In-Depth Notes on Elements and Compounds

FROM ELEMENTS TO COMPOUNDS

Module Overview

• The universe is interconnected, highlighting biological, chemical, and atomic relationships among all elements.

• Key topics include the Big Bang, stellar elements, compound formation, and nucleosynthesis.

Compounds

• Uses of Compounds:

  • Serve as catalysts:

  • Example: Alkylation in producing artemisinin (anti-malarial drug) and cosalane (anti-HIV).

Fundamental Concepts

• Atom: Smallest particle of an element retaining its properties.

• Dalton’s Atomic Theory (1808):

  • Elements consist of tiny particles called atoms.

  • All atoms of a given element are identical.

  • Chemical reactions involve a reorganization of atoms without changing them.

Atomic Structure Development

• J.J. Thomson (1897): Discovered the electron using a cathode ray tube.

• Ernest Rutherford (1908):

  • Conducted the Gold Foil Experiment, discovering the nucleus.

  • Found that some elements emit alpha particles.

• 20th-century advancements like mass spectrophotometers led to the discovery of neutrons.

Modern Atomic Model

• Structure of the Atom:

  • Nucleus: Contains protons and neutrons.

  • Extranuclear Region: Holds electrons in orbit around the nucleus.

Periodic Table Insights

• Periodic Trends

  • Atomic Size:

  • Increases down a group due to additional energy levels.

  • Decreases across a period due to increased nuclear charge attracting electrons closer.

  • Ionization Energy: Energy required to remove an electron. Generally increases across a period and decreases down a group.

  • Electronegativity:

  • Ability to attract electrons; increases left to right across a period but decreases down a group. E.g., Li (1.0) < Na (0.9) < K (0.8).

  • Electron Affinity: Measure of energy change when an electron is added to an atom. Most elements release energy when gaining an electron.

Bonding Fundamentals

• Chemical Bonds: Forces holding atoms together.

  • Ionic Bonds: Electron transfer, forming ions that attract each other. E.g., NaCl has no overall charge as it balances cations and anions.

  • Covalent Bonds: Electron sharing between atoms. Each atom often assumes a noble gas configuration for stability.

  • Types of Bonds:

  • Ionic: Transfer of electrons, resulting in charged particles.

  • Covalent: Sharing of electrons, forming molecules.

  • Metallic: Delocalized electrons allow conductivity.

VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory:

  • Used to predict molecular shapes by minimizing electron pair repulsion.

  • Examples of molecular shapes:

  • Tetrahedral (e.g., CH4)

  • Bent (e.g., H2O)

  • Linear (e.g., CO2).

Polarity in Molecules

• Polar vs Nonpolar:

  • Polar: Molecules with dipoles (e.g., water).

  • Nonpolar: Molecules without charged ends (e.g., methane).

• Electronegativity differences help determine bond type:

  • > 2.0 = Ionic

  • < 0.4 = Nonpolar Covalent

  • 0.5 – 1.7 = Polar Covalent.

Summary Statements

• Understanding atomic and molecular structure is essential for grasping chemical behavior.

• Compounds are formed based on the interactions of electrons and the resulting stability from achieving noble gas configuration.

• Remember the essential principles of bonding and how they relate to the periodic trends.