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4: Atomic Structure

Last updated

Mar 21, 2025
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- 3.18: Significant Figures in Multiplication and Division
- 4.1: Democritus' Idea of the Atom

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4.1: Democritus' Idea of the Atom
This page discusses the philosophical debate on the nature of matter between ancient Greeks Aristotle and Democritus. Aristotle believed matter could be divided infinitely, while Democritus introduced the idea of atoms as indivisible particles. Although Democritus's concepts lacked experimental support and were overshadowed by Aristotle's views for centuries, they eventually became foundational to modern science, taking about 2,000 years for acceptance.
4.2: Law of Conservation of Mass
This page discusses the law of conservation of mass, which asserts that in a chemical reaction, the mass of products equals the mass of reactants, implying mass cannot be created or destroyed. Established in the late 1700s through quantitative analysis, an example involving sodium chloride and silver nitrate demonstrates this principle, showing that their combined product mass matches the initial mass of reactants.
4.3: Law of Multiple Proportions
This page contrasts unicycles and bicycles regarding their design and material differences, then explains the law of multiple proportions with respect to carbon monoxide and carbon dioxide, focusing on their varying toxicities and health impacts. The discussion wraps up with review questions pertaining to the law and the compounds explored.
4.4: Law of Definite Proportions
This page discusses the importance of electricity for daily tasks and the necessity of a stable voltage supply. It explains the law of definite proportions, emphasizing that chemical compounds, like water, maintain fixed elemental ratios by mass. This principle aids in predicting chemical behavior and reactions, illustrated by the example of carbon dioxide production during combustion.
4.5: Mass Ratio Calculation
This page explains the law of multiple proportions in chemistry, illustrating it with examples of copper and chlorine compounds that have different copper-to-chlorine mass ratios. It details the calculations for comparing copper quantities that combine with a fixed mass of chlorine, emphasizing the significance of these ratios in understanding compound composition and encouraging further exploration of mass ratios and molecular formulas.
4.6: Dalton's Atomic Theory
This page outlines the evolution of scientific thought on matter's composition, highlighting the debates preceding the 19th century. It emphasizes John Dalton's advancements in atomic theory, which posited that matter is made of indivisible atoms and that atoms of the same element are identical. Dalton's 1804 theory established foundational principles for modern chemistry, although some aspects, like atomic indivisibility, have been updated.
4.7: Atom
This page explains that atoms are the smallest particles of an element, consisting of protons, neutrons, and electrons. Each element has a unique number of protons, and atoms are electrically neutral due to the equal number of protons and electrons. Although they are generally small, their sizes differ among elements. The helium atom model is used to illustrate the basic atomic structure, featuring a nucleus of protons and neutrons, with electrons orbiting around it.
4.8: Electrons
This page explores the causes of power outages and the evolution of atomic theory, particularly highlighting J.J. Thomson's work on electrons. It details how power outages disrupt electricity flow and describes Thomson's experiments with cathode ray tubes, which identified cathode rays as streams of electrons. His findings, including their deflection by magnetic fields and a consistent charge-to-mass ratio across substances, were pivotal in advancing atomic theory.
4.9: Protons
This page explores the difficulties scientists encounter when explaining invisible entities like atoms versus visible objects. It covers the discovery of electrons via cathode rays and Eugene Goldstein's identification of protons in cathode ray tubes. The text underscores the relationship between protons and electrons in hydrogen atoms and the importance of cumulative research in advancing scientific knowledge of atomic structures.
4.10: Neutrons
This page discusses Sherlock Holmes, a fictional detective by Sir Arthur Conan Doyle, known for his deductive reasoning. It then shifts to the discovery of neutrons, detailing how earlier atomic models failed to reconcile atomic number and weight. Rutherford's theory of extra particles and the subsequent identification of neutrons by German researchers and James Chadwick are outlined.
