Atoms: Early Atomic Theory, Laws, and the Structure of Matter

Atoms and the Structure of Matter

Early Ideas About the Building Blocks of Matter

The concept of atoms as the fundamental building blocks of matter has evolved over centuries. Early philosophers and scientists contributed to our understanding of atomic structure and the nature of matter.

• Leucippus and Democritus (5th century B.C.E.): Proposed that matter is composed of small, indestructible particles called atoms.

• Democritus stated: “Nothing exists except atoms and empty space; everything else is opinion.”

• He suggested that atoms exist in different shapes and sizes and move randomly through empty space.

• Plato and Aristotle: Rejected atomic theory, believing matter had no smallest parts and was composed of varying proportions of four elements: fire, air, earth, and water.

John Dalton and the Atomic Theory

In the early 19th century, John Dalton provided experimental evidence supporting the atomic theory, laying the foundation for modern chemistry.

• Dalton’s Atomic Theory:

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

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

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

  • Atoms of one element cannot change into atoms of another element; in chemical reactions, atoms only change the way they are bound together.

Fundamental Laws of Chemistry

Law of Conservation of Mass

This law states that mass is neither created nor destroyed in a chemical reaction. The total mass remains constant during any chemical or physical change.

• Mass can change forms (solid, liquid, gas), but the total mass remains the same.

• Chemical change: A new compound forms (e.g., burning wood).

• Physical change: The state changes but the substance remains the same (e.g., ice melting to water).

Equation:

$\text{Total mass before reaction} = \text{Total mass after reaction}$

Law of Definite Proportions (Law of Constant Composition)

Proposed by Joseph Proust in 1797, this law states that all samples of a given compound have the same proportions of their constituent elements, regardless of their source or preparation method.

• For example, water ($\mathrm{H_2O}$) always into hydrogen and oxygen in a mass ratio of 1:8 (hydrogen:oxygen).

Example: Decomposition of 18.0 g of water yields 16.0 g of oxygen and 2.0 g of hydrogen. The mass ratio is:

$\text{Mass ratio} = \frac{16.0\ \mathrm{g\ O}}{2.0\ \mathrm{g\ H}} = 8:1$

Law of Multiple Proportions

Formulated by John Dalton in 1804, this law states that when two elements (A and B) form more than one compound, the masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers.

• For example, carbon and oxygen form both carbon monoxide (CO) and carbon dioxide (CO2).

• In CO, the mass ratio of oxygen to carbon is 1.33:1; in CO2, it is 2.67:1. The ratio of these two ratios is 2:1, a small whole number.

Equation:

If compound 1: $\frac{m_B}{m_A} = x$ If compound 2: $\frac{m_B'}{m_A} = y$ Then $\frac{x}{y}$ is a small whole number.

Comparison of the Laws

• Law of Definite Proportions: Applies to multiple samples of the same compound.

• Law of Multiple Proportions: Applies to different compounds made from the same elements.

Discovery of Subatomic Particles

Discovery of the Electron: J.J. Thomson’s Cathode Ray Experiment

J.J. Thomson discovered the electron through experiments with cathode rays.

• Cathode rays are streams of negatively charged particles (later called electrons) that travel from the cathode to the anode in a vacuum tube.

• Properties of cathode rays:

  • Travel in straight lines

  • Independent of the material of the cathode

  • Carry a negative charge

• Thomson measured the charge-to-mass ratio of the electron as $-1.76 \times 10^8$ C/g.

Millikan’s Oil Drop Experiment: Determining the Charge of an Electron

Robert Millikan measured the fundamental charge of the electron using the oil drop experiment.

• He balanced the gravitational and electric forces on tiny charged oil droplets.

• Determined the charge of a single electron as $-1.60 \times 10^{-19}$ C.

• Using Thomson’s charge-to-mass ratio, the mass of the electron was calculated as $9.10 \times 10^{-28}$ g.

Equation:

$\text{mass of electron} = \frac{\text{charge of electron}}{\text{charge-to-mass ratio}}$

Structure of the Atom: Early Models

• Thomson’s Plum-Pudding Model: Electrons are embedded within a positively charged sphere, like plums in a pudding.

• Rutherford’s Gold Foil Experiment (1909): Alpha particles were directed at a thin gold foil. Most passed through, but some were deflected or bounced back, indicating a small, dense, positively charged nucleus.

Nuclear Theory of the Atom

• Most of the atom’s mass and all of its positive charge are contained in a small core called the nucleus.

• Most of the atom’s volume is empty space, with electrons dispersed throughout.

• The number of negatively charged electrons equals the number of positively charged protons, making the atom electrically neutral.

Discovery of the Neutron

• James Chadwick (1932) discovered the neutron, a neutral particle in the nucleus with a mass similar to that of a proton.

• Neutrons account for the previously unexplained mass in the nucleus.

Subatomic Particles: Properties

Atoms are composed of three fundamental subatomic particles: protons, neutrons, and electrons.

Particle	Symbol	Mass (kg)	Charge (C)
Proton	p+	1.67262 × 10-27	+1.60218 × 10-19
Neutron	n	1.67493 × 10-27	0
Electron	e-	9.10938 × 10-31	-1.60218 × 10-19

Key Points:

• Protons and neutrons have nearly identical masses, much greater than that of electrons.

• Protons and electrons have equal but opposite charges; neutrons are neutral.

Summary Table: Comparison of Atomic Laws

Law	Applies To	Description
Law of Conservation of Mass	All chemical and physical changes	Mass is neither created nor destroyed.
Law of Definite Proportions	Multiple samples of the same compound	All samples have the same element ratio by mass.
Law of Multiple Proportions	Different compounds of the same elements	Element mass ratios are small whole numbers.

Example Application: When a log burns, the mass of the ash is less than the original log because much of the matter is released as gases into the air, demonstrating the law of conservation of mass.