Atoms and Elements

Introduction

This chapter explores the foundational concepts of atomic structure, the discovery of subatomic particles, isotopes, and the organization of elements in the periodic table. Understanding these principles is essential for further study in chemistry, as they explain the nature of matter and its transformations.

2.1 Confirming the Existence of Atoms

Modern Evidence for Atoms

• Atoms are the smallest units of elements that retain chemical identity.

• Advanced instruments, such as the scanning tunneling microscope (STM), allow scientists to visualize individual atoms on surfaces.

• Direct imaging of atoms provides experimental confirmation of their existence.

2.2 Early Ideas About the Building Blocks of Matter

Historical Development of Atomic Theory

• Ancient philosophers speculated about the divisibility of matter, but lacked experimental evidence.

• John Dalton (1803) proposed the first modern atomic theory, stating that elements are composed of tiny, indivisible particles called atoms.

2.3 Modern Atomic Theory and the Laws That Led To It

Key Laws and Their Scientists

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

• Law of Definite Proportions (Proust): All samples of a compound have the same proportions of their constituent elements.

• Law of Multiple Proportions (Dalton): When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

2.4 The Discovery of the Electron

Subatomic Particles and Atomic Models

• J.J. Thomson discovered the electron in 1897 using cathode ray experiments, showing that atoms contain negatively charged particles.

• Thomson's "plum pudding" model depicted electrons embedded in a positively charged sphere.

2.5 The Structure of the Atom

Development of Atomic Models

• Ernest Rutherford (1911) discovered the nucleus through gold foil experiments, showing that atoms have a dense, positively charged center.

• Most of the atom's volume is empty space occupied by electrons.

• Later models (Bohr, Schrödinger) introduced quantized orbits and electron orbitals.

2.6 Subatomic Particles: Protons, Neutrons, and Electrons in Atoms

Composition of the Atom

• Protons (p+): Positively charged particles in the nucleus.

• Neutrons (n0): Neutral particles in the nucleus.

• Electrons (e-): Negatively charged particles in orbitals around the nucleus.

• The number of protons defines the element (atomic number Z).

• The sum of protons and neutrons gives the mass number (A).

2.6 Isotopes: When the Number of Neutrons Varies

Definition and Representation

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

• Isotopes are represented as Element-Mass Number (e.g., C-12, C-13, C-14).

• Isotopes have nearly identical chemical properties but different masses.

2.7 Finding Patterns: The Periodic Law and the Periodic Table

Organization of Elements

• The Periodic Law states that when elements are arranged by increasing atomic number, their properties recur periodically.

• The Periodic Table organizes elements into groups (columns) and periods (rows) based on similar properties.

• Metals tend to lose electrons and form cations; nonmetals tend to gain electrons and form anions.

2.7 Ions and the Periodic Table

Formation of Ions

• Ions are charged atoms or molecules formed by the loss or gain of electrons.

• Metals (left side of the table) typically form positive ions (cations).

• Nonmetals (right side of the table) typically form negative ions (anions).

2.8 Atomic Mass: The Average Mass of an Element’s Atoms

Weighted Average and Isotopic Abundance

• The atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

• Atomic mass is expressed in atomic mass units (amu).

• Calculation: $\text{Atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

2.9 Molar Mass: Counting Atoms by Weighing Them

The Mole and Avogadro’s Number

• One mole contains $6.022 \times 10^{23}$ entities (Avogadro’s number).

• The molar mass (g/mol) is numerically equal to the atomic or molecular mass in amu.

• Conversions between mass, moles, and number of atoms use the following relationships:

• $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

• $\text{Number of atoms} = \text{Moles} \times 6.022 \times 10^{23}$

Applications: Radiology and Isotopes

Medical Use of Radioisotopes

• Iodine-131 is a radioactive isotope used in thyroid imaging and treatment.

• The thyroid gland absorbs iodine; radioactive iodine emits particles that can destroy overactive tissue or be detected for imaging.

• Understanding isotopes and atomic counting is essential for radiotracer applications in medicine.

Summary Table: Key Subatomic Particles

Additional info:

• Atomic models evolved from Dalton's solid sphere to Thomson's plum pudding, Rutherford's nuclear model, Bohr's quantized orbits, and Schrödinger's electron cloud model.

• The periodic table is a powerful tool for predicting chemical behavior, ion formation, and element classification.