Atoms and Elements

Leucippus’ and Democritus’ Atomic Philosophy

The concept of the atom originated with Greek philosophers Leucippus and Democritus, who proposed that all matter is composed of tiny, indivisible, and indestructible particles called atoms. However, their ideas were philosophical and lacked experimental evidence or the ability to explain chemical and physical behavior.

Law of Definite Proportions (1797)

The Law of Definite Proportions, formulated by Joseph Proust, states that all samples of a pure compound contain the same elements in the same proportions by mass. This law demonstrates the constant composition of compounds, regardless of their source or method of preparation.

• Definition: The mass ratio of elements in a compound is always the same.

• Example: Isooctane samples always have the same carbon to hydrogen mass ratio.

Law of Multiple Proportions (1804)

John Dalton’s Law of Multiple Proportions states that 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. This law supports the idea of atoms combining in simple, whole-number ratios.

• Example: Nitrogen and oxygen form two compounds: NO and NO2.

John Dalton’s Atomic Theory (1808)

Dalton’s atomic theory provided a scientific basis for the nature of matter. The main postulates are:

1. All elements are composed of tiny, indivisible particles called atoms.

2. Atoms of the same element are identical in their composition.

3. Atoms of different elements can physically mix or chemically combine in simple whole-number ratios to form compounds.

4. Chemical reactions involve the rearrangement of atoms; atoms are not changed into other atoms during chemical reactions.

J.J. Thomson and the Electron (1897)

J.J. Thomson discovered the electron by passing an electric current through gases and observing cathode rays. He found that these rays were negatively charged and present in all atoms, leading to the conclusion that electrons are universal subatomic particles.

J.J. Thomson’s Atomic Theory (1904)

Thomson proposed the "plum pudding" model, where electrons are embedded in a positively charged sphere. This was the first atomic model to include subatomic particles.

Robert Millikan – Determining the Charge of an Electron (1909)

Millikan measured the charge of the electron using the oil drop experiment, finding the charge to be C and the mass g. He showed that electric charge is quantized.

Ernest Rutherford’s Gold Foil Experiment (1909)

Rutherford’s experiment involved firing alpha particles at gold foil. Most particles passed through, but some were deflected, indicating a small, dense, positively charged nucleus at the center of the atom. This disproved the plum pudding model.

Rutherford’s Atomic Theory (1911)

Rutherford proposed that electrons orbit a central nucleus, similar to planets around the sun. This model introduced the concept of the atomic nucleus.

James Chadwick and The Neutron (1934)

James Chadwick discovered the neutron, a neutral subatomic particle found in the nucleus. Neutrons contribute to atomic mass and stability but do not affect chemical properties.

Parts of the Modern Atom

Atoms consist of three main subatomic particles:

• Protons (p+): Determine the element’s identity.

• Neutrons (n0): Contribute to atomic mass and stability.

• Electrons (e-): Responsible for chemical reactivity, found in electron clouds around the nucleus.

Characteristics of the Modern Atom

Atoms are electrically neutral because they contain equal numbers of protons and electrons. Ions are formed by the loss or gain of electrons:

• Cation: Loss of electrons, resulting in a positive charge.

• Anion: Gain of electrons, resulting in a negative charge.

The Periodic Table

History of the Periodic Table

Lothar Meyer and Dmitri Mendeleev independently developed early versions of the periodic table. Mendeleev organized elements by increasing atomic mass and predicted the existence of undiscovered elements. However, some inconsistencies remained when ordering strictly by mass.

Modern Periodic Table

Henry Moseley arranged elements by increasing atomic number, resolving inconsistencies and leading to the modern periodic law: when elements are arranged by atomic number, their properties repeat periodically.

Structure of the Periodic Table

The periodic table is organized into:

• Groups (Families): Vertical columns with similar chemical properties.

• Periods: Horizontal rows.

• Each element is represented by a one- or two-letter symbol (first letter capitalized).

Periodic Table Group Names

Section Names and Classifications

• Transition Metals

• Representative (Main Group) Elements

• Lanthanides

• Actinides

Element Classifications

• Metals: Malleable, ductile, conductive, lustrous; most are solids.

• Nonmetals: Mostly gases at room temperature; solids are brittle and poor conductors.

• Metalloids: Properties intermediate between metals and nonmetals; all are solids, some are semiconductors.

Atomic Structure and Isotopes

Atomic Number

The atomic number (Z) is the number of protons in an atom and determines the element’s identity. In a neutral atom, the number of electrons equals the number of protons.

Mass Number

The mass number (A) is the sum of protons and neutrons in the nucleus:

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers but similar chemical properties.

Proton, Neutron, and Electron Counts

• To find the number of neutrons:

• For ions, the number of electrons differs from the number of protons by the charge.

Atomic Mass (Relative Atomic Mass)

Atomic masses are measured in unified atomic mass units (u), where (1/12 the mass of a C atom). The atomic mass is the weighted average of all naturally occurring isotopes of an element.

Average Atomic Mass

The average atomic mass is calculated using the relative abundances and masses of each isotope:

• Example: Calculate the average atomic mass of sulfur if 95.00% of all sulfur atoms have a mass of 31.972 u, 0.76% have a mass of 32.971 u, and 4.22% have a mass of 33.967 u.

Formula Masses for Compounds

The formula mass of a compound is the sum of the atomic masses of all atoms in its formula. For example:

• XeF4:

• NaCl:

This process applies to all compounds and elements.