The Periodic Table and Periodicity Study Notes

Introduction to the Periodic Table

• Definition: The periodic table is an organized table of chemical elements arranged in order of increasing atomic number.
• Essential Information Provided: The table displays the symbol, full name, mass number, and atomic number for every element.
• Structural Organization:
  • Periods: The horizontal rows in the periodic table.
  • Groups: The vertical columns in the periodic table.
• Predictive Value: The arrangement based on atomic structure and properties allows scientists to predict the chemical properties of an element based solely on its location within the table. Elements in the same group or period share specific characteristics.

Historical Development of the Periodic Table

Throughout the 19th century, chemists categorized elements based on similarities in physical and chemical properties. Key contributors include:

• Antoine Lavoisier (1772–1785):

  • Classified known elements into metals and nonmetals based on physical and chemical reactivity.
  • Nonmetals: Characterized by low melting and boiling points.
  • Metals: Characterized by high melting and boiling points.
  • Other Achievements: Recognized and named oxygen and hydrogen.

• Johann Dobereiner (1829):

  • Proposed the Law of Triads.
  • Observed that certain groups of three elements (triads) exhibited similar chemical and physical properties.
  • Named Triads:
    • Lithium ($Li$), Sodium ($Na$), Potassium ($K$)
    • Chlorine ($Cl$), Bromine ($Br$), Iodine ($I$)
    • Calcium ($Ca$), Strontium ($Sr$), Barium ($Ba$)
    • Iron ($Fe$), Cobalt ($Co$), Manganese ($Mn$)
  • Atomic Mass Relation: He noted that when triads were arranged by increasing relative atomic mass, the mass of the middle element was approximately the average of the other two.

• John Newlands (1865):

  • Proposed the Law of Octaves.
  • Arranged the 56 known elements by increasing atomic weight.
  • Observation: Every element exhibited similar properties to the element eight places ahead of it (e.g., sodium is eight places ahead of lithium).
  • Organizational Contribution: Placed similar elements into vertical columns called groups.
  • Reception: His claim of a repeating pattern was met with "savage ridicule." Scientific peers claimed his classification was as arbitrary as alphabetical order; his paper was rejected by the Chemical Society. His law failed beyond the element calcium.

• Dimitri Mendeleev (1869):

  • Published a Periodic Classification of Elements arranged by increasing relative atomic mass.
  • Placed elements with similar properties into vertical groups.
  • Scientific Breakthroughs:
    • Predicted properties of undiscovered elements and left gaps for them in the table.
    • Occasionally ignored atomic mass order to swap adjacent elements (exchanging them) so they would fit better into chemical families.

• Lothar Meyer (1869–1870):

  • Published a scheme similar to Mendeleev's.
  • Arranged elements by atomic mass and grouped them based on reactivity (combining power) and valency (the number of bonds an element can form).

• Henry Moseley (1914):

  • Arranged elements by increasing atomic number.
  • This arrangement ensured that elements with similar properties fell into the same groups.
  • Proved: Atomic number is a more fundamental property than atomic mass for determining chemical properties.

The Modern Periodic Table Structure

• Scope: Most textbooks show 103 elements, but to date, 118 elements have been discovered or synthesized.
• Natural vs. Synthetic: Elements up to atomic number 98 exist naturally. Elements beyond 98 are synthesized in laboratories or nuclear accelerators.
• Metal/Nonmetal Division:
  • A "step-down" line runs from above Aluminum ($Al$) to Astatine ($At$).
  • Metals: All elements to the left of the line (except Hydrogen).
  • Nonmetals: All elements to the right of the line, including Hydrogen.
• Metalloids (Semi-metals):
  • Elements bordering the division line with properties intermediate between metals and nonmetals.
  • Included Elements: Boron ($B$), Silicon ($Si$), Germanium ($Ge$), Arsenic ($As$), Antimony ($Sb$), Tellurium ($Te$), and Polonium ($Po$).

Physical Properties of Elements

Metals

• Appearance: Lustrous (shiny).
• Mechanical: Malleable (can be hammered) and Ductile (can be stretched into wires).
• Density: High density.
• State: Solid at room temperature (Exception: Mercury ($Hg$) is liquid).
• Thermal/Electrical: Good conductors of heat and electricity.
• Acoustic: Sonorous (makes a ringing sound when hit).
• Thermal points: High melting and boiling points.

Non-metals

• Appearance: Dull in solid state.
• Mechanical: Brittle in solid state.
• Density: Low density.
• State: Usually gases at room temperature; some are solid; Bromine ($Br$) is liquid.
• Thermal/Electrical: Poor conductors of heat and electricity.
• Acoustic: Produce a dull sound when hit.
• Thermal points: Low melting and boiling points.

Group and Period Characteristics

Groups (Vertical Columns)

• There are 8 main groups.
• Valence Electrons: The group number indicates the number of valence electrons (e.g., magnesium in Group 2 has 2 valence electrons).
• Similarity: Elements in the same group have the same number of valence electrons, leading to similar reaction and bonding behaviors (e.g., Sodium and Potassium).
• Vertical Change: Moving down a group, each element adds one more occupied electron shell.
• Transition Metals: There are 10 groups of elements located between Groups 2 and 3.

