Structure of atom

An atom, the basic building block of matter, is made up of a central nucleus and an electron cloud, forming the core of atomic structure. The nucleus, composed of positively charged protons and neutral neutrons, is dense and carries a positive charge. Surrounding the nucleus, electrons orbit in various energy levels or shells, each with a specific capacity for holding electrons. The quantum mechanical model further refines this understanding by describing electrons in terms of probabilities and orbitals within these energy levels. This intricate atomic structure is essential for comprehending the behavior and interactions of atoms in both chemistry and physics.

1.0The Discovery and Nature of Atoms

Democritus, who was a Greek philosopher, introduced the concept of the atom around 450 BC, describing it as solid and varying in size, mass, shape, position, and arrangement. In the 1800s, John Dalton reintroduced the idea, proposing that matter is made of indivisible, indestructible atoms. In 1911, Rutherford provided the modern atomic structure.

2.0Basic Components of an Atom

1. Nucleus: The nucleus is the dense central core of an atom, containing two types of subatomic particles:

• Protons: These are positively charged particles. The number of protons in the nucleus, known as the atomic number, uniquely defines each element. For instance, hydrogen has one proton, while carbon has six. The positive charge of protons provides the overall positive charge of the nucleus.
• Neutrons: These particles are neutral and have no electric charge. The number of neutrons in the nucleus can differ even among atoms of the same element, leading to the formation of isotopes. Isotopes are different forms of a chemical element that have the same number of protons but varying numbers of neutrons. For instance, Carbon-12 and Carbon-14 are isotopes of carbon.

2. Electron Cloud: Surrounding the nucleus is the electron cloud, which is composed of negatively charged particles which called electrons:
3. Electrons: Electrons are basically negatively charged particles that move around the nucleus in regions called orbitals, which are organised into energy levels or shells. These are much smaller and lighter than protons and neutrons. 

3.0Energy Levels/Shells and Electronic Configuration

• Electrons occupy different energy levels, starting with the lowest energy level closest to the nucleus. 
• The first energy level (n=1) can hold up to 2 electrons.
• The second energy level (n=2) can hold up to 8 electrons.
• The third energy level (n=3) can hold up to 18 electrons, and so on.

Orbitals: 

Within each energy level, electrons are distributed among different types of orbitals (s, p, d, f), each with a specific shape and orientation. 

1. s-orbital: Spherical shape, can hold up to 2 electrons.
2. p-orbital: Dumbbell shape, can hold up to 6 electrons (3 orbitals, each holding 2 electrons).
3. d-orbital: More complex shapes, can hold up to 10 electrons (5 orbitals).
4. f-orbital: Even more complex shapes, can hold up to 14 electrons (7 orbitals).

Electron Configuration

• Electrons occupy specific energy levels or shells around the nucleus. The arrangement of electrons in these levels is called the electron configuration.
• The electron configuration of an element is mainly responsible for its chemical properties.
• Electron configuration of an element describes the distribution of electrons in an atom's orbitals, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.
• Electrons fill orbitals starting with the lowest energy level, with each orbital holding a maximum of two electrons with opposite spins. Electrons fill degenerate orbitals singly before pairing. The notation involves the energy level and orbital type (e.g., 1s, 2p) with superscripts indicating the number of electrons (e.g., 1s²). For example, carbon (atomic number 6) has the configuration 1s²2s²2p². A condensed notation uses the nearest noble gas as shorthand, such as sodium (Na) with [Ne]3s¹.

4.0Atomic Number, Mass Number and Isotopes

The atomic number and Mass Number are essential because they uniquely identify an element and identify the specific isotope of an element, which affects the element's atomic mass and nuclear stability.

1. Atomic Number (Z): The number of protons in the nucleus of an atom is known as the atomic number, which defines the element. For example, a carbon atom has an atomic number of 6, meaning they each have 6 protons in their nucleus. This atomic number uniquely identifies the element as carbon, distinguishing it from other elements.
2. Mass Number (A): The total number of protons and neutrons in an atom is called mass number. For example, a carbon atom which has mass number of 12 contains 6 protons and 6 neutrons. The mass number provides information about the overall mass of the atom, as it accounts for the fused mass of its protons and neutrons, which are the heaviest subatomic particles. 
3. Isotopes: An Isotope is a species of chemical element that has the same number of protons but different numbers of neutrons. This results in atoms of the same element having different atomic masses. 

• Isotopes of an element exhibit almost identical chemical behavior but can have different physical properties.
• Hydrogen has three isotopes: Protium, Deuterium and Tritium. Here is a brief introduction:

5.0Evolution of Atomic Models

The understanding of the atom has progressed significantly over centuries through various models proposed by scientists. Here is a brief overview of the key models that have shaped our current understanding of atomic structure:

1. Democritus' Model (c. 450 BC)

• Atoms are indivisible, solid particles varying in size, shape, and mass.
• Atoms are known as fundamental building blocks of matter.

2. Dalton's Model (1803)

• Matter is composed of indivisible atoms, each element consisting of identical atoms.
• Atoms of different elements diverge in mass and properties, and chemical reactions involve the rearrangement of atoms.

3. Thomson's Plum Pudding Model (1904)

• Atoms are composed of electrons scattered within a positively charged "soup."
• Discovered the electron, proposing that atoms are divisible.

4. Rutherford's Nuclear Model (1911)

• Atoms have a small, packed, positively charged nucleus surrounded by electrons.
• Most of an atom's mass is concentrated in the nucleus, with electrons orbiting around it.

5. Bohr's Model (1913)

• Electrons revolve in the nucleus in discrete energy levels or shells.
• Electrons can jump between energy levels by absorbing or emitting energy.

6. Quantum Mechanical Model (1920s)

• Electrons exist in probability clouds called orbitals, not fixed paths.
• Incorporates principles of quantum mechanics, describing electron positions as probabilities rather than precise locations.