Basic Periodic Table Organization and Elemental Properties Study Guide

The Historical Development of the Periodic Table

• Dmitri Mendeleev: A Russian scientist who, around the year $1869$, observed a repetition in the chemical properties of elements (how they react with other chemicals).

  • Organizational Method: He organized the approximately $60$ known elements of his time into groups based on similar chemical properties.
  • Atomic Mass: Because the method for identifying atomic mass had recently been discovered, Mendeleev laid out the elements in order of increasing atomic mass.
  • Predictive Power: He noticed that chemical properties repeated at regular intervals. This led him to leave blank spaces in his table for elements that had not yet been discovered. He was able to accurately predict the properties of these missing elements based on their position relative to known elements.

• Henry Moseley: An English scientist who made the next significant contribution in $1911$.

  • Organizational Shift: Moseley refined Mendeleev's work by organizing elements in order of increasing atomic number rather than atomic mass. This adjustment switched only a small number of Mendeleev's elements but increased the overall accuracy of the table.

• The Periodic Law: Derived from Moseley's work, this law states that when elements are organized by atomic number, there is a periodic repetition of their physical and chemical properties.

  • Definition of "Periodic": In this context, "periodic" means that the elements follow a specific pattern of properties until they reach a certain point, after which the pattern repeats.

Components of a Periodic Table Cell

• The Cell: This is the individual box on the periodic table that contains data for a specific element.
• Atomic Number: Located at the top of the cell.
  • It represents the total number of protons in an atom.
  • The atomic number is a constant identifier for an element; for example, Silver will always have the atomic number $47$. This number is what defines the identity of the atom.
• Chemical Symbol: A shorthand abbreviation for the element.
  • Length: Most symbols are one or two letters long.
  • Notation Rules: The first letter of a chemical symbol is always capitalized. If there is a second or third letter, it must be lowercase. Following this notation is critical for distinguishing between pure elements and compounds (substances containing multiple elements).
  • Linguistic Origins: Some symbols are derived from Latin names. For example, the symbol for Silver is $Ag$, which comes from its Latin name and does not resemble the English name.
• Average Atomic Mass: Located at the bottom of the cell, usually in decimal form.
  • This value is based on the variety of isotopes for that element, their respective masses, and their relative abundance in nature.

Reference Tables and Classroom Procedures

• Usage of Materials: Students are instructed to keep two copies of the chemistry reference tables.
  • Marked Copy: One copy should be used as a working reference to be "marked up" with notes and labels during lessons.
  • Blank Copy: The second copy must be kept blank.
• Testing Protocol: In the North Carolina Final Exam ($NCFE$) and during classroom tests or quizzes, students are provided with a blank table. Using a blank table for practice ensures students know how to extract information without the aid of personal notes.

Major Sections and the Stair-Step Divider

• The Stair-Step: This refers to the zig-zag line on the periodic table that serves as the primary divider for elemental classification.
• Metals:
  • Location: Most elements to the left of the stair-step (including those pull-out rows at the bottom).
  • Exceptions: Hydrogen ($H$), located in the upper left-hand corner, is a nonmetal.
  • Physical Properties:
    • Malleable: They can be pounded or rolled into thin sheets (e.g., aluminum foil).
    • Ductile: They can be drawn into thin wires (e.g., copper wire).
    • Conductors: They are excellent conductors of heat and electricity.
    • State of Matter: Most are solid at room temperature. The primary exception is Mercury ($Hg$), which is liquid at room temperature.
• Nonmetals:
  • Location: Generally located to the right of the stair-step, plus Hydrogen ($H$).
  • Physical Properties:
    • Brittle: They cannot be shaped like metals; if pounded, they will shatter or break.
    • Conductors: They are poor conductors of both heat and electricity.
    • State of Matter: Most are gases at room temperature, with a few exceptions.
• Metalloids:
  • Location: These are the elements that directly touch the stair-step.
  • Exceptions: Aluminum ($Al$) touches the stair-step but is classified as a metal.
  • Specific Elements: Boron ($B$), Silicon ($Si$), Germanium ($Ge$), Arsenic ($As$), Antimony ($Sb$), Tellurium ($Te$), Polonium ($Po$), and Astatine ($At$).
  • Properties: They are semiconductors, meaning their electrical conductivity falls between that of metals and nonmetals. This makes them ideal for use in electronics.

