2.2 Evolution of Atomic Theory

History and Development of Atomic Theory

Atomic Theory Through History

Around 400 BCE, the Greek philosopher Democritus proposed that all matter is made of tiny, indivisible particles he called atomos (meaning "uncuttable"). This was purely philosophical, though. There were no experiments to back it up, and the idea was largely set aside for over two thousand years.

John Dalton revived and formalized atomic theory in the early 1800s with a set of postulates based on experimental evidence:

• All matter is composed of atoms, the fundamental building blocks of matter.
• Atoms of the same element are identical in mass, size, and chemical behavior.
• Atoms of different elements have different properties.
• Atoms cannot be created, divided, or destroyed. (We now know this isn't strictly true for nuclear reactions, but it holds for chemical reactions.)
• Compounds form when atoms combine in simple whole-number ratios (for example, water is always $H_2O$, and carbon dioxide is always $CO_2$).

J.J. Thomson discovered the electron in 1897 using cathode ray tubes. He showed that atoms contain negatively charged particles, which meant atoms are not indivisible after all. To explain how a neutral atom could contain negative charges, he proposed the "plum pudding" model: electrons scattered throughout a uniform sphere of positive charge, like raisins in a pudding.

Ernest Rutherford's gold foil experiment (1911) overturned Thomson's model. Here's what happened:

1. Rutherford's team fired positively charged alpha particles at a very thin sheet of gold foil.
2. Most alpha particles passed straight through, suggesting the atom is mostly empty space.
3. A small fraction deflected at large angles, and a few even bounced straight back.
4. Rutherford concluded that the atom's positive charge and nearly all its mass are concentrated in a tiny, dense core: the nucleus.

This gave us the nuclear model of the atom, with a small, dense, positively charged nucleus surrounded by electrons.

Niels Bohr refined this picture in 1913. His model placed electrons in fixed circular orbits at specific energy levels around the nucleus. Each orbit corresponds to a particular energy, and electrons can jump between levels by absorbing or emitting specific amounts of energy. This successfully explained the discrete line spectrum of hydrogen.

The modern quantum mechanical model replaced Bohr's fixed orbits with orbitals, which are regions of space where an electron is most likely to be found. Electrons behave as both particles and waves (wave-particle duality), and their exact position and momentum cannot be known simultaneously (the Heisenberg uncertainty principle).

Key Contributors to Atomic Understanding

Robert Millikan deserves a mention alongside Thomson and Rutherford. His oil drop experiment (1909) measured the charge of a single electron by balancing gravitational and electric forces on tiny charged oil droplets. This gave us the electron's charge ($1.602 \times 10^{-19}$ C) and, combined with Thomson's charge-to-mass ratio, allowed calculation of the electron's mass.

Quantum Mechanical Model and Atomic Structure

• Energy levels: Electrons exist in discrete (quantized) energy states within an atom. They can only occupy certain allowed energies, not anything in between.
• Valence electrons: The electrons in the outermost energy level. These are the ones responsible for chemical bonding and reactivity.
• Atomic spectra: Each element emits or absorbs a unique pattern of light wavelengths. These patterns act like fingerprints for identifying elements.
• Spectroscopy: The technique of analyzing these light patterns to determine what elements are present in a sample.
• Uncertainty principle: You cannot precisely know both the position and momentum of an electron at the same time. This is why we talk about probability regions (orbitals) rather than exact paths.

Subatomic Particles and Isotopes

Properties of Subatomic Particles

Every atom is built from three types of subatomic particles. Their properties are worth memorizing:

• Protons carry a positive charge ($+1$) and sit inside the nucleus. Each proton has a mass of about $1.673 \times 10^{-27}$ kg, which equals roughly $1$ atomic mass unit (amu). The number of protons defines which element an atom is.
• Neutrons have no charge and also sit inside the nucleus. Their mass is nearly the same as a proton (about $1.675 \times 10^{-27}$ kg, or $\approx 1$ amu). Neutrons help stabilize the nucleus; without enough of them, the positively charged protons would repel each other and the nucleus would fly apart.
• Electrons carry a negative charge ($-1$) and occupy orbitals around the nucleus. They're extremely light compared to protons and neutrons: about $9.109 \times 10^{-31}$ kg, or roughly $\frac{1}{1836}$ amu. Electrons determine an atom's chemical behavior and how it bonds with other atoms.

Isotopes and Chemical Significance

Isotopes are atoms of the same element that have different numbers of neutrons. They have the same atomic number ($Z$, the number of protons) but different mass numbers ($A$, the total of protons + neutrons).

Isotope notation looks like this: $^{A}_{Z}X$

• $X$ = chemical symbol
• $A$ = mass number (protons + neutrons)
• $Z$ = atomic number (protons)

Some common examples:

• Carbon: $^{12}_{6}C$ and $^{13}_{6}C$ are both stable, while $^{14}_{6}C$ is radioactive and used in carbon dating.
• Hydrogen: $^{1}_{1}H$ (protium, the most common), $^{2}_{1}H$ (deuterium, stable but rare), $^{3}_{1}H$ (tritium, radioactive).
• Uranium: $^{235}_{92}U$ is fissile (used in nuclear reactors and weapons), while $^{238}_{92}U$ is the most naturally abundant uranium isotope.

Why do isotopes matter?

• Isotopes of the same element behave almost identically in chemical reactions because they have the same number of electrons and protons. Their physical properties differ, though, because of the mass difference.
• Practical applications include radiometric dating ($^{14}C$ for archaeological samples), medical imaging (radioactive tracers in PET scans), and nuclear energy ($^{235}U$ fission).
• The atomic mass listed on the periodic table is a weighted average of the masses of all naturally occurring isotopes of that element, accounting for how abundant each isotope is in nature.