Subatomic Particles and Atomic Number

Subatomic Particles

• J.J. Thomson discovered the electron in 1897 and proposed the Plum Pudding Model.
• Robert Millikan's oil-drop experiment in 1909 measured the electron's charge.
• Ernest Rutherford's gold foil experiment in 1909 led to the discovery of the nucleus.
• Rutherford proposed the Nuclear Model (or Planetary Model) with a dense, positively charged nucleus and surrounding electrons.
• Eugen Goldstein in 1886 discovered canal rays, indicating positive charges.
• Rutherford in 1917 identified the proton as a positively charged particle in the nucleus.
• James Chadwick in 1932 discovered the neutron, explaining the missing mass.
• All atoms are composed of Protons, Neutrons, and Electrons.
• Protons and neutrons have nearly identical masses.
• The charge of the proton and the electron are equal in magnitude but opposite in sign.
• The neutron has no charge.

Atomic Number

• The number of protons in an atom’s nucleus is its atomic number and is given the symbol Z.
• The number of protons defines the element, because every element has a different number of protons.

Mass Number

• Mass number is the sum of the number of neutrons and protons in an atom and is represented by the symbol A.
• A = number of protons (p) + number of neutrons (n)
• A=p+n

Isotopes

• All atoms of a given element have the same number of protons.
• Atoms with the same number of protons but a different number of neutrons are called isotopes.
• The relative amount of each different isotope in a naturally occurring sample of a given element is roughly constant.
• These percentages are called the natural abundance of the isotopes.

Ions

• The number of electrons in a neutral atom is equal to the number of protons in its nucleus (designated by its atomic number Z).
• During chemical change, atoms can lose or gain electrons and become charged particles called ions.
• Positively charged ions, such as are called cations.
• Negatively charged ions, such as are called anions.

Periodic Law and the Periodic Table

• In 1869, Dmitri Mendeleev noticed that certain groups of elements had similar properties.
• Mendeleev observed that when elements are listed in order of increasing mass, these similar properties recurred in a periodic pattern.
• Mendeleev summarized these observations in the periodic law: When the elements are arranged in order of increasing mass, certain sets of properties recur periodically.
• Mendeleev arranged the known elements in a table and placed elements with similar properties in the same vertical columns.
• Mendeleev's table contained gaps, which allowed him to predict the existence (and even the properties) of yet undiscovered elements.
• In 1886, eka-silicon was discovered by Clemens Winkler, who named it germanium.

Modern Periodic Table

• In the modern table, elements are listed in order of increasing atomic number rather than increasing relative mass.
• The modern periodic table also contains more elements than Mendeleev’s original table.
• The periodic table can also be divided into main-group elements and transition elements or transition metals.
• The periodic table has 18 vertical columns called groups and 7 horizontal rows called periods.
• Main-group elements are in columns 1A–8A (groups 1, 2, 13–18).
• Subgroup elements or transition metals are in columns 3–12 (B groups).

Classification of Elements

• Elements are broadly classified as metals, nonmetals, and metalloids.
• Metals: Lower-left side and middle of the periodic table; good conductors of heat and electricity, malleable, ductile, often shiny, tend to lose electrons.
• Nonmetals: Upper-right side of the periodic table; varied states at room temperature (solids, liquids, gases), poor conductors of heat and electricity, not ductile or malleable, tend to gain electrons.
• Metalloids (Semimetals): Along the zigzag line dividing metals and nonmetals; exhibit mixed properties of metals and nonmetals, often semiconductors with temperature-dependent conductivity.

Group Classifications

• Alkali Metals (Group 1A): Highly reactive metals (e.g., Lithium, sodium, potassium, rubidium, cesium).
• Alkaline Earth Metals (Group 2A): Reactive, but less so than alkali metals (e.g., Magnesium, Calcium, Strontium, Barium).
• Halogens (Group 7A): Very reactive nonmetals (e.g., Fluorine, Chlorine, Bromine, Iodine).
• Noble Gases (Group 8A): Mostly unreactive, chemically stable (e.g., Helium, Neon, Argon, Krypton, Xenon).

Ions and the Periodic Table

• A main-group metal tends to lose electrons, forming a cation with the same number of electrons as the nearest noble gas.
• A main-group nonmetal tends to gain electrons, forming an anion with the same number of electrons as the nearest noble gas.