Atoms and Isotopes

This section explores atoms and isotopes covering, the structure of atoms, the development of the atomic model and isotopes. 

Structure of Atoms

Atoms are the basic building blocks of matter. Every element is made up of atoms, and each atom consists of three main subatomic particles: protons, neutrons, and electrons.

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Atomic Number and Mass Number

• Atomic Number (Z): This is the number of protons in an atom's nucleus and defines the element. The atomic number determines the element's position on the periodic table and its chemical properties. For example, carbon has an atomic number of 6, meaning it has 6 protons.
• Mass Number (A): This is the total number of protons and neutrons in an atom's nucleus. Since protons and neutrons have nearly the same mass, the mass number roughly equals the atom's mass. For instance, a carbon atom with 6 protons and 6 neutrons has a mass number of 12 (6 protons + 6 neutrons = 12).

The relationship between atomic number and mass number can be written as: $$\text{Mass Number} = \text{Number of Protons} + \text{Number of Neutrons}$$

Ions (Positive and Negative)

• Positive Ions (Cations): When an atom loses one or more electrons, it becomes positively charged because the number of protons exceeds the number of electrons. For example, if a sodium (Na) atom loses one electron, it becomes a sodium ion (Na⁺).
• Negative Ions (Anions): When an atom gains one or more electrons, it becomes negatively charged because the number of electrons exceeds the number of protons. For instance, when a chlorine (Cl) atom gains an electron, it becomes a chloride ion (Cl⁻).

The formation of ions is crucial in chemical reactions and the formation of compounds.

The Development of the Atomic Model

The understanding of the structure of the atom has evolved over centuries. Scientists have developed different models to explain atomic structure, each providing a better and more accurate picture of the atom. Here are the key developments:

Dalton’s Model (Early 1800s)

John Dalton proposed that atoms were indivisible particles that combined in fixed ratios to form compounds. His model was based on experimental evidence, but it did not account for the internal structure of atoms.

Thomson’s Model (1897) - The "Plum Pudding" Model

J.J. Thomson discovered the electron in 1897 through his experiments with cathode rays. He proposed a model of the atom in which the atom was a sphere of positive charge with electrons embedded inside, like "plums" in a "pudding." This model explained the existence of electrons but did not accurately describe atomic structure in detail.

Rutherford’s Model (1911)

Ernest Rutherford conducted the famous gold foil experiment, in which he fired alpha particles at a thin sheet of gold. Most particles passed through, but some were deflected, suggesting that atoms have a small, dense, positively charged nucleus. Rutherford’s model showed that atoms have a central nucleus, made up of protons, surrounded by electrons. The nucleus was tiny compared to the size of the atom.

Bohr’s Model (1913)

Niels Bohr refined Rutherford’s model by introducing the concept of electrons orbiting the nucleus in specific energy levels or "shells." Bohr’s model explained the stability of the atom and the spectral lines observed in atomic emission spectra. Electrons could only occupy certain allowed orbits and could move between them by absorbing or emitting a specific amount of energy.

Modern Atomic Model (Quantum Model)

Today, we use the quantum mechanical model, which builds on Bohr’s model but incorporates wave-like properties of electrons. According to this model, electrons do not orbit the nucleus in fixed paths but are instead found in "orbitals," regions of space where there is a high probability of finding an electron. The modern model also includes the concept of sub-levels and energy quantisation, further refining our understanding of atomic structure.

Isotopes

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. This means that isotopes of an element have the same atomic number but different mass numbers.

For example:

• Carbon-12 (₆¹²C) has 6 protons and 6 neutrons.
• Carbon-14 (₆¹⁴C) has 6 protons and 8 neutrons.

Although isotopes of an element have the same chemical properties (because they have the same number of electrons), their physical properties, such as mass and stability, can differ. Some isotopes are stable, while others are radioactive and decay over time, emitting radiation.

Key Points to Remember:

• The atomic number defines the element and is the number of protons in the nucleus.
• The mass number is the sum of protons and neutrons in the nucleus.
• Ions are charged atoms formed when atoms gain or lose electrons.
• The development of the atomic model has evolved from Dalton’s indivisible particles to the quantum mechanical model.
• Isotopes are atoms of the same element with different numbers of neutrons, leading to different mass numbers.

Understanding the structure of atoms and isotopes is essential to understanding the behaviour of matter in both chemistry and physics.