Periods (Horizontal Rows)

• There are 7 periods.
• Electron Shells: The period number indicates the number of occupied electron shells (e.g., Sodium in Period 3 has 3 shells).
• Valence Placement: All elements in a period have their valence electrons in the same electron shell.
• Horizontal Change: Moving across a period, each element has one more valence electron than the one before it.
• Property Shift: Properties change gradually from metals on the left to nonmetals on the right.
  • Example (Period 3): Sodium ($Na$), Magnesium ($Mg$), and Aluminum ($Al$) are metals (conductors, basic oxides). Silicon ($Si$) is a metalloid/semi-conductor. Phosphorus ($P$), Sulfur ($S$), Chlorine ($Cl$), and Argon ($Ar$) are nonmetals (poor conductors, acidic oxides).

Specific Named Groups

• Group 1: Alkali Metals:
  • Forms alkalis (strong bases) when reacting with water.
  • $Na$ and $K$ are the 6th and 7th most abundant elements.
  • Soft, silvery, and so reactive they are only found in compounds, never free in nature.
• Group 2: Alkaline Earth Metals:
  • Named after their oxides (e.g., beryllia, magnesia, lime).
  • Oxides are basic (alkaline) when combined with water.
  • $Mg$ and $Ca$ are 5th and 8th most abundant elements.
  • Shiny, silvery-white, and highly reactive (found only in compounds).
• Group 7: Halogens:
  • Name derived from Greek hal- ("salt") and -gen ("to produce").
  • Found as nonpolar diatomic molecules.
  • 7 valence electrons; very reactive nonmetals; corrosive and poisonous.
• Group 0: Noble Gases:
  • Also called inert gases; they do not readily react due to full valence shells.
  • Colourless (but fluoresce in lights), odourless.
  • Helium ($He$) is the 2nd most abundant element in the universe; Argon ($Ar$) is the 3rd most abundant gas in Earth's atmosphere; Radon ($Rn$) is radioactive.

Periodicity and Trends

Atomic Radius

• Definition: The shortest distance between the nucleus and the outermost electron shell.
• Forces: Smaller atoms have stronger attraction between the positive nucleus and negative electrons.
• Across a Period: Atomic radius decreases. Why? Number of shells remains the same, but nuclear charge (protons) increases, pulling electrons closer.
• Down a Group: Atomic radius increases. Why? More electron shells are added. The "shielding effect" occurs where inner electrons shield outer ones from the nucleus and repel them.

Electronegativity

• Definition: An atom's affinity for electrons; how strongly it attracts electrons from other atoms.
• Trend: Increases across a period; decreases down a group.
• Mechanism: Smaller atoms have greater attractive force at their boundary.
• Metals vs. Nonmetals: Nonmetals (right side) have high electronegativity and attract/take electrons. Metals (left side) exhibit low electronegativity and tend to lose electrons.
• Exceptions: Noble gases do not follow this trend.

Metallic and Non-Metallic Character

• Metallic Character: Increases down a group, decreases across a period.
• Non-Metallic Character: Decreases down a group, increases across a period.

Reactivity and Ionization

• Ease of Ionization: The ease with which an atom loses or gains electrons to achieve a stable noble gas configuration (full outer shell).
• Metals: Reactivity increases down a group (easier to lose electrons as radius increases) and decreases across a period (radius decreases).
  • Group 2 reactivity order: $Be < Mg < Ca < Sr < Ba < Ra$.
• Non-metals: Reactivity increases up a group (easier to gain electrons as radius decreases) and decreases to the left across a period.
  • Group 7 reactivity order: $At < I < Br < Cl < F$.

Displacement Reactions

• Mechanism: A more reactive element in its free state takes the place of (displaces) a less reactive element in a compound.
• Rules:
  1. More reactive displaces less reactive.
  2. Metals only displace other metals; non-metals only displace other non-metals.
• Examples:
  1. $Potassium + Sodium\,Chloride \rightarrow Potassium\,Chloride + Sodium$$K + NaCl \rightarrow KCl + Na$ (Potassium is more reactive than Sodium).        $Sodium + Potassium\,Chloride \rightarrow No\,Reaction$ (Sodium is less reactive).
  2. $Chlorine + Potassium\,Bromide \rightarrow Potassium\,Chloride + Bromine$$Cl_2 + 2KBr \rightarrow 2KCl + Br_2$ (Chlorine is more reactive than Bromine).

Electronic Configuration Calculation

One can determine configuration based on table position:

1. Group Number = Number of valence electrons.
2. Period Number = Number of occupied electron shells.

• Example (Sulphur):
  • Position: Group VI, Period 3.
  • Valence: 6; Shells: 3.
  • Configuration: Shell 1 (2), Shell 2 (8), Shell 3 (6) → 2, 8, 6.

Questions & Discussion

• Similarity in Group VII? All elements have 7 valence electrons.
• Why does Calcium have a larger radius than Beryllium? Calcium is in Period 4 (4 shells) while Beryllium is in Period 2 (2 shells); more shells increase the distance from the nucleus.
• Atomic Number 14 (Element Z)? Configuration 2, 8, 4. It is in Period 3 and Group IV.
• Atomic Number 20? Configuration 2, 8, 8, 2. It is a Period 4, Group II metal (Calcium).
• Why is Chlorine more electronegative than Sodium? Chlorine is smaller and has a higher nuclear charge than Sodium in the same period, allowing it to attract electrons more strongly.
• Which reacts more vigorously with HCl, Magnesium or Calcium? Calcium, because reactivity of metals increases down Group 2.