02.00. Chapter Summary

BioCast:

1. Elements, Compounds & Mixtures

• Classification: All substances are classified as elements, compounds, or mixtures.
• Element:
  • Definition: Pure substances consisting of only one type of atom with the same number of protons.
  • Example: Oxygen (O), Carbon (C).
  • Periodic Table: 118 elements are listed.
• Compound:
  • Definition: Pure substances formed by chemically combining two or more elements in fixed proportions.
  • Examples:
    • Copper(II) sulfate (CuSO₄)
    • Calcium carbonate (CaCO₃)
    • Carbon dioxide (CO₂)
  • Properties: Cannot be separated by physical means.
• Mixture:
  • Definition: Combination of two or more substances (elements or compounds) physically mixed, not chemically bonded.
  • Examples: Sand and water, oil and water, sulfur powder and iron filings.
  • Properties: Can be separated by physical methods like filtration or evaporation.
• Diagrams:
  • Particle Diagrams: Elements, compounds, and mixtures at the particle level.

2. Atomic Structure and the Periodic Table

2.1 Structure of the Atom

• Basic Concepts:
  • Atom: Building block of matter, composed of protons, neutrons, and electrons.
  • Subatomic Particles:
    • Protons:
      • Charge: +1
      • Relative Mass: 1
    • Neutrons:
      • Charge: 0 (neutral)
      • Relative Mass: 1
    • Electrons:
      • Charge: -1
      • Relative Mass: ≈ 0 (negligible)
  • Structure:
    • Nucleus: Contains protons and neutrons.
    • Electron Shells: Electrons orbit the nucleus in energy levels (shells).

• Key Definitions:
  • Atomic Number (Z): Number of protons in an atom; determines the element.
  • Mass Number (A): Total number of protons and neutrons.
  • Isotopes: Atoms of the same element with different numbers of neutrons.
• Diagram:
  • Atomic Structure Diagrams: Protons, neutrons, and electrons within an atom. Notice that there are two ways that this information can be displayed.

• Periodic Table Layout: Illustrates elements with atomic number (Z) and relative atomic mass (Ar).

2.2 Electronic Configuration

• Concept:
  • Electronic Configuration: Arrangement of electrons in an atom’s shells.
  • Notation: Numbers separated by commas indicate electrons in each shell (e.g., Carbon: 2,4).
• Rules:
  • Electrons fill the lowest available energy levels first.
  • Shell Capacities:
    • 1st shell: 2 electrons
    • 2nd shell: 8 electrons
    • 3rd shell: 8 electrons (simplified for IGCSE)

• Periodic Table Relationship:
  • Periods: Indicate the number of occupied shells.
  • Groups: Indicate the number of valence (outer) electrons.
  • Group VIII Noble Gases: Have a full outer shell.
  • Groups I to VII: Number of outer shell electrons corresponds to the group number.

• Examples:
  • Magnesium (Mg): 2,8,2
  • Carbon (C): 2,4

• Diagrams:
  • Electronic Configuration Notation: Represent electron distribution in shells.

3. Isotopes and Relative Atomic Mass

3.1 Isotopes

• Definition: Different atoms of the same element that have the same number of protons but different numbers of neutrons.
• Symbol Representation:
  • Standard Format: Element symbol with mass number as a superscript and atomic number as a subscript

• Example:
  • Carbon-12: ^12_6C
  • Carbon-13: ^13_6C

• Chemical Properties:
  • Same Chemical Properties: Isotopes have identical chemical behavior due to identical electronic configurations.
  • Different Physical Properties: Vary in mass, affecting properties like density and melting point.

3.2 Relative Atomic Mass (Ar)

• Definition: Weighted average mass of an element’s isotopes compared to 1/12th the mass of carbon-12.
• Calculation:

• Example:
  • Rubidium:

• Important Notes:
  • Difference from Mass Number: Relative atomic mass accounts for all isotopes, whereas mass number refers to a specific isotope.
  • No Units: It is a ratio and thus has no units.

4. Ions & Ionic Bonds

4.1 Formation of Ions

• Cations: Positive ions formed by losing electrons (typically metals).
  • Example: Sodium ion (Na⁺)
• Anions: Negative ions formed by gaining electrons (typically non-metals).
  • Example: Chloride ion (Cl⁻)

4.2 Ionic Bonding

• Definition: Strong electrostatic attraction between oppositely charged ions.
• Formation:
  • Process: Transfer of electrons from metal to non-metal.
  • Example:
    • Sodium Chloride (NaCl):
      • Sodium (Na) loses one electron to form Na⁺.
      • Chlorine (Cl) gains one electron to form Cl⁻.
      • Na⁺ and Cl⁻ attract to form NaCl.

• Dot-and-Cross Diagrams:
  • Ionic Bonds: Showing ions with charges enclosed in brackets.
  • Example: NaCl

4.3 Properties of Ionic Compounds

• Physical State: Usually solid at room temperature.
• Melting & Boiling Points: High due to strong electrostatic forces in the lattice.
• Electrical Conductivity:
  • Solid State: Poor conductors.
  • Molten/Aqueous State: Good conductors due to free-moving ions.
• Structure:
  • Giant Lattice Structure: Regular arrangement of alternating positive and negative ions.
  • Example: NaCl lattice.

