C2.2 and C2.3: Atomic Structure, Bonding, and Properties of Matter Study Guide

Atomic Structure and Fundamentals of Matter

• Determining Atomic Components:

  • The atomic identity and mass are used to calculate the subatomic particles present in an atom:
    • Number of Protons: Equivalent to the Atomic Number.
    • Number of Electrons: Equivalent to the Atomic Number (in a neutral atom).
    • Number of Neutrons: Calculated by subtracting the atomic number from the atomic mass: $\text{Atomic mass} - \text{Atomic number} = \text{Number of Neutrons}$.
    • Example Calculation provided: $14 - 7 = 7$, indicating an atom with 7 protons, 7 neutrons, and 7 electrons.

• Steps for Drawing Atomic Structure:

  • Step 1: Identify the exact number of protons, electrons, and neutrons specific to the element.
  • Step 2: Place the correct total of protons and neutrons within the central nucleus.
  • Step 3: Draw the circular electron shells around the nucleus and distribute the electrons correctly into each shell according to the electronic configuration rules.

The Periodic Table and Electronic Structure

• Organization of the Periodic Table:
  • The table is organized into Groups (vertical columns) and Periods (horizontal rows).
  • Groups:
    • Group 1: Alkali Metals (e.g., $Li$, $Na$, $K$, $Rb$, $Cs$, $Fr$).
    • Group 2: Alkali Earth Metals (e.g., $Be$, $Mg$, $Ca$, $Sr$, $Ba$, $Ra$).
    • Group 3-12: Transition Metals (e.g., $Sc$, $Ti$, $V$, $Cr$, $Mn$, $Fe$, $Co$).
    • Group 17: Halogens (e.g., $F$, $Cl$, $Br$, $I$).
    • Group 18 (or 0): Noble Gases (e.g., $He$, $Ne$, $Ar$, $Kr$, $Xe$, $Rn$).
  • Specific Series:
    • Lanthanides: Elements 57–71.
    • Actinides: Elements 89–103.

Chemical Bonding: Ionic Compounds

• Formation of Ionic Bonds:

  • Ionic bonds are formed by the reaction between a metal and a non-metal.
  • Mechanism: It involves the transfer of electrons to achieve a full outer shell.
  • Case Study: Sodium Chloride ($NaCl$):
    • Sodium ($Na$): Has an electronic configuration of $2.8.1$. It has 1 electron in its outer shell. By LOSING this electron, it forms a positive ion ($Na^+$) with a filled outer shell ($[2.8]^+$).
    • Chlorine ($Cl$): Has an electronic configuration of $2.8.7$. It has 7 electrons in its outer shell. By GAINING an electron from sodium, it forms a negative ion ($Cl^-$) with a filled outer shell ($[2.8.8]^-$).

• Ionic Compound Structure:

  • Ionic compounds are composed of ions joined together in a regular arrangement. Positive and negative ions attract each other through strong electrostatic forces of attraction.

Chemical Bonding: Covalent Interactions

• Forces in Covalent Bonding:

  • Covalent bonding involves the electrostatic forces of attraction between the POSITIVE NUCLEUS of each atom and the SHARED ELECTRONS.
  • Bond Strength: The covalent bonds themselves are very STRONG.
  • Intermolecular Forces: Despite strong bonds within molecules, there are WEAK forces between the molecules, known as intermolecular forces.

• Structural Representation Models:

  • Space-filling models: Use circular shapes/spheres to represent the relative volume and the arrangement of atoms.
  • Ball and stick models: Circular balls represent atoms, while long sticks represent chemical bonds.
  • Displayed formula: Uses chemical symbols for atoms and lines (single, double, or triple) to represent covalent bonds.

Giant Covalent Structures and Allotropes

• Properties of Giant Covalent Structures:

  • Consist of a vast amount of non-metal atoms joined by covalent bonds in a repeated, regular pattern.
  • They do not form individual molecules but vast NETWORKS of atoms.
  • Physical Characteristics: Very HIGH melting and boiling points, and they are usually very HARD.

