3.2 Periodic Table, Periodic Trends & Chemical Bonding

3.2 Periodic Table, Periodic Trends & Chemical Bonding

GED Exam Core Concept: The Periodic Table is not merely a chart of elements—it is an organized roadmap of atomic structure and chemical reactivity. Understanding periodic trends and chemical bonding mechanisms allows you to predict how elements interact, what types of bonds they form, and the physical properties of resulting compounds.

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Organization of the Periodic Table

Modern Periodic Law states that when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic (repeating) pattern.

The periodic table is structured into a grid of horizontal rows and vertical columns:

1. Periods (Horizontal Rows 1 to 7)

• A period represents a horizontal row across the table.
• All elements in the same period have the same number of electron energy levels (shells).
• As you move left to right across a period, atomic number increases by 1, adding one proton to the nucleus and one electron to the valence shell.

2. Groups / Families (Vertical Columns 1 to 18)

• A group (or family) represents a vertical column.
• Elements in the same main-group column share the same number of valence electrons (outermost electrons).
• Because valence electrons dictate chemical reactivity, elements in the same group exhibit similar chemical properties and reactivity patterns.

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Key Chemical Families to Know for GED Science

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Classification of Elements: Metals, Nonmetals & Metalloids

A stair-step diagonal line on the periodic table separates elements into three distinct broad classes:

1. Metals (Left and Center):

   • Make up roughly 80% of all known elements.
   • Properties: Shiny luster, high density, high melting points, excellent conductors of heat and electricity, malleable (can be hammered into flat sheets), and ductile (can be drawn into wires).
   • Tend to lose electrons during chemical reactions to form positive ions (cations).

2. Nonmetals (Upper Right Corner + Hydrogen):

   • Properties: Dull appearance, poor conductors of heat and electricity (insulators), low melting points, brittle solids, liquids, or gases at room temperature.
   • Tend to gain or share electrons during chemical reactions.

3. Metalloids / Semimetals (Bordering the Stair-Step Line: $\text{B, Si, Ge, As, Sb, Te}$):

   • Exhibit intermediate properties between metals and nonmetals.
   • Semiconductors: Silicon ($\text{Si}$) and Germanium ($\text{Ge}$) conduct electricity under specific conditions, making them vital for computer chips and solar cells.

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Periodic Trends

Periodic trends are predictable patterns in element properties across the periodic table governed by two competing atomic forces:

1. Effective Nuclear Charge ($Z_{\text{eff}}$): The attractive pull of protons in the nucleus on surrounding electrons (increases left to right across a period).
2. Electron Shielding & Energy Levels: Inner electron shells shield outer valence electrons from nuclear pull (increases top to bottom down a group).

1. Atomic Radius (Atom Size)

• Across a Period (Left $\rightarrow$ Right): Decreases. Increasing nuclear charge (more protons) pulls electrons closer to the nucleus.
• Down a Group (Top $\rightarrow$ Bottom): Increases. Additional electron shells add volume.

2. Electronegativity (Electron Pull in a Bond)

• Electronegativity measures an atom's ability to attract shared electrons in a chemical bond.
• Across a Period (Left $\rightarrow$ Right): Increases. Atoms get closer to completing an octet and have stronger nuclear pull.
• Down a Group (Top $\rightarrow$ Bottom): Decreases. Outer electrons are farther from the nucleus.
• Most Electronegative Element: Fluorine (F) (value of 4.0). Note: Noble gases generally have no assigned electronegativity because they rarely form bonds.

3. Ionization Energy (Energy to Remove an Electron)

• Ionization Energy is the minimum energy required to remove the outermost electron from a neutral gaseous atom.
• Across a Period (Left $\rightarrow$ Right): Increases. Nonmetals hold electrons tightly.
• Down a Group (Top $\rightarrow$ Bottom): Decreases. Valence electrons are far from the nucleus and easily removed.

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Chemical Bonding & The Octet Rule

Atoms bond with one another to achieve thermodynamic stability. The Octet Rule states that atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons, achieving the stable electron configuration of a noble gas.

There are three fundamental types of chemical bonding:

1. Ionic Bonding (Electron Transfer)

• Occurs Between: A Metal + a Nonmetal.
• Mechanism: The metal atom transfers one or more valence electrons completely to the nonmetal atom. This creates a positive metal cation and a negative nonmetal anion.
• Bonding Force: Strong electrostatic attraction between oppositely charged ions.
• Structure: Forms a continuous 3D crystal lattice rather than isolated individual molecules.
• Properties of Ionic Compounds:
  • High melting and boiling points.
  • Hard and brittle solid crystals.
  • Electrical Conductivity: Do not conduct electricity in solid state, but conduct readily when melted (molten) or dissolved in water (aqueous solution) because ions are free to move.

2. Covalent Bonding (Electron Sharing)

• Occurs Between: Two or more Nonmetals.
• Mechanism: Nonmetal atoms share pairs of valence electrons to achieve stable octets, forming discrete molecules.
• Subtypes based on Electronegativity Difference ($\Delta EN$):
  • Nonpolar Covalent Bond ($\Delta EN < 0.5$): Electrons are shared equally between identical or similar atoms (e.g., $\text{H}_2, \text{O}_2, \text{CH}_4$).
  • Polar Covalent Bond ($0.5 \le \Delta EN \le 1.7$): Electrons are shared unequally. The more electronegative atom pulls shared electrons closer, developing a partial negative charge ($\delta^-$), leaving the other atom with a partial positive charge ($\delta^+$) (e.g., $\text{H}_2\text{O}, \text{HCl}$).
• Properties of Covalent (Molecular) Compounds:
  • Lower melting and boiling points (often liquids or gases at room temperature).
  • Poor conductors of heat and electricity (insulators) in all states.

3. Metallic Bonding ("Sea of Electrons")

• Occurs Between: Metal atoms within pure metals or alloys (e.g., $\text{Cu, Fe, Brass}$).
• Mechanism: Metal cations are arranged in a fixed lattice surrounded by a delocalized "sea of mobile valence electrons" that move freely throughout the metallic structure.
• Properties: Explains why metals conduct electricity, conduct heat, and can be bent or stretched (malleable and ductile) without shattering.

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Chemical Bond Properties Summary Table

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Worked Examples

Worked Example 1: Predicting Bond Type

Problem: Determine the bond type for each of the following chemical formulas: (a) $\text{KCl}$, (b) $\text{CO}_2$, (c) $\text{MgBr}_2$.

Solution:

• (a) $\text{KCl}$ (Potassium Chloride): Potassium ($\text{K}$) is a Group 1 metal; Chlorine ($\text{Cl}$) is a Group 17 nonmetal. Metal + Nonmetal = Ionic Bond.
• (b) $\text{CO}_2$ (Carbon Dioxide): Carbon ($\text{C}$) is a nonmetal; Oxygen ($\text{O}$) is a nonmetal. Nonmetal + Nonmetal = Covalent Bond.
• (c) $\text{MgBr}_2$ (Magnesium Bromide): Magnesium ($\text{Mg}$) is a Group 2 metal; Bromine ($\text{Br}$) is a Group 17 nonmetal. Metal + Nonmetal = Ionic Bond.