5. Simple Molecules and Covalent Bonds

5.1 Formation of Covalent Bonds

• Definition: Sharing of electron pairs between non-metal atoms to achieve noble gas electronic configurations.
• Molecules: Two or more atoms bonded covalently (e.g., H₂O, CO₂).
• Types of Covalent Bonds:
  • Single Bond: Sharing one pair of electrons (e.g., H₂).
  • Double Bond: Sharing two pairs of electrons (e.g., O₂).
  • Triple Bond: Sharing three pairs of electrons (e.g., N₂).

5.2 Formation in Simple Molecules

• Examples and Dot-and-Cross Diagrams:
  • Hydrogen (H₂)
  • Chlorine (Cl₂)
  • Water (H₂O)
  • Methane (CH₄)
  • Ammonia (NH₃)
  • Hydrogen Chloride (HCl)

• Supplementary Examples:
  • Methanol (CH₃OH)
  • Ethene (C₂H₄)
  • Carbon Dioxide (CO₂)
  • Nitrogen (N₂)

5.3 Properties of Simple Molecular Compounds

• Physical State: Often gases or liquids at room temperature.
• Melting & Boiling Points: Generally low due to weak intermolecular forces.
• Electrical Conductivity: Poor conductors as they lack free ions or electrons.
• Examples: Water (H₂O), Methane (CH₄), Ammonia (NH₃), Hydrogen Chloride (HCl).

6. Giant Covalent Structures

6.1 Graphite and Diamond

• Graphite:
  • Structure: Layers of carbon atoms bonded in hexagons with delocalized electrons.
  • Properties:
    • Conducts electricity (due to delocalized electrons).
    • Slippery (layers can slide over each other).
    • High melting point.
  • Uses: Pencils, lubricants, electrodes.
• Diamond:
  • Structure: Each carbon atom bonded to four others in a tetrahedral lattice.
  • Properties:
    • Extremely hard.
    • High melting point.
    • Does not conduct electricity.
  • Uses: Cutting tools, jewelry.

6.2 Silicon(IV) Oxide (SiO₂)

• Structure: Each silicon atom bonded to four oxygen atoms, forming a tetrahedral network similar to diamond.
• Properties:
  • Very hard.
  • High melting point.
  • Insoluble in water.
• Uses: Sandpaper, furnace linings, glass production.

6.3 Similarities Between Diamond and SiO₂

• Properties:
  • Both have giant covalent structures.
  • Both are very hard and have high melting points.
  • Both are insoluble in water.
  • Neither conducts electricity.

7. Metallic Bonding

7.1 Structure of Metals

• Metal Lattice: Positive metal ions surrounded by a “sea of delocalized electrons.”
• Bonding: Electrostatic attraction between metal ions and free electrons.

7.2 Properties of Metals

• Physical State: Mostly solid at room temperature (except mercury).
• Melting & Boiling Points: Generally high due to strong metallic bonds.
• Electrical & Thermal Conductivity: Excellent conductors (due to delocalized electrons).
• Malleability & Ductility:
  • Malleability: Can be hammered into sheets.
  • Ductility: Can be drawn into wires.
• Examples: Iron, Copper, Aluminum.
• Diagrams:
  • Metallic Bonding Diagram: Shows metal ions in a lattice with delocalized electrons.

8. Electrical Conductivity in Compounds

• Ionic Compounds:
  • Solid State: Poor conductors.
  • Molten/Aqueous State: Good conductors (free-moving ions).
• Covalent Compounds:
  • All States: Poor conductors (lack free ions or electrons).
• Metals:
  • All States: Good conductors (delocalized electrons).
• Examples of Insulators:
  • Covalent Compounds: Plastic, rubber, wood.
  • Explanation: Lack free charge carriers.

9. Worked Examples

1. Determining Atomic Particles:
   • Example: Element X with atomic number 29 and mass number 63.
     • Protons: 29
     • Electrons: 29 (neutral atom)
     • Neutrons: 63 – 29 = 34
2. Calculating Relative Atomic Mass:
   • Rubidium Example:

1. Electronic Configuration:
   • Magnesium (12 electrons): 2,8,2
2. Ionic Bond Formation:
   • Sodium Chloride (NaCl):
     • Na: 2,8,1 → Na⁺: 2,8
     • Cl: 2,8,7 → Cl⁻: 2,8,8
     • Bond: Na⁺ and Cl⁻ form NaCl.
3. Covalent Bond Formation:
   • Water (H₂O):
     • H: 1 electron each
     • O: 6 electrons
     • Configuration: H:O:H with shared electrons.

10. Exam Advice

• Electron Configurations:
  • Master writing electronic configurations for the first 20 elements.
  • Recognize patterns relating to periods and groups on the Periodic Table.

• Ion Formation:
  • Use group numbers to determine ion charges:
    • Group 1: +1
    • Group 2: +2
    • Group 6: -2
    • Group 7: -1
    • Group VIII: Noble gases (full outer shells).

• Bonding Diagrams:
  • Ionic Bonds: Show ions with charges and use brackets.
  • Covalent Bonds: Use dot-and-cross diagrams to represent shared electrons.

• Properties Correlation:
  • Link bonding types to physical properties (e.g., lattice structures to high melting points in ionic compounds).

• Avoid Common Mistakes:
  • Do not confuse mass number with relative atomic mass.
  • Remember that graphite is a form of carbon, not the metal lead.

• Terminology:
  • Valency Electrons: Electrons in the outer shell that determine bonding.
  • Nucleons: Protons and neutrons collectively.