4: Atoms and Elements
- Page ID
- 47418
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- 4.1: Cutting Aluminum until you get Atoms
- This page explores the idea of infinitely dividing aluminum foil, linking it to ancient philosophical debates on matter, particularly Democritus' theory that matter has a limit to its divisibility. It also explains that most elements comprise singular atoms, while some, such as hydrogen and oxygen, naturally occur as diatomic molecules.
- 4.2: Indivisible - The Atomic Theory
- This page outlines the evolution of atomic theory, starting with Democritus's idea of indivisible particles, which was initially dismissed by Greek philosophers. It highlights John Dalton's 19th-century revival of the theory through experimentation, establishing that atoms are the fundamental units of matter, indivisible, and combine in specific ratios. Dalton's principles, though modified, continue to be essential for understanding chemical behavior today.
- 4.3: The Nuclear Atom
- This page covers the evolution of atomic theory, detailing J.J. Thomson's discovery of the electron and the "plum pudding" model. It also discusses Rutherford's gold foil experiment, which contradicted this model and introduced the nuclear model of the atom, featuring a dense nucleus with orbiting electrons. The page emphasizes atomic structure, including protons, neutrons, and the forces that maintain the nucleus.
- 4.4: The Properties of Protons, Neutrons, and Electrons
- This page explains the three main subatomic particles: electrons (negatively charged, light), protons (positively charged), and neutrons (neutral). It details their locations within an atom, with electrons orbiting outside the nucleus and protons and neutrons within it, contributing to atomic mass. The concept of atomic mass units (amu) is introduced, highlighting the significance of balanced numbers of protons and electrons for electrical neutrality in atoms.
- 4.5: Elements- Defined by Their Number of Protons
- This page covers key concepts of atomic structure, detailing atomic number and mass number as crucial identifiers of elements. It explains that the atomic number defines proton count, while mass number reflects the total of protons and neutrons. The relationship between protons and electrons in neutral atoms is highlighted, along with the introduction of isotopes as variations of elements based on neutron differences.
- 4.6: Looking for Patterns - The Periodic Table
- This page explains the periodic table's organization, emphasizing its significance in chemistry. Elements are arranged by atomic number in groups (1-18) and periods, with classifications as metals, nonmetals, and metalloids. It highlights key groups such as alkali metals, alkaline earth metals, halogens, and noble gases, detailing their traits and common compounds.
- 4.7: Ions - Losing and Gaining Electrons
- This page describes two types of ions: cations, which are positively charged and formed by atoms losing electrons, and anions, which are negatively charged and formed by atoms gaining electrons. Examples include sodium (Na⁺) and chlorine (Cl⁻). The periodic table is useful for predicting ion charges based on group similarities, and the page also covers the naming conventions for these ions.
- 4.8: Isotopes - When the Number of Neutrons Varies
- This page provides an overview of isotopes, detailing their definition as variations of elements with the same number of protons but differing neutron counts, which influence atomic mass. It covers stable and radioactive isotopes, including examples like hydrogen and carbon-14. The significance of isotopes in Dalton's Atomic Theory and the concept of average atomic mass are highlighted.
- 4.9: Atomic Mass - The Average Mass of an Element’s Atoms
- This page covers atomic mass as the weighted average of an element's isotopes, detailing how to calculate it using relative abundance and isotope masses, with examples provided for boron and neon. It highlights the periodic table's listing of atomic mass as a decimal and differentiates atomic mass from atomic number, noting that atomic mass is usually larger and not a whole number. The page concludes with a practice exercise for calculating the atomic mass of chlorine.