Chapter 2: The Development of Modern Atomic Theory

Chapter 2 Overview

• This chapter will be divided into several videos, more than Chapter 1.
• A significant portion will focus on chemical nomenclature, with three different types discussed in separate, shorter videos.
• This initial section will cover the historical development of modern atomic theory.

Modern Atomic Theory

Modern atomic theory can be summarized by the following principles:

• Matter Composition: Matter is composed of small particles called atoms.
• Elements: An element consists of only one type of atom, possessing a characteristic mass for that element.
• Atomic Properties: Atoms of one element have distinct properties from atoms of all other elements (e.g., chlorine atoms differ chemically from oxygen atoms).
• Compounds: A compound comprises atoms of two or more elements combined in small, whole-number ratios (e.g., carbon dioxide, $CO_2$) and can be broken down into its constituent elements but not smaller than individual atoms.
• Conservation of Atoms: Atoms are neither created nor destroyed in a chemical change or reaction; they are merely rearranged into different compounds (e.g., carbon dioxide to carbon monoxide).

Foundations of Atomic Theory: Historical Laws

Prior to modern atomic theory, understanding was based on a series of fundamental laws:

• Law of Conservation of Mass
  • Principle: In any chemical reaction, mass is neither lost nor gained, only changed in form. For example, converting carbon dioxide to carbon monoxide rearranges the atoms, but the 'extra' oxygen mass is still present, likely as oxygen gas.
  • Modern Version: We now understand this as the Law of Conservation of Mass and Energy, recognizing that mass and energy can be converted into one another (e.g., in a nuclear weapon, as described by Einstein's famous equation $E=mc^2$).
• Law of Definite Proportions
  • Principle: A compound always contains the same elements in the same proportion by mass, regardless of the size or source of the sample.
  • Historical Context: Early scientists didn't know the exact chemical formulas (like $H_2O$ for water) but observed consistent mass ratios when compounds were broken down.
  • Example (Water):
    • To calculate the oxygen-to-hydrogen mass ratio in water, we divide the mass of oxygen by the mass of hydrogen.
    • Sample 1: $20.8704$ g Oxygen $/ 2.63$ g Hydrogen $= 7.9355$
    • Sample 2: $0.1661$ g Oxygen $/ 0.02093$ g Hydrogen $= 7.9360$
    • Sample 3 (Kilograms converted to grams): $34014$ g Oxygen $/ 4286$ g Hydrogen $= 7.9360$
    • Conclusion: Within experimental error, all samples yield an oxygen-to-hydrogen ratio of approximately $7.936$, confirming a definite proportion.
• Law of Multiple Proportions
  • Principle: When two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small whole numbers.
  • Example (Water vs. Hydrogen Peroxide):
    • Water ($H_2O$): The oxygen-to-hydrogen ratio was determined to be $7.936$ grams of oxygen per gram of hydrogen.
    • Hydrogen Peroxide ($H<em>2O</em>2$): Performing similar calculations for hydrogen peroxide yields an oxygen-to-hydrogen ratio of approximately $15.87194$ grams of oxygen per gram of hydrogen.
    • Comparison: Hydrogen peroxide has twice as much oxygen per hydrogen atom compared to water, which is reflected in the ratio.
    • Applying the Law: Divide the larger ratio by the smaller ratio:
      $15.87194 / 7.936 hickapprox 1.9999 hickapprox 2.0$
    • Conclusion: The resulting ratio of approximately $2.0$ (a small whole number) indicates that water and hydrogen peroxide are distinct compounds made from the same elements but with different small whole-number ratios of atoms. This evidence helped lead scientists toward understanding specific chemical formulas like $H<em>2O$ and $H</em>2O_2$.

Early Atomic Theory Development

Combining these laws led to the early development of atomic theory:

• Early Postulates:
  • Atoms are small particles that constitute all matter.
  • Atoms combine in constant ratios for specific compounds.
  • Atoms have specific properties that change only when they form new compounds.
  • Compounds can be broken down to recover the original atoms.
• Discovery of Subatomic Particles:
  • The existence of particles smaller than an atom was discovered, initially through physical means like bombarding atoms with radiation, as chemical methods could not break down atoms.
  • Electrons: Experiments by scientists like J. J. Thomson and Robert Millikan (the oil drop experiment) helped determine the mass and negative charge of the electron.
  • Charge Balance: Since atoms are neutral, the discovery of negatively charged electrons implied the existence of a balancing positive charge within the atom.
• Plum Pudding Model (Thomson):
  • Description: This model proposed that atoms were a diffuse, nebulous sphere of positive charge (like the 'pudding') with individual, negatively charged electrons (like 'plums') scattered throughout it.
  • Analogy: A sphere of gelatinous positive charge with negative charges dispersed within.
• Challenging the Plum Pudding Model: Rutherford's Gold Foil Experiment
  • Purpose: Ernest Rutherford conducted this experiment to confirm the Plum Pudding model.
  • Radiation Types: He utilized alpha particles, which are heavy, positively charged particles (one of three main types of radiation: alpha positive, beta negative, gamma neutral light).
  • Experimental Setup: An alpha source directed a beam of positively charged alpha particles towards an incredibly thin piece of gold foil (chosen for its malleability, allowing it to be hammered thin, ideally close to one atom thick). A photographic film surrounded the gold foil to detect the alpha particles after interaction.
  • Expected Results (based on Plum Pudding Model): Based on the idea of a diffuse positive charge and tiny electrons, Rutherford expected most alpha particles to pass straight through the foil with minimal deflection, causing a very concentrated dark spot on the photographic film, perhaps with a slight, diffuse scattering.
  • Actual Results: While a significant number of alpha particles did pass straight through (creating a strong dark spot), a small but distinct number were deflected at large angles, and some even bounced directly back towards the source. This contradicted the Plum Pudding model, which predicted only minor deflections.
  • Discovery of the Nucleus: The unexpected scattering led Rutherford to propose a new atomic model:
    • Atoms have a tiny, dense, positively charged nucleus at their center.
    • The vast majority of an atom's mass and all of its positive charge are concentrated in this nucleus.
    • Electrons orbit this nucleus at a relatively large distance, making most of the atom empty space.
    • The strong deflections observed occurred when an alpha particle directly hit or passed very close to this dense, positively charged nucleus.
  • Revised Model: The atom transitioned from a diffuse 'plum pudding' to a model with a solid, positive core (the nucleus) with negative charges (electrons) scattered around it rather than within it.
  • Modern Perspective Note: While this was a vast improvement, even this Rutherford model (often depicted as electrons orbiting like planets) is not entirely accurate by modern quantum theory standards. The familiar atomic symbol with orbiting electrons is still wrong, despite its use by some organizations like the International Atomic Energy Agency. A true understanding requires quantum mechanics, which will be discussed in later chapters.

Conclusion of this Section

• This discussion on atomic theory will be resumed when covering quantum theory later in the course.
• The next parts of this chapter will delve into topics like the periodic table.