Organizations of the Table: Periods and Groups

• Periods: The horizontal rows on the periodic table.
  • Numbering: There are $7$ labeled periods.
  • Trends: As you move from left to right across a period, atomic number increases, and chemical properties change systematically. For example, a period might begin with a highly metallic metal, lose metallic properties as you move right, and end with a nonmetal. The pattern then resets at the start of the next period.
  • Period 1: Includes only Hydrogen ($H$) and Helium ($He$).
  • Periods 6 and 7: Include the "pull-out" rows at the bottom of the table.
• Groups (Families): The vertical columns on the periodic table.
  • Properties: Elements in the same group share very similar chemical properties.
  • Numbering Systems:
    • Modern: Labeled $1$ through $18$ from left to right.
    • Older System: Uses numbers and letters ($1A$ through $8A$ and $1B$ through $8B$).

Specific Elemental Groups and Their Properties

• Group 1 (Group 1A) - Alkali Metals:

  • Inclusion: Includes the metals in the first column, excluding Hydrogen.
  • Physical Properties: Silver-colored and very soft; they are soft enough to be cut with a knife.
  • Reactivity: They are the most reactive metals on the table. They are never found in their pure elemental form in nature.
  • Handling: They oxidize (react with oxygen) in air. In laboratories, they are stored under oil or kerosene to prevent them from reacting with the atmosphere.
  • Water Reaction: All alkali metals react with water to produce hydrogen gas ($H_2$).
  • Electrons: They easily lose $1$ electron, which accounts for their high reactivity.

• Group 2 (Group 2A) - Alkaline Earth Metals:

  • Physical Properties: Silver-colored and harder than the alkali metals.
  • Reactivity: They are reactive and are not found pure in nature, though they are less reactive than Group $1$.
  • Electrons: They lose $2$ electrons when bonding.

• Groups 3–12 (B Groups) - Transition Metals:

  • Location: The middle block of the table.
  • Reactivity: Less reactive than Group $1$ and $2$; consequently, they can be found in nature in their pure form.
  • Physical Properties: High luster (very shiny) and excellent conductors of electricity.
  • Conductivity Examples: Copper ($Cu$), Silver ($Ag$), and Gold ($Au$).
  • Magnetism: Particular transition metals are magnetic, specifically Iron ($Fe$), Cobalt ($Co$), and Nickel ($Ni$).
  • Compounds: When they form compounds, they often produce very bright, vibrant colors (e.g., cobalt compounds are often a deep blue). These are frequently used in paints.

• Inner Transition Metals:

  • Location: The two rows pulled out at the bottom. They technically belong in periods $6$ and $7$, inserted between the second and third columns. They are moved to the bottom to keep the table compact.
  • Lanthanides: The first row (Period $6$), atomic numbers $57$ through $71$. They have properties similar to Group $2$ metals.
  • Actinides: The second row (Period $7$), atomic numbers $89$ through $103$. Most are man-made (synthetic) and not found in nature. They have applications in nuclear power plants and nuclear weapons.

• Group 17 (Group 7A) - Halogens:

  • Reactivity: The most reactive nonmetals because they easily gain $1$ electron.
  • Diatomic Nature: In their elemental form in nature, they exist in pairs (e.g., $F_2$, $Cl_2$).
  • Salt Formers: The name "halogen" means "salt former." They react with metals to form salts, such as Sodium Chloride ($NaCl$), formed from Chlorine gas and Sodium.

• Group 18 (Group 8A) - Noble Gases:

  • Properties: Extremely stable and chemically inert (they do not naturally bond with other elements).
  • Safety: They are non-flammable and non-reactive.
  • Applications: When their electrons are excited, they produce distinct colors of light. "Neon" lights may contain Neon ($Ne$) or other noble gases depending on the desired color. Their unreactivity makes them useful in various industrial applications.