Atomic Structure and Nuclear Chemistry: Foundations and Applications

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Atomic Structure and Nuclear Chemistry

Who Thought of the Atom?

The concept of the atom has evolved over centuries, beginning with philosophical ideas and advancing through scientific experimentation. Early theories laid the groundwork for modern atomic theory.

Ancient Greek Philosophers: Proposed that all matter consists of combinations of four elements: air, earth, fire, and water.
Democritus (~460–370 B.C.): Suggested that matter could be subdivided until reaching an indivisible particle, called the atom (from Greek a-tomos, meaning "uncuttable").
Antoine Lavoisier (1743–1794): Discovered the law of conservation of mass, stating that mass is conserved 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.

Example: Water (H2O) always contains hydrogen and oxygen in a mass ratio of approximately 1:8.

John Dalton's Atomic Theory

In 1808, John Dalton formalized atomic theory based on previous ideas and experimental evidence. His postulates formed the basis for understanding chemical reactions and the nature of matter.

All matter consists of solid and indivisible atoms.
Atoms are indestructible and retain their identity in chemical reactions.
All atoms of a given chemical element are identical in mass and properties.
Atoms of different elements differ in mass and other properties.
Compounds consist of elements combined in small whole-number ratios.

Additional info: Modern atomic theory has modified Dalton's postulates, recognizing that atoms can be divided (into subatomic particles) and that isotopes exist.

What's in an Atom?

Subatomic Particles

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

Particle	Mass (kg)	Charge (C)	Relative Charge	Location
Electron	9.109382 × 10−31	−1.602176 × 10−19	−1	Outside nucleus
Proton	1.672622 × 10−27	+1.602176 × 10−19	+1	Nucleus
Neutron	1.674927 × 10−27	0	0	Nucleus

Most of an atom is empty space. The nucleus is extremely small compared to the overall size of the atom, but contains most of its mass.
Nucleus: Contains protons and neutrons, bound together in a region of positive charge (as shown by Rutherford's gold foil experiment).
Electrons: Move around the nucleus, balancing the overall charge of the atom.

Formula:

Example: A neutral carbon atom has 6 protons and 6 electrons, so its charge is zero.

Defining an Element

Atomic Number and Mass Number

Each atom is characterized by two numbers:

Atomic number (Z): Number of protons in the nucleus. Defines the element.
Mass number (A): Total number of nucleons (protons + neutrons).

Formula:

Changing the number of protons (as in nuclear reactions) changes the element.
Atoms of the same element can have different mass numbers; these are called isotopes.

Example: Carbon has three naturally occurring isotopes:

Isotope	Protons (Z)	Neutrons	Mass Number (A)
12C	6	6	12
13C	6	7	13
14C	6	8	14

Additional info: The term nucleon refers to either a proton or a neutron.

Isotopes and Their Applications

Uses of Isotopes

Isotopes have important applications in science and technology, including tracing, dating, and forensic analysis.

Isotope ratios are used to trace and date samples in biology, geology, paleontology, and archaeology.
Forensic application: The amount of 14C in tooth enamel can determine the year of birth, due to changes in atmospheric 14C from nuclear testing.

Example: Radiocarbon dating uses the decay of 14C to estimate the age of organic materials.

Measuring Isotopes: Mass Spectrometry

Mass spectrometry is a technique used to determine the isotopic composition of a sample by measuring the mass and abundance of its atoms.

Produces a spectrum showing the proportion of atoms belonging to each isotope.
Allows calculation of average atomic mass and identification of isotopic ratios.

Example: Chlorine has two main isotopes, and mass spectrometry reveals their relative abundances.

Average Atomic Mass

Weighted Average Calculation

The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes of an element, factoring in both mass and abundance.

Formula:

Example: Silicon has three naturally occurring isotopes:

92.23% 28Si (27.9769 u)
4.67% 29Si (28.9765 u)
3.10% 30Si (29.9738 u)

Average atomic mass calculation:

Additional info: Most elements have more than one naturally occurring isotope; only a few (e.g., F, Na, P) are monoisotopic.

Gallium Isotope Example

Gallium has two naturally occurring isotopes:

69Ga (mass = 68.926 u)
71Ga (mass = 70.925 u)

Given the average atomic mass (69.723 u), the more abundant isotope can be predicted and the natural abundance of each calculated using the weighted average formula.