Foundations of Atomic Theory and the Periodic Table: Key Laws, Experiments, and Concepts

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Atomic Theory and Fundamental Laws

Law of Conservation of Mass

The Law of Conservation of Mass states that mass cannot be created or destroyed in a chemical or physical process. This foundational principle underlies all chemical reactions.

Antoine Lavoisier established this law through careful experimentation.
All mass must remain constant and accounted for during any reaction.
Example: When burning wood, the total mass of the products (ash, gases) equals the mass of the original wood and oxygen.

Law of Definite Proportions

The Law of Definite Proportions states that a given compound always contains the same proportion by mass of its constituent elements, regardless of its source or method of preparation.

Joseph Proust formulated this law.
The mass fraction of an element in a compound is constant.
Example: Water (H2O) always contains 11.2% hydrogen and 88.8% oxygen by mass.

Law of Multiple Proportions

The Law of Multiple Proportions states that when two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

John Dalton introduced this law.
Atoms combine in simple whole-number ratios to form compounds.
Example: Carbon and oxygen form CO (carbon monoxide) and CO2 (carbon dioxide); the ratio of oxygen masses that combine with a fixed mass of carbon is 1:2.

Dalton's Atomic Theory

Dalton's Atomic Theory provided the first modern scientific description of the atom and its role in chemical reactions.

Each element is made up of tiny particles called atoms.
Atoms of a given element are identical in mass and properties.
Compounds are formed when atoms of different elements combine in fixed, whole-number ratios.
Chemical reactions involve the reorganization of atoms; atoms themselves are not changed in chemical reactions.

Important Atom Characterization Experiments

J.J. Thomson's Cathode Ray Tube Experiment

This experiment demonstrated the existence of negatively charged subatomic particles, later called electrons, common to all elements.

Measured the mass-to-charge ratio of the electron:
Proposed the "plum pudding model" of the atom, where electrons are embedded in a positively charged sphere.

Robert Millikan's Oil Drop Experiment

Millikan determined the absolute charge of an electron by observing the behavior of oil droplets in an electric field.

Measured the elementary charge:

Ernest Rutherford's Gold Foil Experiment

Rutherford's experiment revealed that the atom's mass is concentrated in a small, dense nucleus, and that atoms are mostly empty space.

Disproved the plum pudding model.
Led to the nuclear model of the atom.

Modern Atomic Theory

Structure and Properties of Atoms

Atoms are roughly spherical and electrically neutral.
Composed of three principal subatomic particles:
- Protons (p+): positively charged, located in the nucleus
- Neutrons (n0): neutral, located in the nucleus
- Electrons (e-): negatively charged, found outside the nucleus
Atomic notation:
- = atomic number = number of protons in the nucleus
- = mass number = number of protons + neutrons
- For a neutral atom:

Isotopes

Atoms of the same element with different mass numbers () due to varying numbers of neutrons.
Isotopes have the same chemical and physical properties.
Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Ions

Cations: Positively charged ions, formed by losing electrons (fewer electrons than protons).
Anions: Negatively charged ions, formed by gaining electrons (more electrons than protons).

Modern Atomic Theory (Updated)

Atoms are the smallest particle of matter with a unique identity.
Atoms of the same element have the same atomic number but can have different masses (isotopes).
Atoms of one element cannot be converted to atoms of another element by chemical processes (only nuclear reactions can change elements).
Compounds are formed when atoms of different elements combine in whole-number ratios.

The Periodic Table and Periodic Law

Periodic Law

The Periodic Law states that when elements are arranged in order of increasing atomic mass (now atomic number), certain sets of properties recur periodically.

Classification of Elements

Metals: Good conductors, malleable, ductile, shiny.
Non-metals: Poor conductors, brittle, dull.
Metalloids: Properties intermediate between metals and non-metals.

Groups of Elements and Common Names

Group Name	Group Number
Alkali Metals	Group 1 (IA)
Alkaline Earth Metals	Group 2 (IIA)
Halogens	Group 17 (VIIA)
Noble Gases	Group 18 (VIIIA)
Transition Metals	Groups 3-12 (B groups)
Main Group Elements	Groups 1, 2, 13-18
Lanthanides & Actinides	Separate rows at bottom

Ions and the Periodic Table

Main group metals tend to lose enough electrons to achieve the same number of electrons as the nearest noble gas (form cations).
Main group nonmetals tend to gain enough electrons to achieve the same number of electrons as the nearest noble gas (form anions).

Average Atomic Mass

The average atomic mass of an element is the weighted average of the masses of all naturally occurring isotopes, based on their relative abundances.

No atom of an element has exactly the average atomic mass, but for calculations, we use this value as a representative mass.
Formula:
Example: Chlorine has two main isotopes, Cl-35 and Cl-37; its average atomic mass is about 35.45 amu.

Stoichiometry

Stoichiometry is the study of the quantitative relationships between the amounts of reactants and products in chemical reactions.

The mole is the SI unit for amount of substance, defined as the number of atoms in exactly 12 grams of carbon-12.
Avogadro's number: entities per mole.
Conversions:
- 1 atom = average mass (amu)
- 1 mole of atoms = average mass (g)
Example: 1 mole of carbon-12 atoms has a mass of exactly 12 g and contains atoms.