1.4 Bonding

• Ionic bond: Strong electrostatic forces of attraction between oppositely charged ions.
• Covalent bond: attraction of the nucleus to the shared pair of electrons
• The charge on ions in each group is the number of electrons it needs to gain/lose to have a full outer electron shell.
• $Ca^{3+}$is unlikely to exist because the ionisation energy to lose the electron from the inner electron shell is too high.
• Ionic bonds have high melting points because they have strong electrostatic forces/ionic bonds which require a lot of energy to break.
• Ionic bonds are brittle: External force moves rows, like charges repel so crystal shatters

  
• Ionic bonds dissolve in water because polar water molecules break the crystal lattice

  
• Ionic bonds act as electrolytes in electrolysis because they are able to conduct electricity when molten or dissolved in water.
• A covalent bond is a shared pair of electrons where one electron comes from each atom
• Electron density - The measure of the probability of an electron being present in a specific location.
• A coordinate covalent bond is a shared pair of electrons where both electrons come from one atom only
• If an atom is electron deficient, it doesn't have a full outer shell of electrons (e.g. BCl3)
• Lone pairs - Electron pairs that are not involved in bonding (they are more negative than bonding pairs as they are only under the influence of one nucleus)
• Differences between covalent and coordinate bonding - no mutually shared electrons Similarities between covalent and coordinate bonding - between two non-metalsCoordinate bonding is a type of covalent bond.
• Electronegativity - The measure of the tendency of an atom to attract a bonding pair of electrons
• Pauling scale measures electronegativity on a scale from 0 to 4, 4 being the most electronegative. Generally small atoms attract electrons more effectively so have greater electronegativity values
• Electronegativity increases across a period
• Electronegativity decreases down a group
• Pure covalent:

  • Non - polar molecules
  • Forms temporary dipoles
  • Van der waal's forces between molecules
  • Difference is zero or very small (less than 0.5)
  • e.g. H-H = 0, C-H = 0.4
• Polar covalent:

  • Polar molecule
  • Forms dipoles (partial charges)
  • Dipole - dipole forces between molecules
  • Difference is less than 2.0, more than 0.5
  • e.g. H-Cl = 1.4
• Ionic bonding:Occurs when there's a large difference in electronegativity (greater than 2.0)e.g. NaCl has a difference in electronegativity of 2.1
• The type of bond that forms is dependent on the difference in electronegativity between the atoms
• Polar covalent - Electrons spend more time around the more electronegative atom. Polarity exists. Dipoles are fromed
• Intra - forces inside compounds

  Inter - forces between molecules
• The forces between molecules govern the physical properties of compounds including:

  • melting point
  • boiling point
  • solubility of a compound

  The stronger the forces between the molecules, the more energy needed to break them, so the higher the melting and boiling points of a substance
• Covalent bonds > Hydrogen bonding > Dipole-dipole >Van der waal's

  Strongest Weakest
• Itermolecular forces are stronger when the molecules are larger (larger Mr) because there are more electrons in the molecules and therefore a stronger intermolecular force and more energy is needed to break them.
• Intermolecular forces are stronger when the molecules have a larger surface area - greater surface area for forces to occur
• A hydrogen bond is a strong intermolecular force of attraction between hydrogen and the lone pairs of a highly electronegative atom (N, O, F) in another molecule.
• Hydrogen bond:

  
• H2O has a higher boiling point than NH3 because it has more lone pairs that it can form hydrogen bonds with.
• Alcohols are soluble because the OH group is able to form hydrogen bonds with the water molecules
• Molecules that demonstrate Van der Waal's forces or dipole-dipole forces DO NOT dissolve in water as they are unable to form hydrogen bonds with water molecules.
• Hydrogen bonds form when a hydrogen atom is attached to a highly electronegative atom (O, N, F,)
• non-miscable - does not mix
• Bonding pairs - Electron pair shared between two atoms (covalent or coordinate)
• Lone pairs - Electron pair that us it used in bonding. Lone pairs are more negative that bonding pairs of electrons so effrect shape of that atom.
• Bonded pair-bonded pair << lone pair-bonded pair << lone pair-lone pair
• Electron pairs repel because they are all negatively charged.
• Vesper - Valence shell electron pair repulsion theory

  Electrons will arrange themselves around a central atom to bring about minimum repulsion between the electron pairs and maximum separation.