Bonding

• Bonding


  The strong electrostatic attraction that holds atoms together in a molecule or compound
• Types of bonding


  • Metallic
  • Ionic
  • Covalent
  • Macromolecular
  • Simple molecular
• Recognising the type of bonding

  1. Single elements from the left of the periodic table

  2. A metal element and a non-metal element

  3. Diamond, graphite, silicon and SiO2

  4. Non-metal elements
• Metallic bonding


  The strong electrostatic attraction of positive metal ions surrounded by a sea of delocalised electrons
• Metallic structure


  • Giant lattice (regular arrangement of particles)
  • Strong metallic bonds
• Metallic bonding examples


  • Sodium
  • Magnesium
• Strength of metallic bonds in Magnesium vs Sodium

  • Mg has a greater charge of 2+
  • Mg has twice as many electrons in the sea of delocalised electrons
  • Mg ions are smaller, meaning there is a greater charge density
  • Therefore the attraction between the Mg2+ ions and the delocalised electrons is stronger
• Properties of metals


  • Conductivity (good electrical and thermal conductors)
  • Strength (majority are very strong due to strong electrostatic attraction)
  • Malleable and ductile (layers of metal ions can slide past one another)
  • High melting and boiling points (linked to strength of metallic bonds)
• Covalent bond


  A shared pair of electrons between two atoms
• Covalent structure


  • Macromolecular or simple molecular
  • Strong covalent bonding
• Macromolecular covalent structures


  • Diamond
  • Graphite
• Simple molecular structure


  Made up of molecules held together by intermolecular forces
• Ionic bond


  The strong electrostatic attraction between oppositely charged ions
• Ionic structure


  • Giant lattice
• Ionic bonding examples


  • Lithium Fluoride
  • Magnesium Oxide
• Physical properties of ionic compounds

  • High melting and boiling points
  • Electrical conductivity (in aqueous or molten state)
  • Brittleness and tendency to shatter easily
• Writing ionic formulas


  1. Write the ions side by side

  2. Draw arrows that cross each other

  3. Write the number of the charge at the arrow end

  4. Write the formula of the ionic compound
• Writing formation equations for ionic compounds

  1. Mg(s) + Cl2(g) MgCl2(S)

  2. 2Na(s) + 1/2O2(g) Na2O (S)
• Coordinate bond


  A shared electron pair which have both come from the same atom
• Valence shell electron pair repulsion (VSEPR) theory

  Bonding pairs and lone pairs of electrons repel each other, determining the shape of molecules and ions
• Molecular shapes based on VSEPR theory

  • Trigonal planar
  • Tetrahedral
  • Trigonal bipyramidal
  • Octahedral
  • V-shaped
  • Pyramidal
  • Seesaw
  • Square pyramidal
• Determining the shape and bond angle of a molecule

  1. Determine the central element

  2. Determine the number of valence electrons

  3. Add 1 for each covalent bond

  4. Add/subtract charges for ions

  5. Divide by 2 to get number of electron pairs

  6. Determine bonding vs lone pairs

  7. Select the appropriate 3D shape
• Determining the number of electron pairs around a central atom

  1. Step 1 - Determine the central atom

  2. Step 2 - Determine how many valence electrons the atom has

  3. Step 3 - Add 1 for every covalent bond the central atom forms

  4. Step 4 - Add 1 for every - charge or take one for every + charge if you have an ion

  5. Step 5 - Divide the total value by 2 to determine the number of electron pairs

  6. Step 6 - Determine how many electron pairs are bonding, and how many are lone pairs

  7. Step 7 - Select the appropriate 3D shape for the molecule
• Covalent bond


  A shared pair of electrons
• Electronegativity


  The power of an atom to attract the pair of electrons in a covalent bond
• The Pauling Scale is used as a measure of electronegativity
• Factors determining electronegativity


  • Nuclear charge (number of protons)
  • Atomic radius (distance between nucleus and outer shell)
  • Shielding (number of electrons between nucleus and outer shell)
• The closer an element is to Fluorine the more electronegative it is
• Across Period 2 the electronegativity increases because the number of protons increases and the shielding remains the same
• Polarity


  Molecules made of atoms with a difference in electronegativity have their electrons distributed unevenly, producing a polar covalent bond
• Determining if a molecule is polar

  1. Is there a difference in electronegativity in the molecule?

  2. Is the central atom only surrounded by the same types of atom?

  3. Are there lone pairs?
• Intermolecular forces


  Forces between molecules that influence melting and boiling points
• Types of intermolecular forces

  • Induced dipole-dipole (Van der Waals) forces
  • Permanent dipole-dipole forces
  • Hydrogen bonding
• Determining the type of intermolecular force a molecule will have


  Consider the difference in electronegativity and presence of lone pairs
• Hydrogen bonding


  The strongest intermolecular attraction, occurs between H (bonded to N, O, or F) and lone pair on a N, O, or F atom on another molecule
• Permanent dipole-dipole forces


  Occur between polar molecules due to difference in electronegativity leading to bond polarity
• Induced dipole-dipole (Van der Waals) forces

  The weakest force but can be stronger than hydrogen bonds and permanent dipole-dipole if the molecule is large, occur between all molecules due to random movement of electrons
• Hydrogen bonding is extremely important despite being a weak bond, as it enables processes like the structure of proteins and DNA
• As the central atoms get bigger in molecules containing hydrogen

  The boiling point increases
• Explaining the trend in melting and boiling points across Period 3

  1. Na to Mg to Al: Increasing metallic bonding strength

  2. Silicon: Highest melting point due to strong covalent bonds

  3. P, S, Cl, Ar: Decreasing intermolecular forces (Van der Waals) from largest to smallest molecules