Chemistry Chapter: Atomic History and Structure

History of the Atom

• Democritus proposed the concept of indivisible particles called "atomos," leading early atomic theory.
• John Dalton formulated the atomic theory in the early 1800s, introducing ideas of atoms as discrete units, compounds as combinations of atoms, and conservation of mass.
• J.J. Thomson discovered the electron in 1897, leading to the "plum pudding" model, which suggested atoms contained positive and negative charges.
• Ernest Rutherford conducted the gold foil experiment in 1911, revealing that atoms have a dense nucleus, leading to the nuclear model of the atom.
• Niels Bohr introduced the Bohr model in 1913, proposing that electrons orbit the nucleus in fixed paths, quantized energy levels.

Key Terms

• Cation: a positively charged ion formed by losing one or more electrons.
• Anion: a negatively charged ion formed by gaining one or more electrons.
• Valence electrons: electrons in the outermost shell, responsible for chemical bonding.
• Octet rule: atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons, promoting stability.

Bohr-Rutherford and Lewis Dot Diagrams

• Bohr-Rutherford diagrams illustrate electron arrangements around the nucleus, showing energy levels and orbits.
• Lewis dot diagrams depict valence electrons as dots around the chemical symbol, highlighting bonding potential and electron sharing.

Chart with Ions

• Ions are represented with their charge and the number of electrons compared to their neutral atom version.
• Common ions include Na⁺ (sodium cation), Cl⁻ (chloride anion), Ca²⁺ (calcium cation), and O²⁻ (oxide anion).

Isotopes and Radioisotopes

• Isotopes are variants of an element with the same number of protons but different neutron counts, affecting atomic mass.
• Radioisotopes are unstable isotopes that decay over time, emitting radiation and transforming into different elements.

Nuclear Reactions

• Three types of nuclear reactions include:
  • Fusion: Combining small nuclei to form a heavier nucleus, releasing energy.
  • Fission: Splitting a heavy nucleus into smaller nuclei, also releasing energy; utilized in nuclear power.
  • Radioactive decay: Spontaneous transformation of an unstable nucleus into a more stable form while emitting particles or electromagnetic radiation.

Isotopic Abundance

• Isotopic abundance refers to the relative proportion of each isotope of an element in nature.
• The weighted average mass incorporates the mass of isotopes and their respective abundances, used to determine the average atomic mass of an element.

Groups on the Periodic Table

• Elements are organized into groups (columns) on the periodic table, indicating similar chemical properties.
• Key groups include alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), and noble gases (Group 18).

Periodic Trends

• Atomic radius tends to decrease across a period (left to right) and increase down a group (top to bottom).
• Ionization energy increases across a period and decreases down a group; it's the energy required to remove an electron from an atom.
• Electron affinity generally increases across a period and varies in groups, indicating the energy change when an electron is added to an atom.

This quiz covers the key concepts of atomic theory, exploring the contributions of scientists such as Dalton, Thomson, Rutherford, and Bohr. It includes definitions of important terms like cations, anions, and the octet rule, as well as Bohr-Rutherford and Lewis dot diagrams. Additionally, you'll learn about isotopes, nuclear reactions, and periodic trends.