Study Notes for General Chemistry 1

General Chemistry 1 Course Material

ATOMS AND ELEMENTS
Learning Objectives

  • Understand the fundamental laws of matter:

    • Law of Conservation of Mass: Mass is neither created nor destroyed during chemical reactions.

    • Law of Constant Composition: A chemical compound always contains its component elements in fixed ratio by mass.

    • Law of Multiple Proportions: When elements combine, they do so in the ratio of small whole numbers.

  • Dalton’s Atomic Theory: Explanation of how these laws led to the theory.

  • Comparison of Atomic Concepts: Distinguish atomic number, mass number, and isotopes; understanding their roles in defining elements.

  • Components of Matter: Differentiate among atoms, molecules, and ions with examples.

Module Contents

  • Learning Outcomes

  • Pre-activity: Recall concepts related to atoms.

  • Fundamentals of Atoms

  • Subatomic Particles

  • Electron Configuration

  • Isotopes

  • Lesson Summary

  • References

Fundamentals of Atoms

  • Definition: Atoms are the fundamental building blocks of all matter.

  • Structure of an Atom:

    • Dense positively charged nucleus composed of protons and neutrons.

    • Surrounding the nucleus is a cloud of negatively charged electrons.

  • Atomic Number:

    • The number of protons in the nucleus; it defines the element's identity and position on the Periodic Table.

  • Physical Size: Atoms are incredibly small, measuring in tenths of a nanometer, and are mostly empty space.

  • Historical Background:

    • Leucippus and Democritus (around 420 BCE) proposed that:

    • Matter consists of indivisible particles called atoms.

    • Atoms vary in size and shape.

    • Atoms exist in voids allowing for their movement.

    • Atoms are in constant motion.

Focus Questions

  • How are atoms formed?

  • What are the contributions of philosophers and scientists to atomic theory?

  • What are the components of an atom?

  • How to determine atomic number and mass?

Historical Development of Atomic Theory (17th to Mid-19th Century)

  • Scientific Revolution: Recognized as beginning in the seventeenth century, characterized by experimental science.

  • Sir Francis Bacon: Father of empiricism; popularized inductive methodologies, leading to the scientific method.

  • Robert Boyle:

    • Quantitatively studied gas behavior, contributing to the definition of an element as a substance that cannot be broken down further.

  • Antoine-Laurent Lavoisier:

    • Conducted quantitative experiments that refined the definition of an element and supported the atomic theory.

  • John Dalton: Introduced operational perspective on atoms in 1803, setting the groundwork for modern physical science.

Dalton's Atomic Theory Postulates

  • Matter comprises extremely small, indestructible atoms.

  • All atoms of a given element are identical.

  • Atoms combine to form compounds but remain unchanged.

  • Atoms combine in small numerical ratios.

Progress in Atomic Structure (Mid-19th Century to Present)

  • Emerging Evidence:

    • Development from Dalton’s theory to various models including Thomson's, Rutherford's, Bohr’s, and Quantum Mechanical models.

Discovery of Electrons (J.J. Thomson)

  • Cathode Rays: Discovered negatively charged particles (electrons) which contradicted the indivisibility postulate.

  • Robert Millikan: Measured the charge of an electron (1.60x10⁻¹⁹ coulomb) through the oil-drop experiment, which demonstrated its mass (9.11x10⁻²⁸ grams).

Rutherford's Nuclear Model

  • Gold Foil Experiment: In 1911, tested Thomson’s model and concluded:

    • Atoms are mostly empty space.

    • Atoms contain a small dense nucleus that is positively charged.

    • Nuclei are surrounded by electrons, confirming the presence of protons.

Neutrons and Atomic Properties

  • James Chadwick: Discovery of neutrons (1932), particles with no charge, leading to the atomic mass definitions.

Atomic Identity and Structure

  • Atomic Number (Z):

    • Defined by the number of protons; determines the identity of an element.

  • Mass Number (A): Calculated by protons and neutrons, affects atomic structure.

Electronic Structure of Atoms

  • Electron Configuration: Governs the chemical properties of the atom by distributing electrons according to specific rules.

  • Key Principles:

    • Aufbau Principle: Electrons fill lower energy orbitals before higher ones.

    • Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers; each orbital can hold a maximum of two electrons with opposite spins.

    • Hund's Rule: Electrons occupy degenerate orbitals singly with parallel spins before pairing.

Isotopes

  • Definition: Variations of an element with the same number of protons but different numbers of neutrons.

  • Characteristics:

    • Identical chemical behavior due to the same electronic configuration.

    • Different mass numbers due to varying neutron counts.

  • Notation: Represented in nuclear or name-mass notation.

Lesson Summary

  • Development of Atomic Theory: From early Greek philosophies through the experimental science age to modern atomic understanding through progressive models.

  • Subatomic Particles: Atoms consist of protons, neutrons, and electrons, with defined roles for each.

  • Electronic Configuration: Arranged according to energy principles; uses orbital diagrams for visual representation.

  • Isotopes Defined: Atoms with the same protons but varied neutrons; relevant for understanding chemical identity and properties.

References

  • Chang, R., & Overby, J. (2021). Chemistry.

  • Silberberg, M. S. (2007). Principles of general chemistry.

  • Tro, N. J. (2018). Introductory Chemistry in SI Units. Pearson.