4.11: Cathode Ray Tube
This page outlines the history and importance of cathode ray tubes (CRTs) in television technology, detailing early contributions from Heinrich Geissler and Sir William Crookes. It emphasizes that CRTs revealed cathode rays as streams of particles with mass and mentions innovations like Karl Ferdinand Braun's oscilloscope and Wilhelm Roentgen's discovery of X-rays that emerged from CRT research, underscoring their significant influence on science and technology.
4.12: Oil Drop Experiment
This page discusses Robert Millikan's oil drop experiment conducted between 1908 and 1917, which aimed to measure the charge of an electron. By using oil droplets and electrical charges, Millikan determined the electron's charge to be about $1.5924 \times 10^{-19} \: \text{C}$ , aligning closely with current values. His work also contributed to the understanding of atomic structure by indicating the presence of another positively charged particle.
4.13: Plum Pudding Atomic Model
This page discusses the evolution of model construction, transitioning from balsa wood to plastics, and how models, such as J.J. Thomson's "plum pudding" model, help visualize concepts like atomic structure. While these models simplify complex ideas for better understanding, they are not functional replicas and may be replaced as scientific knowledge advances, as seen with Rutherford's atomic model. Ultimately, models serve as valuable tools across different scientific disciplines.
4.14: Gold Foil Experiment
This page discusses Rutherford's 1911 gold foil experiment, which challenged the prevailing atomic model by demonstrating that some alpha particles were significantly deflected. This led to the proposal of a nuclear model of the atom, with a dense, positively charged nucleus made up of protons and neutrons, and electrons surrounding it in a cloud, highlighting that most of the atom is empty space. This was a pivotal advancement in atomic theory.
4.15: Atomic Nucleus
This page likens science to a jigsaw puzzle, illustrating how researchers' discoveries enhance our understanding of complex concepts like the atomic nucleus. It traces the evolution of atomic models from Thomson's to Rutherford's, highlighting how the introduction of neutrons clarified the coexistence of protons in the nucleus. This led to the recognition of the strong nuclear force, which binds protons and neutrons despite the repulsion between protons.
4.16: Atomic Number
This page explores individuality through identifiers such as cell phone numbers and DNA, then shifts to atomic theory, explaining how atomic numbers define elements based on proton counts. It emphasizes the organization of the periodic table for predicting element properties and notes that atoms are neutral, with equal numbers of protons and electrons, underscoring the foundational aspect of atomic structure.
4.17: Mass Number
This page explains methods for determining chemical mass and traces the historical development of atomic weight from John Dalton's initial concepts based on hydrogen to modern understandings involving protons and neutrons. It defines mass number as the sum of protons and neutrons, providing examples from the first six elements of the periodic table, including a detailed look at helium, and illustrates how to calculate the number of neutrons based on mass and atomic numbers.
4.18: Isotopes
This page explains that isotopes are variants of the same element with identical proton counts but differing neutron numbers, leading to varied atomic masses. It describes the term "nuclide" as referring to a specific isotope's nucleus. Using carbon as an example, it mentions three natural isotopes: carbon-12, carbon-13, and carbon-14. While isotopes change atomic mass, they do not affect chemical reactivity, which is determined by the number of electrons and protons.
4.19: Atomic Mass Unit
This page highlights the historical importance of standardized measurements in the U.S., particularly in science for consistent data comparison. It establishes the carbon-12 atom as the reference for atomic mass, defining one atomic mass unit as one twelfth of carbon-12's mass. The variations in atomic masses, such as oxygen-16, are discussed due to nucleus interactions, and a mass spectrometer is identified as the tool for measuring these atomic masses.
4.20: Calculating Average Atomic Mass
This page defines atomic mass as the weighted average of an element's isotopes based on their natural abundances, using hydrogen and chlorine as examples. It explains the calculation process for average atomic mass, demonstrating how one isotope can dominate the average. Additionally, it includes key concepts and review questions related to the topic.

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