bonding

• Ionic bonding

  Many, strong electrostatic forces of attraction between oppositely charges ions in a lattice
• Metallic bonding

  Many, strong, electrostatic force of attraction between the positively charged metal ions and the sea of delocalised electrons in a lattice
• Covalent bonding

  A shared pair of electrons between two atoms
• Dative covalent bond

  A covalent bond where the shared pair of electrons are donated by ONE ATOM in the bond
• Crystal structure

  • A solid material made from either atoms, molecules, or ions arranged in a highly ordered three dimension lattice
• Ionic compound formulas

  • Ammonium ion NH4+
  • Hydroxide ion OH-
  • Sulphate ion SO42-
  • Nitrate ion NO3-
  • Hydrogen sulphate ion HSO4-
  • Carbonate ion CO32-
• Ionic substances have high melting points because many of strong electrostatic forces of attraction between oppositely charged ions in a lattice that require a lot of energy to break
• Factors affecting ionic bond strength

  • Charge on the ions - the bigger the charge the stronger the ionic bond
  • Size of the ions - the smaller the ionic radius the stronger the bond
• Solid ionic compounds cannot conduct heat/electricity because their ions are not mobile when bonded in a lattice
• Liquid (molten) or aqueous solutions of ionic compounds can conduct because their ions are mobile
• Solubility rules for ionic compounds

  • Soluble: All NO3– compounds, All NH4+, Na+, K+ compounds, All Cl–, Br– and I– compounds, except with Ag+ and Pb2+, All SO42– compounds, except with Ca2+, Sr2+, Ba2+ and Pb2+
  • Insoluble: All CO32– compounds, except with NH4+, Na+ and K+, All OH – compounds, except with NH4+, Na+ and K+
• Factors affecting metallic bond strength

  • Charge on the ions - the bigger the charge the stronger the electrostatic force of attraction between the positive metal ion and the delocalised electrons
  • Size of the ions - the smaller the stronger the electrostatic force of attraction between the positive metal ion and the delocalised electrons
• Ductility of metals

  Layers of metal ions can roll/slide over each other without breaking the metallic bonds
• Metals are good electrical conductors in solid or liquid state due to the delocalised electrons
• Covalent bond

  Atoms are held together by the balance of electrostatic attractive (electron and nucleus) and repulsive forces (between the two nuclei and between the electrons pair) within a bond
• Macromolecular substances have very high melting points


  A lot of energy is needed to break the many, strong covalent bonds in a lattice
• Diamond

  • Each carbon has four electrons in the outer shell, each carbon makes four covalent bonds, there are no free electrons, so diamond does not conduct electricity
• Graphite

  • Each carbon has four electrons in the outer shell, it makes layers of hexagonal rings, the layers are held together by weak IMF, each carbon makes three covalent bonds, each carbon has one unbonded electron which is delocalised between the layers, so graphite does conduct electricity
• Electron pairs

  Also called charge clouds, can be bonding or non-bonding (lone pairs)
• Electron pairs repel

  Electron pairs move as far apart as possible to minimise repulsion
• Types of repulsion between charge clouds

  • Lone pair - lone pair repulsion is greater than

  Lone pair - bonding pair repulsion which is greater than

  Bonding pair - bonding pair repulsion
• Electronegativity

  The power of an atom to attract a pair of electrons in a covalent bond
• Factors affecting electronegativity

  • Atomic Radius: The bigger the atomic radius the lower the electronegativity

  Nuclear Charge: The greater the number of protons the higher the electronegativity

  Shielding: the greater the shielding the lower the electronegativity
• Electronegativity changes across a period

  Electronegativity increases across a period because the nuclear charge increases, shielding is constant, and atomic radius decreases, therefore greater attraction for the bonding electrons
• Electronegativity changes down a group

  Electronegativity decreases down a group because the atomic radius increases, therefore less attraction for the bonding pair of electrons
• Fluorine is the most electronegative element
• Electronegative elements

  • N, F, O, I, Cl, Br
• Polar bond

  An unequal distribution of the pair of electrons in a covalent bond due to a difference in electronegativity between the atoms
• Non-polar bond

  When the electronegativity of the atoms in the covalent bond are very similar (or the same)
• Molecule with polar bonds not having an overall dipole

  When the molecule is symmetrical and the dipoles cancel
• Intermolecular forces

  Only present in simple molecules (and graphite, macromolecular, has IMFs between the layers)
• Van der Waals forces

  Arise from the uneven distribution of electrons in a molecule at any instant, producing an instantaneous dipole that induces a dipole in neighbouring molecules, with opposite dipoles attracting
• Van der Waals forces increase

  Melting point/boiling point increases
• Factors affecting Van der Waals force strength

  • The bigger the molecule/greater Mr (the more electrons) the stronger the Van der Waals

  Straight chain organic molecules have stronger Van der Waals than branched chain as they pack closer together
• Permanent dipoles

  Arise when the atoms in a bond have different electronegativities, leading to a molecule having an overall permanent dipole (providing the molecule is not symmetrical)
• Hydrogen bonding

  Occurs in compounds that have hydrogen covalently bonded to either nitrogen, oxygen or fluorine, where the ∂+H attracts the lone pair of electrons on a N,F or O in an adjacent molecule (or the same strand in things such as DNA)
• Hydrogen bonding in water and ammonia

  • Water: H-bonding between water molecules

  Ammonia: H-bonding between ammonia molecules
• Hydrogen bonding in ice

  In ice, water molecules form more (four) hydrogen bonds with neighbouring water molecules forming an open lattice structure, so the water molecules in ice are further apart/more spread out, causing ice to float
• Electrical conductivity

  Requires mobile (delocalised) electrons or mobile ions
• Types of crystal structures

  • Ionic

  Metallic

  Macromolecular

  Molecular