UNIT 2

Democritus (c. 400 BCE): Proposed that matter is finite and made up of indivisible particles called atoms, relying on reasoning rather than experimental evidence. He laid the foundation for the concept of the atom.
Modern Understanding: Determined that the atom is divisible and consists of smaller subatomic particles. Current research suggests that even subatomic particles can be further divided.
Structure of the Atom:
- Nucleus: Contains protons (positively charged) and neutrons (no charge).
- Electrons: Much lighter than protons and neutrons and orbit the nucleus.
- The majority of an atom’s volume is empty space.

Four Elements: Early philosophers believed matter was composed of four basic elements: earth, air, fire, and water.
Continuous Theory of Matter (Aristotle, Plato): Proposed that matter was infinite and could be divided indefinitely.
Discontinuous Theory of Matter (Democritus): Argued against the continuous theory, stating there is a limit to how matter can be divided, leading to his definition of atoms.

Joseph Louis Proust (1754-1826): Studied mass ratios in chemical compounds, concluding that compounds always contain the same proportions of elements by mass.
- Example: Water (H2O) is consistently 1/9 by mass hydrogen and 8/9 by mass oxygen.

Antoine Lavoisier: Observed that in chemical reactions, the mass of reactants equals the mass of products. This concept transformed chemistry by accounting for gases and highlighting their existence as real matter.

John Dalton: Advanced atomic theory based on previous work, proposing that:
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical, while atoms of different elements vary.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements can combine in simple whole-number ratios to form compounds.
5. In chemical reactions, atoms are rearranged.
Law of Multiple Proportions: If two elements can form multiple compounds, the masses of one element that combine with a fixed mass of the other can be expressed as ratios of whole numbers.

William Crookes (1877): Conducted experiments in vacuum tubes, leading to the discovery of cathode rays, which were later identified as electrons.
J.J. Thomson (1897): Demonstrated that cathode rays could be deflected by magnetic fields, proving the existence of negatively charged particles (electrons).
- Introduced the Plum Pudding Model, proposing that electrons are embedded in a positively charged 'soup'.

Robert Millikan: Measured the charge of an electron, finding it to be approximately 1.602176487 imes 10^{-19} coulombs and determining its mass as 9.10938215 imes 10^{-31} kg.

Ernest Rutherford’s Gold Foil Experiment (1911): Proposed a new atomic model where a small, dense nucleus houses protons, surrounded by electrons. This replaced the Plum Pudding Model due to observed deflections of alpha particles.
James Chadwick (1932): Discovered the neutron, expanding the understanding of atomic structure by showing that the nucleus contains both protons and neutrons.

The Modern Atomic Model recognizes electrons, protons, and neutrons as fundamental particles within the atom, with electrons primarily influencing chemical properties. The experiments of Crookes, Thomson, Rutherford, and Chadwick contributed to this understanding.
Electron Configuration and the Periodic Table: Different groups of the periodic table are determined by the configuration of electrons, influencing chemical behavior.
The Structure of the Periodic Table is linked to the filling order of atomic sublevels (s, p, d, f) and reflects trends in elemental properties.

Elements are organized into periods (horizontal rows) and groups (vertical columns), indicating recurring properties based on electron configuration.
- Atomic Radius: Tends to increase down a group and decrease across a period from left to right.
- Ionization Energy: Increases across a period; decreased down a group due to increased atomic size and electron shielding.
- Electronegativity: Follows similar trends, affecting how an atom attracts shared electrons in bonds.

Ionic compounds consist of cations and anions. The naming convention usually states that the cation name precedes the anion name.
For monatomic ions, cations retain their elemental name, while anions take the suffix -ide.
Common Polyatomic Ions: Include ions like sulfate (SO4^2-) and nitrate (NO3^-).
Acids: Defined as substances that produce H+ ions in aqueous solutions. Binary acids stem from hydrogen and a single nonmetal, while oxyacids consist of hydrogen, oxygen, and another element.

Molecular (covalent) compounds consist of two nonmetals. Their naming system uses prefixes to indicate the number of atoms from each element in the compound.

The evolution of atomic theory showcases a gradual understanding of atomic structure, the relationships between atomic components, and their binding characteristics, paving the way for modern chemistry practices. This foundational knowledge continues to influence ongoing research in chemistry, biochemistry, and materials science.

Empirical formula: A formula representing the simplest whole-number ratio of different elements in a compound.
Binary ionic compounds: Compounds formed from two different elements, typically including a metal and a nonmetal, bonded ionically.
Ternary ionic compounds: Compounds composed of three or more elements that include at least one polyatomic ion.
Monatomic ions: Ions made up of a single atom that carries a charge due to the loss or gain of electrons.
Polyatomic ions: Ions formed from two or more atoms that are covalently bonded and carry a charge.
Oxyanions: Polyatomic ions that contain oxygen along with another element.
Covalent bond: A type of chemical bond formed when two atoms share pairs of electrons.
Molecule: A structure consisting of two or more atoms that are covalently bonded together.
Binary molecular compound: A compound comprised of two nonmetals that are connected by covalent bonds.
Molecular formula: A representation revealing the exact number of each type of atom in a molecule.
Acid: A compound that releases H+ ions when dissolved in water.
Binary acids: Acids formed by hydrogen combined with a single nonmetal.
Oxoacids: Acids containing hydrogen, oxygen, and another element.
Base: A substance that can accept protons or donate electron pairs during chemical reactions.

Joseph Priestley: An early chemist known for his discovery of oxygen and contributions to the understanding of gases.
Alfred Stock: A chemist recognized for his work on coordination compounds and the development of chemical nomenclature.

Understanding the formation and naming of ionic and molecular compounds: Focus on the methods for properly identifying and naming different types of compounds.
The characteristics of acids and bases, and naming conventions for each type: Emphasizes the attributes that define acids and bases and outlines how to name them accordingly.