and still exhibit the chemical properties of that particular substance. Molecules are formed when two or more atoms chemically bond together through shared electrons. Consequently, molecules possess distinct physical and chemical characteristics that differ from the individual atoms they are composed of.

The atomicity of an element refers to the number of atoms present within a single molecule of that element. Depending on the specific element, the atomicity can vary. For instance, certain elements exist in a monatomic state, meaning they consist of single atoms in their molecular form. These elements, such as helium (He) and neon (Ne), exist as independent atoms and do not readily bond with other atoms.

In contrast, some elements have a diatomic atomicity, meaning their molecules consist of two atoms bonded together. Well-known examples include oxygen (O₂), nitrogen (N₂), and hydrogen (H₂). These elements are most stable when they form pairs of atoms, sharing electrons between them to achieve a more stable electron configuration.

Furthermore, there are elements with a triatomic atomicity, where three atoms are bonded together within a single molecule. Notable examples include ozone (O₃) and carbon dioxide (CO₂). In these cases, three atoms form a stable molecular structure, exhibiting unique chemical and physical properties distinct from monatomic or diatomic substances.

Understanding the atomicity of elements provides valuable insights into their behavior, reactivity, and the types of bonds they form. By categorizing elements based on their atomicity, scientists can better predict their chemical properties and study their interactions with other substances.

In summary, matter is composed of discrete particles, including atoms, molecules, and ions. Atoms represent the smallest units of elements, capable of participating in chemical reactions. Molecules, formed through the bonding of atoms, retain the chemical properties of the substances they comprise. The atomicity of an element refers to the number of atoms within its molecule, with monatomic, diatomic, and triatomic states representing different arrangements of atoms. This knowledge helps scientists categorize elements and understand their unique characteristics and behaviors.


Element                                              Atomicity

Hydrogen                                             Diatomic

Oxygen                                                Diatomic

Nitrogen                                               Diatomic

Neon                                                    Monoatomic

Helium                                                  Monoatomic

Argon                                                   Monoatomic


An ion, whether it is a single atom or a collection of atoms, is characterized by the presence of an electric charge. This charge arises from the gain or loss of electrons, leading to an imbalance between the number of protons and electrons. Ions play a crucial role in various chemical processes and are classified based on their charge. When a group of atoms carries either a positive or negative charge, they are referred to as radicals.

Specifically, an acid radical represents a small group of atoms that retains its identity and carries a negative charge. Acid radicals are commonly found in acids and contribute to their distinctive chemical properties. These radicals are capable of reacting with other substances, participating in chemical reactions, and influencing the overall behavior of the acid.

For instance, one example of an acid radical is the sulfate ion (SO₄²⁻). It consists of one sulfur atom bonded to four oxygen atoms, forming a tetrahedral structure. The negative charge arises from the gain of two electrons, giving the sulfate ion its characteristic charge. Sulfate ions play a significant role in various chemical processes, such as in the formation of salts and as an indicator of acidity in solutions.

Another example of an acid radical is the nitrate ion (NO₃⁻). It consists of one nitrogen atom bonded to three oxygen atoms, resulting in a trigonal planar structure. Similarly, the negative charge on the nitrate ion is due to the gain of one electron. Nitrate ions are commonly found in fertilizers and explosives, as well as serving as a source of nitrogen in biological systems.

These examples illustrate the significance of acid radicals in chemical systems. By carrying a negative charge while maintaining their structural integrity, acid radicals contribute to the reactivity and behavior of acids. Their presence influences the acidity, stability, and various other properties of the compounds they are part of.

Overall, ions, including radicals, play a fundamental role in chemistry, influencing the behavior and properties of substances. Acid radicals, characterized by a group of atoms carrying a negative charge, maintain their identity and contribute to the distinctive properties of acids. Understanding these acid radicals is vital for comprehending the behavior of acids and their involvement in various chemical reactions.

Generally, ions are grouped as cations and anions. Cations are positively charged ions e.g. Ca2+, Na+, NH4+ e.t.c.

Anions are negatively charged ions e.g… C032-, S042-, Cl, OH, etc.


John Dalton, British Physicist and Chemist (1808) proposed the atomic theory thus:

  1. All elements are made up of small indivisible particles called atoms.
  2. Atoms can neither be created nor destroyed in any chemical reaction.
  3. Atoms of the same elements are exactly alike in aspect and are different from atoms of all other elements.
  4. Atoms of different elements can combine in simple whole-number ratios to form compounds.
  5. All chemical changes result from the combination or separation of the atom


The scientific advancements and discoveries made in the twentieth century have brought to light new information that challenges the completeness of Dalton’s atomic theory. Consequently, modifications to the theory have become necessary in order to account for these novel findings.

Rutherford’s groundbreaking experiments shattered the notion that an atom is an indivisible solid piece, as proposed by Dalton. Through his experiments, Rutherford discovered the presence of subatomic particles within the atom, namely protons, electrons, and neutrons. This understanding revolutionized our perception of atomic structure, revealing that atoms consist of a positively charged nucleus, containing protons and neutrons, orbited by negatively charged electrons. This newfound knowledge demonstrated that atoms possess internal complexity and are not indivisible entities.

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While Dalton’s theory remains valid for ordinary chemical reactions, the advent of nuclear reactions has challenged its applicability. During nuclear reactions, such as nuclear fission, the nucleus of an atom can be split, releasing a tremendous amount of heat energy. This process leads to the destruction of atoms involved, contradicting the idea that atoms are indestructible entities. The recognition of nuclear reactions and their impact on atomic structure necessitates an expansion of Dalton’s theory to encompass these unique phenomena.

The discovery of isotopes further undermines Dalton’s third statement, which asserts that atoms of the same element have identical masses. Isotopes are variants of an element that differ in the number of neutrons present in their nuclei while retaining the same number of protons. For instance, chlorine has two isotopes, chlorine-35 and chlorine-37, with different relative atomic masses. This revelation invalidates the assumption that all atoms of a given element possess identical masses, highlighting the need to modify Dalton’s atomic theory to incorporate the concept of isotopes.

The fourth statement put forth by Dalton holds true for inorganic compounds that typically contain only a few atoms per molecule. However, when considering organic compounds, such as those containing carbon, Dalton’s theory falls short. Carbon is capable of forming highly complex organic molecules, including proteins, starch, and fats, which can consist of thousands of atoms intricately bonded together. These large organic molecules challenge the simplicity and limited scope of Dalton’s original theory, necessitating an expanded understanding of atomic structure to encompass the complexities of organic compounds.

In summary, the twentieth-century discoveries, including Rutherford’s identification of subatomic particles, the recognition of nuclear reactions, the discovery of isotopes, and the understanding of large organic molecules, have collectively rendered Dalton’s atomic theory incomplete. The theory required modifications to account for the presence of subatomic particles, the occurrence of nuclear reactions, the existence of isotopes, and the complexities of organic compounds. These advancements have expanded our understanding of atomic structure and chemical behavior, leading to a more comprehensive and accurate model of the atom.

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