2.2 Evolution of Atomic Theory | Chemistry: Atoms First 2e

Learning Objectives

Outline milestones in the development of modern atomic theory including Thomson, Millikan, and Rutherford experiments
Describe the three subatomic particles (protons, neutrons, electrons) and define isotopes

Core Concepts & Principles

Electrons and Cathode Rays

In the late 1800s, J. J. Thomson used a cathode ray tube to discover that high-voltage electrical discharges in low-pressure gases produce a visible beam. This beam is deflected by electric and magnetic fields, showing it consists of negatively charged particles—electrons—that are fundamental constituents of all atoms and far less massive than atoms. Thomson calculated their charge-to-mass ratio.

Elementary Charge and Mass of the Electron

In 1909, Robert Millikan measured the charge on microscopic oil droplets, discovering that charges are always multiples of $1.602 \times 10^{-19}\text{ C}$ (the elementary charge of a single electron). Combining this with Thomson's charge-to-mass ratio yielded the mass of the electron ( $9.107 \times 10^{-31}\text{ kg}$ ).

Early Atomic Models

Thomson proposed the "plum pudding" model (negative electrons embedded in a positive sphere), while Hantaro Nagaoka proposed a Saturn-like model (positive sphere surrounded by an electron ring). Both preceded the modern nuclear view.

Rutherford Scattering and the Nucleus

Ernest Rutherford, Hans Geiger, and Ernest Marsden fired high-speed alpha particles ( $\alpha$ particles)—positively charged particles consisting of two protons and two neutrons—at thin gold foil. While most passed straight through (revealing that atoms are mostly empty space), a tiny fraction deflected sharply, showing that atoms contain a tiny, massive, positively charged nucleus. Rutherford later identified the proton as the hydrogen nucleus.

Isotopes and Neutrons

Frederick Soddy discovered isotopes—atoms of the same element that differ in mass due to varying numbers of neutrons. James Chadwick discovered uncharged neutrons in 1932, explaining why isotopes have essentially the same chemical properties (the same electron arrangement in neutral atoms) yet differ in mass.

Core Mental Model: The Nuclear Atom

Atoms are not indivisible; they consist of an extremely dense, positively charged nucleus (protons and neutrons) surrounded by a vast region of space containing low-mass electrons.

Problem-Solving Routines & Methods

Finding Electron Mass

1
Use the measured electron charge magnitude and charge-to-mass ratio.
2
Divide charge by charge-to-mass ratio: m = e / (e/m).
3
Cancel coulombs to leave kilograms.

Calculation of Electron Mass

\text{Mass}_{\text{electron}} = 1.602 \times 10^{-19}\text{ C} \times \frac{1\text{ kg}}{1.759 \times 10^{11}\text{ C}} = 9.107 \times 10^{-31}\text{ kg}

Combines Millikan's elementary charge with Thomson's charge-to-mass ratio.

Variables & Constants

\text{C}

=Coulombs;

\text{kg}

=kilograms

Practice & Concept Checks

Concept Check

What did the vast majority of alpha particles passing undeflected through gold foil reveal about atomic structure?

Concept Check

How do isotopes of the same element differ structurally, and why are their chemical properties very similar?

Key Terms & Vocabulary

alpha particle (α particle)Subatomic Particles

positively charged particle consisting of two protons and two neutrons

electronSubatomic Particles

negatively charged, subatomic particle of relatively low mass located outside the nucleus

isotopesAtomic Structure

atoms that contain the same number of protons but different numbers of neutrons

neutronSubatomic Particles

uncharged, subatomic particle located in the nucleus

nucleusAtomic Structure

massive, positively charged center of an atom made up of protons and neutrons

protonSubatomic Particles

positively charged, subatomic particle located in the nucleus