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 definite proportions: compounds always contain the same elements in the same proportions 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; different elements have different atoms with different masses.

  • Compounds are formed from 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 mass is in the tiny, dense nucleus (contains protons and neutrons); most of the atom's volume is empty space.

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

• Atomic charge is calculated as: Charge=#protons-#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.

Applications and Importance of Isotopes

• Isotope ratios are used in fields like biology, geology, paleontology, and archaeology for tracing and dating samples.

• Forensic example: 14C in tooth enamel can determine year of birth due to changes in atmospheric 14C from nuclear testing.

Measuring Isotopes: Mass Spectrometry

• Mass spectrometry separates isotopes based on mass, producing a spectrum that shows the proportion of each isotope in a sample.

• This technique allows determination of the isotopic composition and calculation of average atomic mass.

Average Atomic Mass

• Most elements have more than one naturally occurring isotope; atomic mass on the periodic table is a weighted average of all isotopes.

• Weighted average atomic mass is calculated as: 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 relative abundances.