Atoms & Elements

What are Atoms?

An atom is the smallest identifiable unit of an element, serving as the fundamental building block of matter. There are approximately 91 naturally occurring elements, with scientists having synthesized about 20 additional elements. Each element is composed of atoms unique to that element.

History of Atomic Theory

The concept of the atom originated with the Greek philosopher Democritus, who proposed that matter was composed of small, indivisible particles called "atomos." This idea laid the foundation for later scientific developments.

Dalton's Model of the Atom

John Dalton formalized atomic theory in the early 19th century, stating:

• All matter is made up of tiny particles called atoms or molecules.

• Molecules can be broken down into atoms by chemical processes, but atoms cannot be broken down further.

• All atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.

• Atoms combine in simple, whole-number ratios to form compounds.

Dalton's theory relied on the law of conservation of mass and the law of definite composition. For example, carbon dioxide (CO2) always has a fixed ratio of carbon to oxygen, and water (H2O) always has two hydrogen atoms and one oxygen atom.

Thomson's Model of the Atom

J.J. Thomson discovered the electron, a negatively charged particle much smaller and lighter than atoms. He proposed the "plum pudding model," where electrons are embedded in a sphere of positive charge, maintaining overall neutrality.

Thomson's experiments showed that electrons are present in many substances, and he calculated the mass-to-charge ratio of electrons and protons.

Rutherford's Model of the Atom

Ernest Rutherford conducted the gold foil experiment, firing alpha particles at a thin sheet of gold. Most particles passed through, but some were deflected, indicating the presence of a dense, positively charged nucleus surrounded by empty space.

The experiment demonstrated that atoms are mostly empty space, with a tiny, dense nucleus at the center.

Discovery of Subatomic Particles

Rutherford predicted the existence of neutral particles (neutrons) in the nucleus, which were discovered about 30 years later. Neutrons are similar in mass to protons but have no charge.

Atomic Notation

Atomic number (Z) is the number of protons in the nucleus and identifies the element. Mass number (A) is the sum of protons and neutrons. Atomic notation is a shorthand for expressing the composition of an atomic nucleus:

• Element symbol: e.g., Na for sodium

• Atomic number (Z): number of protons

• Mass number (A): protons + neutrons

For sodium (Na): Z = 11, A = 23, so neutrons = 23 - 11 = 12.

Isotopes

Atoms of the same element with different numbers of neutrons are called isotopes. Isotopes have the same atomic number but different mass numbers. Not all isotopes are stable.

• Hydrogen isotopes: Protium (1 proton, 0 neutrons), Deuterium (1 proton, 1 neutron), Tritium (1 proton, 2 neutrons, radioactive)

• Carbon isotopes: C-12 (6 protons, 6 neutrons), C-13 (6 protons, 7 neutrons), C-14 (6 protons, 8 neutrons, radioactive)

Atomic Mass

The atomic mass is the weighted average mass of all naturally occurring isotopes, measured in atomic mass units (amu). The formula is:

Example: Chlorine consists of 75.77% Cl-35 (34.97 amu) and 24.23% Cl-37 (36.97 amu):

The Periodic Table

Mendeleev's Periodic Table

Dmitri Mendeleev arranged elements in order of increasing atomic mass, grouping elements with similar properties. He left gaps for undiscovered elements and predicted their properties.

Modern Periodic Table

The modern periodic table arranges elements by increasing atomic number (Z), following the periodic law: properties of elements repeat in a predictable pattern. The table is organized into groups (vertical columns) and periods (horizontal rows).

Classification of Elements

Elements are classified as metals, nonmetals, and metalloids based on their properties.

Metals

• Solids at room temperature (except Hg)

• Shiny, reflective surface

• Conduct heat and electricity

• Malleable and ductile

• Lose electrons to form cations

• Located on the lower left of the periodic table

Nonmetals

• Found in all three states (solid, liquid, gas)

• Poor conductors of heat and electricity

• Solids are brittle

• Gain electrons to form anions

• Located on the upper right of the periodic table (except H)

Metalloids (Semi-metals)

• Show properties of both metals and nonmetals

• Often semiconductors

• Example: Silicon is shiny, conducts electricity, but is brittle and does not conduct heat well

Groups and Periods

Groups are vertical columns (18 total), and elements in the same group exhibit similar properties. Periods are horizontal rows (7 total).

Group A Elements (Representative/Main Group)

• IA: Alkali metals (very reactive)

• IIA: Alkaline earth metals (fairly reactive)

• VA: Pnictogens

• VIA: Chalcogens

• VIIA: Halogens (very reactive)

• VIIIA: Noble gases (inert)

Group B Elements (Transition and Inner Transition)

• Transition elements: Middle of the periodic table, unpredictable properties

• Inner transition elements: Lanthanide (rare earth) and actinide (radioactive) series

• Transuranium elements: Z > 93, man-made

Element Names and Symbols

Each element has a unique name and symbol, often derived from Latin, Greek, or geographic locations. Symbols are one or two letters, with the first letter capitalized and the second lowercase.

Ions and Ionic Charge

Formation of Ions

Metals tend to lose electrons to achieve a stable noble gas configuration, forming positively charged cations. Nonmetals gain electrons to form negatively charged anions. The charge of an ion is related to the number of valence electrons.

• Group IA metals lose one electron to form +1 ions:

• Group VIA nonmetals gain two electrons to form -2 ions:

Examples of Ionic Formulas

• Chlorine: Cl-

• Calcium: Ca2+

• Phosphorous: P3-

• Sodium: Na+

• Oxygen: O2-

• Aluminum: Al3+

Additional info:

This study guide covers the foundational concepts of atoms, elements, isotopes, atomic mass, the periodic table, classification of elements, and the formation of ions, all of which are essential for an Introduction to Chemistry course.