Atomic Structure and Nuclear Chemistry

Introduction

This study guide covers the foundational concepts of atomic structure and nuclear chemistry, including the historical development of atomic theory, the structure and properties of atoms, the definition of elements and isotopes, and the measurement and application of isotopic data. These topics are essential for understanding the nature of matter and the principles underlying chemical reactions.

Who Thought of the Atom?

Historical Development of Atomic Theory

• Ancient Greek Philosophers: Early thinkers proposed that all matter was composed of four elements: air, earth, fire, and water.

• Democritus (~460–370 B.C.): Suggested that matter could be divided repeatedly until reaching an indivisible particle, which he called the atom (from Greek a-tomos, meaning "uncuttable").

• Antoine Lavoisier (1743–1794): Formulated the law of conservation of mass, stating 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 elements in the same proportions by mass.

Dalton's Atomic Theory (1808)

• 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 properties.

• Atoms of different elements differ in mass and other properties.

• Compounds are formed by the combination of atoms of different elements in small whole-number ratios.

Note: While Dalton's theory laid the groundwork for modern chemistry, some postulates have been modified (e.g., atoms are divisible into subatomic particles, and isotopes exist).

What’s in an Atom?

Subatomic Particles

Atoms are composed of three main types of subatomic particles:

• Most of an atom is empty space. The nucleus is extremely small compared to the overall size of the atom, but contains nearly all of its mass.

• Nucleus: Contains protons and neutrons, is positively charged, and is the dense center of the atom.

• Electrons: Move around the nucleus in regions of space, balancing the overall charge of the atom.

Atomic charge formula:

Defining an Element

Atomic Number and Mass Number

• Atomic Number (Z): The number of protons in the nucleus of an atom. This defines the element.

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

For example, the notation for carbon-12 is:

• Changing the number of protons (atomic number) changes the element (as in nuclear reactions).

• Atoms of the same element can have different mass numbers; these are called isotopes.

Isotopes

• Isotopes: Atoms of the same element (same number of protons) but different numbers of neutrons (different mass numbers).

• Most elements have more than one naturally occurring isotope.

Examples:

• Hydrogen: (protium), (deuterium), (tritium)

• Carbon: , ,

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

How Useful are Isotopes?

Applications of Isotopes

• Isotope ratios are used in fields such as biology, geology, paleontology, and archaeology to trace and date samples.

• Forensic applications: For example, the amount of in tooth enamel can be used to determine the year of birth, based on atmospheric levels that changed during nuclear bomb testing (1955–1963).

Example: Measuring in tooth enamel can estimate the year the enamel was formed, and thus the individual's year of birth, to within about 1.6 years.

How Can We Measure Isotopes?

Mass Spectrometry

• Mass spectrometry (MS): A technique that separates isotopes of an element based on their mass-to-charge ratio, producing a spectrum that shows the proportion of each isotope in a sample.

• Main difference between isotopes: Their mass (number of neutrons).

• Application: Determining the isotopic composition and average atomic mass of elements.

“Average” Samples and Atomic Mass

Weighted Average Atomic Mass

• Most elements exist as mixtures of isotopes.

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

• The average atomic mass is calculated using the mass and relative abundance of each isotope:

Where:

• Isotopic mass: The mass of a specific isotope (in atomic mass units, u).

• Fractional abundance: The fraction of atoms of that isotope in a natural sample.

Examples

• Silicon: Has three naturally occurring isotopes:

  • 92.23% (27.9769 u)

  • 4.67% (28.9765 u)

  • 3.10% (29.9738 u)

  Calculation: Estimate and calculate the average atomic mass using the formula above.

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

  • : 68.926 u

  • : 70.925 u

  Application: Predict which isotope is more abundant and calculate the natural abundance of each isotope.

Summary Table: Subatomic Particles

Additional info: Some context and explanations have been expanded for clarity and completeness, including the explicit formula for average atomic mass and the summary table of subatomic particles.