Historical Development of Atomic Theory

• Ancient Greek philosophers believed all 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; 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, in nucleus)

  • Neutrons (no charge, 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.

• The overall charge of an atom is determined by the difference between the number of protons and electrons: Charge=#protons-#electrons

Defining Elements and Isotopes

• Each atom has an atomic number (Z) equal to the number of protons, and a mass number (A) equal to the total number of protons and neutrons (nucleons):

  • A=#protons+#neutrons

  • Z=#protons

• An element is defined by its atomic number; 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 applications include using 14C in tooth enamel to estimate year of birth, based on atmospheric nuclear testing history.

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 average atomic mass of elements, which is a weighted average of all naturally occurring isotopes: Average atomic mass=∑iisotopes(fractional abundance of i×mass of i)

• Most elements have more than one naturally occurring isotope, so their atomic masses are not whole numbers.

Example Calculations

• Given isotopic abundances and masses (e.g., for silicon or gallium), the average atomic mass can be estimated and calculated using the formula above.