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_{reactants}=mass_{products}

• Joseph Proust demonstrated the law of definite proportions: compounds have constant composition by mass.

• John Dalton's atomic theory (1808):

  • All matter consists of solid, indivisible atoms.

  • Atoms are indestructible and retain identity in chemical reactions.

  • Atoms of the same element are identical; atoms of different elements differ in mass and properties.

  • 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 ≈ 1.672622×10−27kg

  • Neutrons: no charge, mass ≈ 1.674927×10−27kg

  • Electrons: negative charge, mass ≈ 9.109382×10−31kg

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

• Electrons move around the nucleus, balancing the atom's charge: Charge=Number of protons−Number of electrons

Defining Elements and Isotopes

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

• Mass number (A): total number of protons and neutrons (nucleons): A=p+n

• Isotopes: atoms of the same element (same Z) with different mass numbers (A), due to different numbers of neutrons.

Applications and Measurement of Isotopes

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

• Forensic example: 14C in tooth enamel can determine year of birth within 1.6 years, based on atmospheric nuclear testing history.

Measuring Isotopes: Mass Spectrometry

• Mass spectrometry separates isotopes by mass, producing a spectrum showing the proportion of each isotope in a sample.

• Atomic mass on the periodic table is a weighted average of all naturally occurring isotopes: Atomic mass=∑fractional abundance × isotope mass

• Most elements have more than one naturally occurring isotope; only a few (e.g., F, Na, P) have just one.

Calculating Average Atomic Mass

• Example: Silicon has three isotopes with known abundances and masses; average atomic mass is calculated using the weighted average formula.

• Example: Gallium has two isotopes; their natural abundances can be determined from the average atomic mass and isotope masses.