Atoms and Elements: Foundations of Atomic Theory

Chapter 2: Atoms and Elements

Introduction

This chapter introduces the fundamental concepts of atomic theory, the historical development of our understanding of atoms, and the laws that govern their behavior. It covers the evolution of atomic models, key experiments, and the basic structure of matter.

Fundamental Laws of Atomic Theory

Nature of Matter

The study of matter seeks to answer whether it is continuous or particulate. Modern chemistry supports the particulate nature of matter, composed of atoms.

• Brownian Motion (1827): Observed by Einstein (1905) and Perrin (1908), provided evidence for the existence of atoms.

• Conservation of Mass (Lavoisier, 1787): In chemical reactions, matter is neither created nor destroyed.

• Law of Definite Proportions (Proust, 1797): All samples of a compound have the same proportions of their constituent elements.

• Law of Multiple Proportions (Dalton, 1804): When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

• Dalton's Atomic Theory (1808): Each element is composed of tiny, indestructible particles called atoms.

Example: Carbon and oxygen can form carbon monoxide (CO) and carbon dioxide (CO2), with the ratio of oxygen masses being a small whole number.

Dalton’s Atomic Theory (1808)

Postulates of Dalton’s Atomic Theory

• Atoms: Each element is composed of tiny, indestructible particles called atoms.

• Identity: All atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.

• Combination: Atoms combine in simple, whole-number ratios to form compounds.

• Rearrangement: In chemical reactions, atoms are not changed into atoms of another element; they only rearrange how they are bonded.

Example: Water (H2O) is always composed of two hydrogen atoms and one oxygen atom.

Early Atomic Models

Dalton’s Billiard Ball Model

Dalton envisioned atoms as solid, indestructible spheres, similar to billiard balls.

• Atoms: Tiny, indestructible particles.

JJ Thomson’s Cathode Ray Tube Experiment (1906 Nobel)

Thomson discovered the electron using cathode ray tubes, showing that atoms contain smaller, negatively charged particles.

• Discovery: Electrons are components of atoms.

• Implication: Atoms are divisible and contain internal structure.

Thomson’s Plum Pudding Model (1904)

Thomson proposed that atoms are composed of electrons embedded in a positively charged "pudding," like chocolate chips in a cookie.

• Model: Electrons scattered within a sphere of positive charge.

Example: The chocolate chip cookie model illustrates electrons (chips) in a positively charged matrix (cookie).

Millikan Oil Drop Experiment

Millikan measured the charge of the electron by observing tiny oil droplets in an electric field.

• Result: Determined the fundamental charge of the electron ($e = 1.602 imes 10^{-19}$ C).

Practice Problems

Stoichiometry and Law of Conservation of Mass

Practice problems help apply the laws of atomic theory to real chemical reactions.

• Example 1: If 7.7 g of Na(s) react with 11.9 g Cl2(g), how many grams of NaCl(s) are formed?

• Example 2: If 10.0 g sample of calcite contains 4.0 g of calcium, how much calcite contains 40.8 g of calcium?

• Example 3: If 24.0 g sample of water contains 23.1 g of oxygen, how much water contains 40.8 g of oxygen?

Application: These problems reinforce the concept of mass relationships and stoichiometry in chemical reactions.

Summary Table: Fundamental Laws of Atomic Theory

Law	Description	Example
Conservation of Mass	Matter is neither created nor destroyed in a chemical reaction.	Mass of reactants equals mass of products.
Law of Definite Proportions	A compound always contains the same proportion of elements by mass.	Water is always 11% hydrogen and 89% oxygen by mass.
Law of Multiple Proportions	When two elements form more than one compound, the ratios of the masses of one element that combine with a fixed mass of the other are small whole numbers.	CO and CO2 have oxygen-to-carbon mass ratios of 1.33:1 and 2.66:1.

Key Terms and Definitions

• Atom: The smallest unit of an element that retains its chemical properties.

• Element: A substance composed of only one type of atom.

• Compound: A substance formed from two or more elements chemically combined in fixed proportions.

• Electron: A negatively charged subatomic particle found in atoms.

• Stoichiometry: The calculation of reactants and products in chemical reactions.

Equations

• Law of Conservation of Mass:

$\text{Total mass of reactants} = \text{Total mass of products}$

• Law of Definite Proportions:

$\frac{\text{Mass of element A in compound}}{\text{Mass of element B in compound}} = \text{constant for all samples}$

• Law of Multiple Proportions:

$\frac{\text{Mass of element A that combines with 1 g of element B in compound 1}}{\text{Mass of element A that combines with 1 g of element B in compound 2}} = \text{small whole number}$

Additional info: Some details and examples have been expanded for clarity and completeness.