Historical Development of Atomic Theory

• Ancient Greek philosophers believed matter was made of four elements (air, earth, fire, water), but Democritus proposed matter is composed of indivisible particles called atoms.

• Antoine Lavoisier established the law of conservation of mass: mass is neither created nor destroyed in chemical reactions.

• Joseph Proust demonstrated the law of constant composition (law of definite proportions): compounds always contain the same proportion of elements by mass.

• John Dalton's atomic theory (1808) stated:

  • All matter consists of solid, indivisible atoms.

  • Atoms are indestructible and retain identity in chemical reactions.

  • Atoms of a given element are identical in mass and properties.

  • Atoms of different elements differ in mass and properties.

  • Compounds are combinations of elements in small whole-number ratios.

Structure of the Atom

• Atoms are made of three subatomic particles:

  • Protons (positive charge, mass ≈ 1.67 × 10-27 kg)

  • Neutrons (no charge, mass ≈ 1.67 × 10-27 kg)

  • Electrons (negative charge, mass ≈ 9.11 × 10-31 kg)

• Most of the atom's volume is empty space; the nucleus (tiny, dense, positively charged) contains protons and neutrons.

• Electrons move around the nucleus, balancing the atom's overall charge.

• Atomic charge is calculated as: Charge=Number of protons−Number of electrons

Defining Elements and Isotopes

• Each atom has:

  • Atomic number (Z): number of protons (defines the element).

  • Mass number (A): total number of protons and neutrons (nucleons).

• Changing the number of protons changes the element.

• Atoms of the same element with different numbers of neutrons are called isotopes.

• Most elements have more than one naturally occurring isotope.

Applications and Measurement of Isotopes

• Isotope ratios are used in fields like biology, geology, paleontology, archaeology, and forensics (e.g., dating tooth enamel using 14C levels).

• Mass spectrometry is used to measure isotopic composition by separating isotopes based on mass and displaying their relative abundances as a spectrum.

Atomic Mass and Isotopic Abundance

• Most elements exist as mixtures of isotopes; atomic mass on the periodic table is a weighted average of all naturally occurring isotopes.

• The average atomic mass is calculated as: Average atomic mass=∑iall isotopes(fractional abundance×isotope mass)

• Example: Silicon has three isotopes with different abundances and masses; the average atomic mass is calculated using their respective values.