Atoms and Elements

Definition and Historical Context

Atoms are the fundamental building blocks of matter, originating from the Greek word atomos, meaning "indivisible." Early philosophers such as Leucippus and Democritus proposed that matter was composed of small, indestructible particles, while Plato and Aristotle believed in continuous matter composed of fire, air, earth, and water. The scientific approach, championed by John Dalton, provided evidence for the atomic theory, establishing atoms as the smallest identifiable units of elements.

Laws Leading to Atomic Theory

Three key laws underpin the development of atomic theory:

• Law of Conservation of Mass: In a chemical reaction, matter is neither created nor destroyed. The total mass of reactants equals the total mass of products.

• Law of Definite Proportions: All samples of a given compound have the same proportions of their constituent elements, regardless of source or preparation.

• Law of Multiple Proportions: When two elements form different compounds, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

Dalton's Atomic Theory

Dalton's atomic theory explains the above laws:

• Each element is composed of tiny, indestructible particles called atoms.

• All atoms of a given element have the same mass and properties.

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

• Atoms of one element cannot change into atoms of another element; they only rearrange in chemical reactions.

Discovery of Subatomic Particles

The Electron

J. J. Thomson discovered the electron through cathode ray experiments, revealing that these particles travel in straight lines, are independent of the cathode material, and carry a negative charge. He measured the charge-to-mass ratio of the electron as C/g.

Millikan's Oil Drop Experiment

Robert Millikan determined the charge of a single electron ( C) using the oil drop experiment. Combining this with Thomson's ratio, the mass of the electron was calculated as g.

Atomic Models

Thomson proposed the "plum-pudding" model, where electrons are embedded in a sphere of positive charge. Rutherford's gold foil experiment disproved this, showing that most of the atom is empty space with a dense, positively charged nucleus.

Neutrons

James Chadwick discovered the neutron, a neutral particle with a mass similar to the proton, accounting for the previously unexplained atomic mass.

Subatomic Particles Table

Atoms are composed of protons, neutrons, and electrons. Their properties are summarized below:

Elements and the Periodic Table

Atomic Number and Chemical Symbols

The atomic number (Z) is the number of protons in an atom's nucleus and defines the element. Each element has a unique chemical symbol.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons. The mass number (A) is the sum of protons and neutrons. Isotopes are represented as or as the chemical symbol/name followed by a dash and the mass number (e.g., Ne-20).

Ions

Atoms can gain or lose electrons to form ions. Cations are positively charged (e.g., Na+), and anions are negatively charged (e.g., F−).

Periodic Law and Table Organization

Elements are arranged in the periodic table by increasing atomic number. Elements with similar properties recur periodically and are grouped in columns (families/groups) and rows (periods).

Classification of Elements

• Metals: Good conductors, malleable, ductile, shiny, tend to lose electrons.

• Nonmetals: Poor conductors, not ductile or malleable, tend to gain electrons.

• Metalloids: Exhibit mixed properties, often semiconductors.

Special Groups

• Noble Gases (Group 8A): Unreactive, stable, e.g., helium, neon, argon.

• Halogens (Group 7A): Very reactive nonmetals, found as salts, e.g., chlorine, bromine, iodine, fluorine.

• Alkali Metals (Group 1A): Highly reactive metals, e.g., sodium, lithium, potassium.

• Alkaline Earth Metals (Group 2A): Fairly reactive metals, e.g., calcium, magnesium.

Ions and the Periodic Table

Main-group metals tend to lose electrons and form cations; main-group nonmetals tend to gain electrons and form anions. The charge of ions is predictable based on group number.

Atomic Mass and Counting Atoms

Atomic Mass

Atomic mass is the weighted average of the masses of an element's isotopes, based on their natural abundance. It is calculated using:

The Mole and Avogadro's Number

The mole (mol) is a counting unit for atoms, defined as particles (Avogadro's number). The mass of 1 mol of atoms (molar mass) in grams is numerically equal to the atomic mass in amu.

Counting Atoms by Weighing

To determine the number of atoms in a sample:

1. Measure the mass of the sample.

2. Convert mass to moles using molar mass.

3. Convert moles to number of atoms using Avogadro's number.

Additional info: These notes cover the foundational concepts of Chapter 2: Atoms and Elements, including historical development, atomic theory, subatomic particles, isotopes, ions, and the organization of the periodic table. They also introduce the mole concept and methods for counting atoms, which are essential for quantitative chemistry.