Atoms, Atomic Theory, and the Periodic Table: Foundations of Chemistry

주석 실무 비디오 수업

스터디 가이드 - 스마트 노트

자료에 맞춘 맞춤형 노트, 핵심 정의, 예시, 맥락을 확장해 제공합니다.

Atoms: Ideas from the Ancient Greeks

Early Concepts of Matter

The concept of the atom originated in ancient Greece, where philosophers debated the nature of matter.

Four Main Components: Air, water, earth, and fire were believed to be the fundamental elements of nature.
Leucippus and Democritus: Proposed that matter is composed of tiny, indivisible particles called atomos.
Aristotle: Argued that matter was continuous and could be divided infinitely.
Scientific Limitation: There was no experimental method to prove either theory at the time.

Diagram of four elements: air, water, earth, fire

Scientific Laws: Conservation of Mass and Definite Proportions

Law of Conservation of Mass

The law of conservation of mass states that in a chemical reaction, matter is neither created nor destroyed.

Antoine Lavoisier: Formulated this law, which is fundamental to chemical calculations.
Implication: New materials are created by rearranging atoms, not by creating or destroying them.

Diagram showing conservation of mass in a chemical reaction Example of conservation of mass with mercuric oxide

Law of Definite Proportions

Joseph Proust observed that elements in a compound always occur in fixed proportions.

Definition: All samples of a given compound have the same proportions of their constituent elements, regardless of source or preparation.
Example: Silver sulfide always forms with a 2:1 ratio of silver to sulfur.

Lead and sulfur combining in fixed proportions

John Dalton and the Atomic Theory of Matter

Law of Multiple Proportions

Dalton's law states that when two elements form more than one compound, the masses of one element that combine with a fixed mass of the other can be expressed as ratios of small whole numbers.

Example: Carbon and oxygen form both CO and CO2, illustrating the law.

Carbon monoxide and carbon dioxide molecules Mass ratios for carbon and oxygen in compounds

Dalton's Atomic Theory

Dalton's atomic theory provided a foundation for modern chemistry.

All matter is composed of extremely small and indestructible particles called atoms.
Each element consists of just one type of atom, unique to that element.
Compounds are formed when atoms of different elements combine in fixed proportions.
Chemical reactions involve rearrangement of atoms; atoms are not created or destroyed.

Dalton's Atomic Theory table

Practice Questions

The term atom comes from the Greek word atomos, meaning "cannot be cut".
CO and CO2 formation illustrates the Law of Multiple Proportions.
2 g H2 + 16 g O2 → 18 g H2O illustrates the Law of Conservation of Mass.
Silver sulfide always forms with a 2:1 ratio of Ag to S, illustrating the Law of Definite Proportions.

Atoms, Isotopes, and Atomic Mass

Atomic Number and Mass Number

Atomic Number (Z): Number of protons in an atom's nucleus; defines the element.
Mass Number (A): Sum of protons and neutrons in the nucleus.

Atomic number, mass number, and chemical symbol notation

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

Example: Natural neon consists of three isotopes with different neutron counts and abundances.
Mass Number: Used to distinguish isotopes (e.g., chlorine-35, chlorine-37).

Atomic Mass

Atomic Mass: The average mass of an element's atoms, weighted by natural abundance.
Atomic Mass Unit (amu): Defined as 1/12 the mass of a carbon-12 atom.
Example: Chlorine has an atomic mass of 35.45 amu, reflecting its isotopic composition.

Periodic table element box for chlorine

The Mole and Molar Mass

Definition and Avogadro's Number

Chemists use the mole to count atoms, molecules, or particles.

Mole (mol): Amount of substance containing particles (Avogadro's number).
Analogy: Just as a dozen means 12, a mole means .

Eggs analogy for mole and Avogadro's number

Molar Mass

Molar Mass: Mass of one mole of atoms of an element, expressed in grams per mole (g/mol).
Example: 1 mol of copper atoms = 63.55 g; 1 mol of hydrogen atoms = 1.0079 g.
Molecular Weight: Mass of a single molecule, calculated by summing atomic masses of constituent atoms.

Relationship between grams, moles, and number of atoms

Using Moles in Calculations

The mole provides a bridge between mass and number of atoms.

Conversion: Use molar mass to convert between grams and moles, and Avogadro's number to convert between moles and number of particles.
Example: For methane (CH4), 12.0107 g C and 4.03176 g H (4 moles) combine in a 3:1 mass ratio.

Conversion factor example for methane

Mendeleev and the Periodic Table

Development of the Periodic Table

Dmitri Mendeleev organized elements by increasing atomic mass, revealing periodic patterns.

Gaps: Predicted undiscovered elements, such as eka-silicon (later found to be germanium).
Periodic Law: Elements with similar properties recur in a regular pattern.

Mendeleev's original periodic table Modern periodic table showing time of discovery

Prediction and Discovery

Mendeleev's predictions for eka-silicon closely matched the observed properties of germanium.

Property	Predicted for Eka-Silicon (1871)	Observed for Germanium (1886)
Atomic mass	72	72.6
Density (g/cm3)	5.5	5.47
Color	Dirty gray	Grayish white
Density of oxide (g/cm3)	EsO2: 4.7	GeO2: 4.703
Boiling point of chloride	EsCl4: below 100 °C	GeCl4: 86 °C

Comparison table of predicted and observed properties for germanium

Summary

The concept of the atom has evolved from philosophical speculation to scientific theory.
Key laws—conservation of mass, definite proportions, and multiple proportions—form the basis of chemical understanding.
Dalton's atomic theory, isotopes, atomic mass, and the mole are fundamental to quantifying chemical reactions.
The periodic table organizes elements by atomic number and reveals recurring chemical properties.