In this lesson, students will explore the foundations of Atomic Theory and the structure of the atom. The journey begins with the ancient Greek philosopher Democritus, who proposed that matter is made up of tiny, indivisible particles called atoms, a term derived from the Latin word “atomos.” In the early 19th century, John Dalton expanded on this idea with his Atomic Theory, asserting that all matter consists of these small particles. For a long time, it was believed that atoms could not be divided. However, groundbreaking experiments in the early 20th century by scientists such as Goldstein, J.J. Thomson, Rutherford, and Bohr revealed that atoms are made up of even smaller subatomic particles: electrons, protons, and neutrons. This lesson will cover Rutherford’s contributions to Atomic Theory, how Bohr’s theory differs, the structure of the atom, the concept of isotopes, and the electronic configurations of the first 18 elements in the Periodic Table.

Atom

According to Dalton, an atom is an indivisible, hard, and dense sphere. Atoms of the same element are identical and can combine in different ways to form compounds. However, in the late 1800s and early 1900s, scientists discovered new subatomic particles. In 1886, Goldstein found positively charged particles called protons, and in 1897, J.J. Thomson discovered negatively charged particles known as electrons. This established that electrons and protons are fundamental particles of matter. Based on these findings, Thomson proposed his “plum pudding” theory, suggesting that atoms are solid structures with positive charge and tiny negative particles (electrons) embedded within them, similar to plums in a pudding.

In 1895, Sir William Crookes conducted experiments by passing electric current through gases in a discharge tube at very low pressure. He used a glass tube with two metallic electrodes connected to a high-voltage battery. With the pressure inside the tube set to 10^-4 atm, shining rays were emitted from the cathode when the current was passed, traveling toward the anode. These rays were named “cathode rays” because they originated from the cathode.

Discovery of Proton

In 1886, Goldstein discovered that besides cathode rays, there were other rays in the discharge tube moving in the opposite direction. Using a discharge tube with a perforated cathode, he found that these rays passed through the holes and created a glow on the tube’s walls. He named these “canal rays.” The properties of canal rays are as follows: they travel in straight lines opposite to cathode rays, they are positively charged as shown by their deflection in electric and magnetic fields, and their nature depends on the type of gas in the discharge tube. These rays do not come from the anode; instead, they are produced when cathode rays collide with residual gas molecules, ionizing them. The mass of the particles in canal rays is similar to that of a proton, which is about 1840 times more than that of an electron. Therefore, these rays consist of positively charged particles, with their mass and charge varying based on the gas used. For instance, canal rays produced by hydrogen contain protons.

Discovery of Neutron

Rutherford noted that the atomic mass of elements couldn’t be explained solely by the masses of electrons and protons. In 1920, he predicted the existence of a neutral particle with a mass equal to that of a proton. Scientists searched for this particle, and in 1932, Chadwick discovered the neutron by bombarding a beryllium target with alpha particles. He observed highly penetrating radiation, which he identified as neutrons. Neutrons have no charge, making them neutral, and they are highly penetrating. Their mass is nearly equal to that of a proton.

Rutherford’s Atomic Model

Rutherford conducted the famous “Gold Foil” experiment to explore how negative and positive charges coexist in an atom. He bombarded a very thin gold foil (0.00004 cm thick) with alpha particles, which are helium nuclei emitted by radioactive elements like radium and polonium. To observe the effects of these alpha particles, he used a photographic plate or a zinc sulfide-coated screen. His findings demonstrated that the “plum-pudding” model of the atom was incorrect.

Rutherford made several key observations: most alpha particles passed through the foil without any deflection, while only a few were deflected at large angles, and very few bounced back after hitting the gold foil. Based on these results, he proposed a planetary model of the atom with the following conclusions: most of an atom is empty space; there is a central nucleus containing positive charges; the nucleus is very dense and hard; the nucleus is small compared to the overall size of the atom; electrons revolve around the nucleus; and the atom is neutral, meaning the number of electrons equals the number of protons. Other particles in the nucleus, except for electrons, are called nucleons.

Despite its significance, Rutherford’s model had some defects. According to classical radiation theory, electrons, being charged particles, should continuously emit energy and spiral into the nucleus. If they did this, a continuous spectrum would be expected, but instead, a line spectrum was observed. This discrepancy led scientists to question the validity of Rutherford’s model.

Electronic Configuration

Before diving into electronic configuration, it’s important to understand shells and subshells within an atom. An atom consists of a tiny nucleus at its center, with electrons revolving around it. These electrons occupy different energy levels, or shells, based on their potential energy. Each energy level is represented by the principal quantum number ( n ), which can take values like 1, 2, 3, and so on. These levels are designated by letters: K, L, M, etc. The shell closest to the nucleus (K shell) has the lowest energy, and energy increases as you move further from the nucleus.

In summary, electrons revolve around the nucleus in distinct energy levels, with each shell representing a specific amount of energy. Understanding these concepts sets the stage for learning about electronic configuration, which describes how electrons are distributed among these shells and subshells.

Isotopes

Isotopes are atoms of the same element that have the same atomic number but different mass numbers. This means they have the same number of protons and electrons but differ in the number of neutrons. Because their electronic configurations are the same, isotopes exhibit similar chemical properties. However, they have different physical properties due to their varying mass numbers. Most elements have isotopes, and here we will focus on the isotopes of hydrogen, carbon, chlorine, and uranium.

For example, hydrogen naturally exists as a mixture of three isotopes: protium, deuterium, and tritium. All three have 1 proton and 1 electron, but they differ in their number of neutrons. Protium has no neutrons, deuterium has one neutron, and tritium has two neutrons. These isotopes are represented as follows:

• Protium: ( ^1_1H )
• Deuterium: ( ^2_1H )
• Tritium: ( ^3_1H )

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