Atomic Structure and the Development of Atomic Theory

Historical Perspectives on the Atom

The concept of the atom has evolved over centuries, beginning with philosophical ideas and advancing through scientific experimentation and theory.

• Ancient Greek Philosophers: Proposed that all matter was composed of four elements: air, earth, fire, and water. Democritus (~460–370 B.C.) introduced the idea of the atom (from Greek: a = not; tomos = cut), an indivisible particle that makes up all matter.

• Antoine Lavoisier (1743–1794): Discovered the law of conservation of mass, which states that mass is neither created nor destroyed in chemical reactions.

• Joseph Proust (1754–1826): Demonstrated the law of constant composition (also known as the law of definite proportions), which states that a chemical compound always contains the same proportion of elements by mass.

Dalton's Atomic Theory

In 1808, John Dalton synthesized earlier ideas into the first modern atomic theory, which laid the groundwork for our understanding of matter.

• All matter consists of solid and indivisible atoms.

• Atoms are indestructible and retain their identity in all chemical reactions.

• All atoms of a given chemical element are identical in mass and in all other properties.

• Different elements have different kinds of atoms; these atoms differ in mass from element to element.

• Compounds consist of elements combined in small whole-number ratios.

Note: While Dalton's theory was foundational, later discoveries (such as subatomic particles and isotopes) have led to modifications of some postulates.

Structure of the Atom

Subatomic Particles

Atoms are composed of three fundamental subatomic particles, each with distinct properties:

• Most of an atom is empty space; the nucleus is extremely small compared to the overall size of the atom (about 1/1,000,000,000,000,000th of the atom's volume).

• The nucleus contains most of the atom's mass and is composed of protons and neutrons bound together in a region of positive charge.

• Electrons travel around the nucleus, balancing the overall charge of the atom.

Charge of an atom:

Defining Elements and Isotopes

Atomic Number and Mass Number

Each element is uniquely defined by its atomic number and mass number:

• Atomic number (Z): Number of protons in the nucleus. Determines the identity of the element.

• Mass number (A): Total number of nucleons (protons + neutrons) in the nucleus.

Changing the number of protons (as in a nuclear reaction) changes the element.

Isotopes

Atoms of the same element can have different numbers of neutrons, resulting in different mass numbers. These variants are called isotopes.

• Isotopes have the same atomic number but different mass numbers.

• Most elements have more than one naturally occurring isotope.

Nucleon is a general term for a proton or neutron (i.e., a particle in the nucleus).

Applications and Measurement of Isotopes

Uses of Isotopes

Isotope ratios are valuable in various scientific fields for tracing and dating samples:

• Applications in biology, geology, paleontology, archaeology, and forensic science.

• Forensic example: Atmospheric 14C levels increased during nuclear bomb testing (1955–1963) and decreased after testing was banned. The amount of 14C in tooth enamel can be used to determine the year of formation and, thus, the year of birth (within 1.6 years).

Measuring Isotopes: Mass Spectrometry

To determine which isotopes are present in a sample, chemists use mass spectrometry:

• Isotopes of an element differ in mass.

• Mass spectrometry separates isotopes based on their mass-to-charge ratio and provides a spectrum showing the proportion of each isotope in the sample.

Example: The average atomic mass of chlorine is 35.4527 u, reflecting the weighted contributions of its isotopes.

Atomic Mass and Isotopic Abundance

Weighted Average Atomic Mass

Most elements exist as mixtures of isotopes. The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes, factoring in both mass and abundance:

Where:

Examples of Isotopic Calculations

• Silicon: Has three naturally occurring isotopes:

  • 92.23% 28Si (27.9769 u)

  • 4.67% 29Si (28.9765 u)

  • 3.10% 30Si (29.9738 u)

  Task: Estimate and then calculate the average atomic mass of silicon using the weighted average formula.

• Gallium: Has two naturally occurring isotopes and an average atomic mass of 69.723 u.

  • 69Ga: atomic mass = 68.926 u

  • 71Ga: atomic mass = 70.925 u

  Task: Predict which isotope is more abundant, then calculate the natural abundance of each isotope using the average atomic mass.

Additional info: These calculations are fundamental for understanding the composition of elements and are frequently tested in introductory chemistry courses.