Chapter Four: Atoms and Elements

Democritus (Ancient Greece):
- Proposed the original atomic theory; material can be divided into smaller parts until reaching "atomos" (undivided).
- Idea lost to Aristotle’s four-element theory (earth, air, water, fire).
- Aristotle's theory was popular for nearly 2000 years.

Origins in Ancient Egypt:
- Egyptian practices related to life after death likely contributed to early chemical knowledge.
- Interaction of Greek philosophy with Egyptian beliefs resulted in alchemy (Greek word "chemia").
Contribution of Arab scholars in the development of alchemy (added "al" to "kemiya").
- Attributed to black land (fertile soil).
Goals included seeking immortality and transmutation of base metals.

Robert Boyle (1661):
- Wrote The Skeptical Chemist.
- Advocated for experimental method; elements made of corpuscles.
Antoine Lavoisier (1785):
- Established the law of conservation of matter.
- Example: Combustion of hydrogen and oxygen yields water, maintaining mass.
Joseph Proust (1799):
- Proposed the law of definite proportions; compounds contain elements in fixed ratios.
- Example ratio in water: 8 grams of oxygen for every 1 gram of hydrogen.
John Dalton (1803):
- Developed first atomic theory with key postulates:
1. Elements consist of indivisible atoms.
2. Atoms of the same element are identical.
3. Atoms are neither created nor destroyed, only rearranged.
4. Law of multiple proportions: atoms combine in small whole number ratios.
- Example of water: represented as $ext{H}_2 ext{O}$ (2 hydrogen, 1 oxygen).

Nature of Charge:
- Charged particles: positive (protons) and negative (electrons).
- Nature of oppositely charged particles — attraction, like charges — repulsion.
- Concept of electrically neutral atoms: equal numbers of protons and electrons.

J.J. Thomson (1897):
- Discovered electrons using a cathode ray tube; proposed the Plum Pudding model (electrons in a positively charged sphere).
Ernest Rutherford (1909-1911):
- Conducted the gold foil experiment; discovered the nucleus.
- Concluded atoms consist of a dense nucleus of protons surrounded by mostly empty space, electrons orbiting in a planetary model.
James Chadwick (1932):
- Discovered neutrons; pivotal in completing the model of the atom with protons and neutrons in the nucleus.

Structure of Atoms:
- Nucleus: contains protons and neutrons; diameter $ext{10}^{-13} ext{cm}$
- Electrons: orbit at approximately $ext{10}^{-8} ext{cm}$ from the nucleus; contribute negligibly to mass.
- Overall atomic mass from protons and neutrons (99% or more).

Defined by the number of protons; atomic number ( ext{Z}) indicates number of protons.

Definition:
- Mass of a proton/neutron ≈ 1 atomic mass unit (amu); exact definition: 1/12 of a carbon-12 atom.

Chemical symbols derived from Latin or English; often contain one or two letters (e.g., Na for sodium).
Atomic number (e.g., 19 for potassium) indicates number of protons; written at the top.
Atomic mass (weighted average) appears below the symbol; not identical to mass number.

Each element has its unique symbol, with distinct categories:
- Metals: Left of the staircase on the table.
- Nonmetals: Right of the staircase, with hydrogen as an exception.
- Metalloids: Located on the staircase line, exhibiting properties of both metals and nonmetals.
Major Groups:
- Alkali metals (Group 1): Very reactive, soft metals.
- Alkaline earth metals (Group 2): Harder, less reactive.
- Halogens (Group 17): Very reactive nonmetals.
- Noble gases (Group 18): Inert and non-reactive gases.

Elements are organized by increasing atomic number.
Each row is called a period, and each column is known as a group.
Similar chemical and physical properties present among elements in the same group.

Definition of Ions: Charged atoms; cations (positive) form by losing electrons; anions (negative) form by gaining electrons.
Examples:
- Sodium (Na): Neutral atom with 11 protons/electrons; forms Na$^{+}$ by losing one electron.
- Fluorine (F): Neutral atom with 9 protons/electrons; forms F$^{-}$ by gaining one electron.

For main group metals, the charge correlates with their position on the periodic table:
- Group 1: +1 charge
- Group 2: +2 charge
- Group 3: +3 charge
For nonmetals:
- Group 15: -3 charge
- Group 16: -2 charge
- Group 17: -1 charge

Isotopes: Atoms of the same element with different numbers of neutrons, thus differing in mass.
Example of Carbon Isotopes:
- Carbon-12: 6 protons, 6 neutrons; atomic mass 12.
- Carbon-13: 6 protons, 7 neutrons; atomic mass 13.
- Carbon-14: 6 protons, 8 neutrons; atomic mass 14; used in radiocarbon dating.

Format:
- Chemical symbol (e.g., $^{14}_{6} ext{C}$ ) indicates mass number and atomic number.
Alternative naming notation: carbon-12, carbon-13, carbon-14.

Mass Number: Total number of protons and neutrons in a specific isotope.
Atomic Mass: Weighted average of all naturally occurring isotopes of an element based on relative abundance.

Extent considered; practicing calculation not required. Examples include conversions of % abundance into decimals for calculation of average atomic mass.

Recap on key laws (law of conservation of matter, law of definite proportions).
Summary of atomic structures, definitions, and the importance of the periodic table.