Bonding

There are 3 things that give evidence of ionic bonding:

• Physical properties of ionic compounds: high melting temperatures, soluble in polar solvents, conduct electricity when molten or in aqeuous solution.
• Electron density maps of ionic compounds shows zero electron density between ions- meaning complete electron transfer.
• Migration of ions in electrolysis- Electrolysis of green copper chromate attracts yellow chromate ions to anode and blue copper ions to cathode. 

• An ion is formed when an atom gains or loses one or more electrons.
• A positive ion is called a cation, a negative ion is called an anion.
• When ions are formed, they tend to have a full outer shell of electrons. This is the octet rule.
• Two ions/atoms that have the same electronic configuration are isoelectronic. 

• An ionic bond is an omnidirectional electrostatic force of attraction between oppositely charged ions.
• The forces of attraction act equally in all directions.
• In ionic compounds each ion is surrounded by ions of the opposite charge.
• Ionic compounds form giant ionic lattices in the solid state. 

• The ionic radius is the radius of an ion in crystal form.
• Cations are smaller than the original atom because the atom loses electrons. It is the opposite for anions.
• Down the group, ionic radius increases because the number of shells being used increases.

• Isoelectronic ions have different atomic radii. 
• The additional electrons in an anion makes the ion larger because there is greater repulsion and the electrons are less tightly bound to the atom.
• The loss of electrons in a cation means the remaining electrons are attracted more closely to the nucleus.
• For isoelectronic anion, atom and cation: 
  anion>atom>cation for ionic radius. 

• The formation of an ionic crystal from its elements is exothermic (energy is released).
• The lattice energy is the energy released when 1 mole of an ionic crystal is formed from its ions in the gaseous state, under standard conditions. 
• This process can be broken down into a number of stages. The application of these stages to find the lattice energy of a crystal is called a Born-Haber cycle. 

• Born-Haber cycles can be used to predict the relative stabilities of ionic compounds, or even if a particular formula will exist as a compound.
• The most stable compound is that with the most exothermic enthalpy of formation.
• An ionic compound with an endothermic enthalpy of formation cannot exist. 

• Born-Haber is an experimental way of calculating lattice energies.
• Lattice energies can also be calculated using Coulomb's Law (electrostatic attraction), which assumes complete electron transfer in ionic compounds. This is theoretical.
• In some ionic compounds, the theoretical lattice energies are smaller than the experimental energies.
• This suggests the bonding in some ionic compounds is stronger than theory would suggest, it is not completely ionic, it has some covalency. 

• Coulomb's law assumes that ions are completely spherical and seperate.
• Experiment values suggest there is a degree of electron sharing (covalency).
• This degree of covalency is due to polarization of ions.
• Polarization of an ion is the distortion of its electron cloud away from completely spherical.
• A cation will distort an ion, because the cation has polarizing power.
• The anion is polarizable. 

• The polarizing power of a cation depends on its charge density.
• A large nuclear charge acting over a small ionic radius will be more polarizing than a small charge acting over a large radius.
• The polarizability of an anion depends on its size only
• A large anion is very easily polarized, because its electron cloud is further from the nucleus and is held less tightly than on a smaller anion.  
• More covalency means a stronger bond. 

• A covalent bond is formed when a pair of electrons is shared between two atoms.
• This happens when the electron clouds of two atoms overlap and electron density is greatest beteen the nuclei.
• This region of high electron density attracts each nucleus and therefore keeps the atoms together.
• Covalent bonding is the strong electrostatic force of attraction between the nuclei of the bonded of the bonded atoms and the shared pair of electrons between them. 

• Dative covalent bonds are formed when both of the shared electrons come from just one atom.
• Aluminium chloride, AlCl3, will form dimers (combination of identical monomers) of Al2Cl6.
• The Al atom is AlCl3 is electron deficient (only has 6 electrons), but by forming dative covalent bonds the octet rule is fufilled.
• 1 Cl atom from each molecule gives an electron pair to the Al atom on the other molecule.
• Atoms can share more than one pair of electrons to form a double (oxygen) or triple (nitrogen) bond.  

• The physical properties of giant molecular structures (carbon and silicon) provide evidence for the strong electrostatic attraction in covalent bonding.
• They are very hard and have very high melting temperatures. A lot of energy is required to break the bonds that hold the atoms in place as a solid.
• Electron density maps show high electron density between atoms that are covalently bonded (more evidence). 

• In some molecules, not all the electrons in the outer shell may be involved in bonding.
• A non-bonding pair of electrons is called a lone-pair.
• Lone pairs affect the shape of a molecule.

• Metals consist of giant lattices of metal ions in a sea of delocalised electrons.
• The electrons vibrate around the metal ions (also vibrating) about a fixed point. The two hold eachother in place.
• It is the outer electrons of the metal that have become delocalised- they are no longer associated with one particular atom.
• Metallic bonding is the strong electrostatic attraction between metal ions and the surrounding sea of delocalised electrons.

• This model of metallic bonding can be used to explain the typical characteristics of metals.
• Electrical conductivity- the delocalised electrons are free to move in the same direction when an electric field is applied to the metal. The movement of charged particles is a current.
• Thermal conductivity- the delocalised electrons transmit kinetic energy (heat) through the metal, from a hot to a cold region, by colliding with eachother. 

• High melting temperatures- the metal ions are held together by their strong attraction to the surrounding delocalised electrons; it takes a lot of energy to break the metallic bonds.
• Malleability and ductility- Metals can be hammered into shape (malleable) or stretched into wire (ductile) because the layers of positive ions can be forced to slide across eachother while staying surrounded by the sea of delocalised electrons.