Historical Development of Atomic Theory

• Ancient Greek philosophers believed all matter was made of four elements (air, earth, fire, water).

• Democritus proposed matter is composed of indivisible particles called atoms.

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

• Proust demonstrated the law of definite proportions: compounds always contain the same elements in the same proportions by mass.

• Dalton's atomic theory (1808) stated:

  • All matter consists of solid, indivisible atoms.

  • Atoms retain their identity in chemical reactions.

  • Atoms of the same element are identical; different elements have different atoms.

  • Compounds are formed from elements in small whole-number ratios.

Structure of the Atom

• Atoms are made of three subatomic particles:

  • Protons (positive charge, in nucleus)

  • Neutrons (neutral, in nucleus)

  • Electrons (negative charge, outside nucleus)

• Most of the atom's mass is in the tiny, dense nucleus; most of its volume is empty space.

• Atomic charge is determined by: Charge=#protons-#electrons

Defining Elements and Isotopes

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

• Mass number (A) = number of protons + neutrons (nucleons).

• Isotopes: atoms of the same element with different numbers of neutrons (different mass numbers).

Applications and Importance of Isotopes

• Isotope ratios are used in fields like biology, geology, archaeology, and forensics.

• 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 showing the proportion of each isotope in a sample.

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

Average Atomic Mass

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

• Weighted average atomic mass is calculated as: Atomic mass=∑iisotopes(fractional abundance × isotope mass)

• Example: Silicon has three naturally occurring isotopes; their abundances and masses are used to calculate the average atomic mass.