• Carbon Allotropes:

  • Allotropes are defined as different forms of the same element.
  • Diamond and Graphite: Both are made of Carbon ($C$) atoms. They exhibit different properties because the atoms are bonded in different arrangements, creating distinct giant structures.

• Silicon Dioxide (Sand/Quartz):

  • Sand is largely composed of the mineral Quartz ($SiO_2$).
  • Structure: A giant covalent lattice where each Silicon ($Si$) atom (configuration $2.8.4$) is bonded to four Oxygen ($O$) atoms. Conversely, each Oxygen atom (configuration $2.6$) is bonded to two Silicon atoms.

Polymers

• Etymology and Definition:
  • Derived from the Greek words "poly" (meaning many) and "meros" (meaning parts).
  • Natural Polymers Examples:
    • Keratin: Found in hair and fur.
    • Proteins: Large molecules made of amino acids.
    • Starch: A carbohydrate molecule made of sugar subunits.

Metallic Bonding

• Structure of Metals:
  • The outer electrons of metal atoms detach to form a "sea of electrons."
  • These electrons are DELOCALIZED, meaning they are free to move throughout the entire structure.
  • The metal atoms become positively charged ions and are held in a regular LATTICE structure.
  • Bonding Definition: The attraction between the positive ions and the sea of delocalized electrons is called metallic bonding.

Fullerenes

• Buckyballs ($C_{60}$):

  • Resemble a sheet of graphene closed into a hollow ball.
  • Contains carbon atoms arranged in pentagons as well as hexagons.
  • Conductivity: Buckyballs do NOT conduct electricity because their ball structure does not allow electrons to flow freely.

• Practical Applications of Fullerenes:

  • Lubricants: Non-stick, slippery coatings for machinery, functioning like miniature ball bearings.
  • Drug Delivery: Act as "cages" to hold drug molecules for direct delivery into the body.
  • Molecular Sieves: Used to trap large particles (like viruses) while allowing smaller, healthy particles to pass through.
  • Chemical Sponges: Used to soak up toxic substances within the body.

States of Matter and Energy Changes

• Phase Transitions:

  • Melting: Transition from Solid to Liquid.
  • Boiling: Transition from Liquid to Gas.
  • Freezing: Transition from Liquid to Solid.
  • Condensing: Transition from Gas to Liquid.

• Bonding and Physical Change:

  • When a substance changes state, forces of attraction between particles must be broken.
  • Solid to Liquid: Some forces or bonds break.
  • Liquid to Gas: ALL remaining bonds break.

• State Determinants:

  • Substances with high melting/boiling points (Metals, Ionic compounds, Giant Covalent structures) are typically SOLID at room temperature.

Physical Properties: Malleability, Brittleness, and Conductivity

• Malleability in Metals:

  • Metals can be shaped because their ions are arranged in layers within the lattice.
  • Delocalized electrons allow these layers to SLIDE over one another when force is applied.
  • Because the sea of electrons remains intact, NO bonds are broken during this displacement.

• Brittleness:

  • Giant Covalent Structures (e.g., Glass): Atoms are in fixed positions held by strong bonds. When force is applied, atoms cannot slide; instead, many covalent bonds break at once, causing the substance to shatter.
  • Ionic Compounds: These are also usually brittle.

• Flexibility:

  • Simple molecules and polymers are flexible because they are held together by WEAK intermolecular forces.
  • If atoms are not arranged in a rigid lattice, the shape can be altered easily.

• Electrical Conductivity:

  • Conductivity requires charged particles (electrons or ions) that are FREE to move.
  • Metals: Good conductors due to delocalized electrons.
  • Giant Covalent Structures: Poor conductors because they lack delocalized electrons.
  • Ionic Compounds:
    • Solid state: Poor conductors; ions are fixed in place.
    • Molten/Dissolved state: Good conductors; ions are free to move.