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 elements in the same proportions 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 the combination of atoms in small whole-number ratios.

  Note: Modern science has modified some of Dalton's postulates, recognizing the existence of subatomic particles and isotopes.

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, dense nucleus.

• Protons: Positively charged particles located in the nucleus.

• Neutrons: Neutral particles also found in the nucleus.

• Electrons: Negatively charged particles that occupy the space around the nucleus.

• Most of the atom is empty space. The nucleus is extremely small compared to the overall size of the atom.

Charge of an atom:

Experimental Evidence: The Gold Foil Experiment

Ernest Rutherford's gold foil experiment provided crucial evidence for the nuclear model of the atom, demonstrating that atoms have a small, dense, positively charged nucleus.

• Most alpha particles passed through gold foil, indicating that most of the atom is empty space.

• Some particles were deflected, suggesting a dense, positively charged nucleus.

Defining an Element and Isotopes

Each element is defined by its atomic number (Z), which is the number of protons in its nucleus. The mass number (A) is the total number of protons and neutrons (nucleons).

• Atomic Number (Z): Number of protons; determines the element.

• Mass Number (A): Number of protons + neutrons.

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A).

• Changing the number of protons changes the element; changing the number of neutrons creates isotopes.

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

Applications of Isotopes

Isotope ratios are used in various scientific fields for tracing, dating, and forensic analysis.

• Radiocarbon Dating: levels in biological samples can determine the age of archaeological finds.

• Forensic Science: The amount of in tooth enamel can estimate the year of birth, especially after nuclear bomb testing altered atmospheric $^{14}\text{C}$ levels.

Measuring Isotopes: Mass Spectrometry

Mass spectrometry is a technique used to determine the isotopic composition of elements in a sample by separating ions based on their mass-to-charge ratio.

• Atoms are ionized and accelerated through a magnetic field.

• Ions of different masses are deflected by different amounts and detected separately.

• The resulting spectrum shows the relative abundance of each isotope.

Average Atomic Mass

The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes of an element, taking into account both their masses and relative abundances.

• Formula:

• Example (Silicon):

  • 92.23% Si (27.9769 u)

  • 4.67% Si (28.9765 u)

  • 3.10% Si (29.9738 u)

  • Calculate the average atomic mass using the formula above.

• Example (Gallium): Given the average atomic mass and the masses of Ga and Ga, predict the more abundant isotope and calculate their natural abundances.

Summary Table: Subatomic Particles

Key Takeaways:

• Atoms are composed of protons, neutrons, and electrons.

• Isotopes are atoms of the same element with different numbers of neutrons.

• Mass spectrometry is used to determine isotopic composition.

• The atomic mass on the periodic table is a weighted average of all isotopes.