Atomic Structure and Nuclear Chemistry

Historical Development of Atomic Theory

The concept of the atom has evolved over centuries, beginning with philosophical ideas and culminating in scientific theories based on experimental evidence.

• Ancient Greek Philosophers: Proposed that all matter was composed of four elements: air, earth, fire, and water. Democritus introduced the idea of the atom as an indivisible particle.

• Law of Conservation of Mass (Lavoisier, 1785): Mass is neither created nor destroyed in chemical reactions.

• Law of Constant Composition (Proust, 1794): A chemical compound always contains the same proportion of elements by mass.

• Dalton's Atomic Theory (1808):

  • All matter consists of solid, indivisible atoms.

  • Atoms are indestructible and retain their 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 formed by combinations of atoms in small whole-number ratios.

• Modern Modifications: Some postulates have been revised (e.g., atoms can be divided in nuclear reactions, isotopes exist).

Structure of the Atom

Atoms are composed of three fundamental subatomic particles: protons, neutrons, and electrons. Most of the atom's mass is concentrated in a tiny nucleus, while electrons occupy the surrounding space.

• Proton: Positively charged particle located in the nucleus.

• Neutron: Neutral particle also found in the nucleus.

• Electron: Negatively charged particle found in the outer regions of the atom.

Charge of an atom:

Experimental Evidence: Rutherford's Gold Foil Experiment

Rutherford's experiment demonstrated that atoms have a small, dense, positively charged nucleus, with most of the atom being empty space.

• Alpha particles were directed at a thin gold foil.

• Most particles passed through, but some were deflected, indicating a dense nucleus.

Defining Elements and Isotopes

Each element is defined by its atomic number (number of protons). Atoms of the same element can have different numbers of neutrons, resulting in isotopes.

• Atomic Number (Z): Number of protons in the nucleus.

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

• Isotopes: Atoms of the same element with different numbers of neutrons (and thus different mass numbers).

Example: Carbon has three naturally occurring isotopes: , , and .

Applications of Isotopes

Isotope ratios are used in various scientific fields for tracing and dating samples. For example, the amount of in tooth enamel can be used to estimate the year of birth.

Additional info: This method relies on the spike in atmospheric due to nuclear bomb testing in the mid-20th century.

Measuring Isotopes: Mass Spectrometry

Mass spectrometry is a technique used to determine the relative abundance of isotopes in a sample by separating ions based on their mass-to-charge ratio.

• Sample is ionized and accelerated through a magnetic field.

• Ions are deflected according to their mass-to-charge ratio and detected.

The resulting spectrum shows the proportion of each isotope present.

Average Atomic Mass

The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes of an element.

• Weighted Average Formula:

Example: Silicon has three isotopes: 92.23% (27.9769 u), 4.67% (28.9765 u), and 3.10% (29.9738 u). The average atomic mass is calculated using the formula above.

Example: Gallium has two isotopes, and , with an average atomic mass of 69.723 u. The more abundant isotope can be predicted by comparing the average atomic mass to the masses of the isotopes.

Summary Table: Key Terms and Concepts