Ionic Bonding - Evidence for ions and ionic compou
The evidence that ions exist:
- Physical Properties of ionic compounds:
- -high melting temperatures; showing strong forces of attraction between ions, as lots of energy needs to be put in to break the attraction between ions. Also evidence it exists as giant ionic lattice not discrete ionic compounds.
- -soluble in polar solvents, Soluble in water (Covalent molecules are insoluble in water but soluble in non-polar solvents). New bonds made between ions and water molecules
- -conduct electricity when molten or in aqueous solution. If ions free to move between + or - they conduct electricity.
- -electrolysis - the migration of ions in electrolysis e.g. of green aqueous copper (II) chromare (VI) attracts a yellow colour (Chromate (VI) ions) to the anode and a blue colour (copper(II) ions) to the cathode.
- Electron density maps of compounds produced from X-ray diffraction patterns show zero electron density between ions - meaning complete electron transfer.
- Metal + Non Metal
Metal loses outermost e-s to non-metal and formes +ve cation so its configuration becomes isoelectronic to a noble gas. As the non-metal gains an e- it forms a -ve anion and becomes isoelectronic to a noble gas.
When ions are formed they tend to have a full outer shell (i.e. 8 electrons) this is the octet rule.
Ionic Dot and Cross Diagrams
Only electrons in the outer shell are shown. they are drawn at the four points of the compass and paired up. the reactions of elements to form ionic compounds can be represented by dot and cross diagrams. the electrons from one reactant are shown with crosses and the other with dots. (The ion has sqaure brackets around it and a sign i.e. [Mg]2+)
The charge on each ion must be shown.
Ionic Bonding and lattices
An ionic bond is an omnidirectional short range electrostatic force between oppositely charged ions.
- Ions are charged spheres. they attract in all directions. So ionic compounds do not just exist as 2 or 3 ions attracted to each other, but as giant ionic lattices with each cation surrounded by anions and vice versa.
- In ionic compounds each ion is surrounded by ions of the opposite charge.
- Giant ionic lattices are also known as ionic crystals and are formed in the solid state.
Trends in ionic radii
The ionic radius is the radius of an ion in its crystal form.
- Cations are smaller than the original atom as they lose electrons. Usually a whole electron shell has been lost and the remaining electrons have been pulled in towards the nucleus more strongly.
- Anions are larger than the original atom since the atom gains electrons and there is more repulsion in the electron cloud.
Going down a group, the ions become larger- the number of shells increases.
Isoelectronic ions have different ionic radii:
- the additional electrons in anions make the ions larger as there is greater repulsion and all the electrons are less tightly bound than in the atom.
- the loss of electrons to form cations means the nucleus attracts those electrons that remain more strongly.
Lattice energies and Born-Haber Cycles
The formation of an ionic crystal from its elements is exothermic. (energy 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 can be broken into a number of stages each associated with a particular energy change:
- Atomization of the metal (turning to a gas)
- Ionization of the gaseous metal (loses electron so add in + e-)
- Atomization of the non-metal (turning to a gas)
- Electron Affinity of non-metal (gain the e- from before)
- Lattice enthalpy
- Forming the crystal from the gaseous ions.
Having measured all of these, the lattice energy can be calculared using an enthalpy level diagram known as a Born-Haber Cycle. similar to Hess' law.
A more exothermic lattice enthalpy means stronger ionic bonds.
Factors affecting the strength of the ionic bond (and how exothermic the lattice enthalpy is) are:
- Ionic radii - smaller the ionic radii are, the closer they can approach each other so the greater the attraction.
- Charge on ions - greater the +/- charge on the ion, the greater the attraction between the ions so the greater the ionic bond.
Lattice Enthalpy cannot be directly measured under standard conditions so we use the Born Haber Cycle. There are 2 ways to work out lattice enthalpy:
1) if the enthalpy formation= sum of all enthalpy changes in cycle,
the lattice enthalpy = enthalpy formation - sum of all other enthalpy changes
2) the enthalpys going up one way is equal to the lattice enthalpy plus the electron affinity.
Testing the Ionic Model
We can also calculate, using Coulumb's law (electrostatic attraction) a theoretical value for lattice enthalpy which assumes the ions are perfectly charged spheres.
Coulumb's law assumes complete electron transfer in ionic compounds, and the size of the ionic radii. Coulumb's Law calculates the force of attraction between ions as a function of their charges and the distance between them.
However, this model falls apart as cations can distort an anion by polarisation. Polarisation of an ion is the distortion of its electron cloud away from completely speherical.
Cations are good at polarising if they have a:
- high +ve charge
- small ionic radius
Anions easy to polarise have :
- high -ve charge
- large ionic radius.
Covalent Bonding - Formation of covalent bonds
A covalent bond is formed when a pair of electrons is shared between 2 atoms. (non-metal and non-metal)
This happens when 2 atoms approach each other and their electron clouds overlap and electron density is greatest between the nuclei. This region of high electron density attracts each nucleus and keeps the atoms together.
- Covalent bonding is a strong electrostatic attraction between the nuclei of the bonded atoms and the shared pair of electrons.
- The distance between the two nuclei is the bond length and is the seperation at which the energy of the system is at its lowest.
Dative Covalent Bonds are formed when both the shared electrons come from just one of the atoms. e.g. AlCl3 formed dimers (combination of two identical molecules) of Al2Cl6. Aluminium atom is electron-deficient but by forming dative covalent bonds the octet rule is fulfilled.
Atoms can share more than one pair of electrons to form double or triple covalent bonds.
Evidence for covalent bonds
The physical properties of giant atomic structures e.g. diamond, provide evidence for strong electrostatic attraction in covalent bonding.
Giant Atomic Structures are known as giant molecular structures.
- They are very hard
- have high melting temperatures
- covalent bonds are very strong
- require a lot of energy to break before atoms can move in a liquid.
Electron Density Maps show high electron density between atoms that are covalently bonded.
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 are shown on dot and cross diagrams as one atom only (not shared)
- Lone pairs affect the shape of molecules.
Metals consist of giant lattices of metal ions in a sea of delocalised electrons. The metal ions vibrate about fixed points in the solid lattice, being held in place by electrons around them.It is the outer electrons that become delocalised- they are no longer associated with one particular atom.Metallic bonding is the strong attraction between metal ions and the sea of delocalised electrons.
The typical characteristics of metals can be explained byusing this simple model of metallic bonding:
- 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- an electric current.
- Thermal conductivity: delocalised e-s transmit kinetic energy through the metal from hot region to a cooler one by colliding with each other.
- High melting temperatures : +ve ions are strongly held together by attraction of delocalised e-s; takes a lot of energy to break metallic bonds and allow particles to move around in liquid state.
- Malleability and ductility: metals can be hammered into shape (malleable) or stretched into wire (ductile) as layers of +ve ions can slide across each other while being surrounded by sea of e-s.