Bonding and physical properties
The four types of crystal structure are ionic, metallic, giant covalent,and molecular. Examples of these structure are diamond, graphite, ice, iodine, magnesium and sodium chloride.
Ionic structures have high melting points due to the strong electrostatic attrection between them. They also conduct electricity if molten or dissolved, as the ions are free to move and carry the current.
Metallic structures have fairly high melting points due to the attraction between the cations and delocalised electrons, however it is not as strong as an ionic bond. They also conduct electricity because the delocalised electrons are free to move and carry the current.
Molecular covalent structures have low melting points due to weak inter molecular forces that are broken as opposed to the actual covalent bond and do not conduct electricity due to the lack of ions.
Giant covalent structures have high melting points due to the strong covalent bonds that have to be broken, they also don't conduct electricity due to the lack of ions.
Giant covalent layers can slide over each other and can conduct electricity due to the presence of delocalised electrons.
When a substance moves from a solid through to a gas it gradually gains kinetic energy and begins to vibrate more and more.
Electronegativity is the power of an atom to attract the pair
of electrons in a covalent bond
We can predict whether the bond will be covalent, polar covalent or ionic using the electronegativity values. :
Some atoms attract electrons more strongly than others hence the electrons do not lie in the middle of the bond. where this occurs we say the bonds are polar as the end where the electrons are closer to has a partial negative charge (delta negative) so the other end has a partial positive charge (delta positive).
Electronegativity depends on the atomic radius and its nuclear charge.
A temporary dipole can occur in any molecule when, by chance, all of the electrons are on one side of a molecule. It is temporary because it lasts for a very short time as the electrons are constantly moving. Temporary dipoles are constantly appearing and disappearing.
A permanent dipole occurs when one side of a molecule is constantly delta positive and negative, however if a molecule with polar bonds completely surrounds the central atom the molecule is not a permanent dipole.
Shapes of simple molecules and ions
Bonding pairs and lone (non-bonding) pairs of electrons as
charge clouds that repel each other.
Pairs of electrons in the outer shell of atoms arrange
themselves as far apart as possible to minimise repulsion
Lone pair–lone pair repulsion is greater than lone pair–bond
pair repulsion, which is greater than bond pair–bond pair
If there is a lone pair the bond angle is reduced by 2.5 degrees.
VESPR Theory: 1. Decide which central atom. 2. Using he periodic table, write down the number of outer shell e- of the central atom. 3. Add e- if the particle is an anion or take away e- if the particle is a cation. 4. Add e- for each atom covanlently bonded to the central atom. 5. You now have the number of e-, to get the number of e- pairs you must divide by 2. 6. Now you can predict the shape and bond angles of the molecule.
Forces between molecules
in a permanent dipole-dipole force a weak intermolecular force is created between more electronegative atoms and less electronegative atoms due to the small electrostatic attraction. Due to the fact that it is a weak bond, it can easily be broken.
An induced dipole appears as a result of of attraction between two molecules at the end of two contrasting dipoles. This is known as Van der Waal's forces. They are very weak and are present between all molecules and are the reason that compounds can be liquefied and solidified.
Hydrogen bonding is a bond between a N,O or F, and a hydrogen atom (the atoms with the highest and lowest electronegativities). Although it is fundamentally a dipole-dipole bond, it is much stronger due to the large differences between the electronegativities. This is the strongest intermolecular force and therefore massively increases melting and boiling points.
Ice has a low density due to the presence of more hydrogen bonding than water as it wishes to maximise this type of bond. This creates very open hexagonal structures which accounts for the low density.
Hydrogen bonding also causes anomalous melting and boiling points due to the fact that the strength of the intermolecular force is so strong.
Ionic bonding involves electrostatic attraction between
oppositely charged ions in a lattice.
Sulphate: [S04]2-, Hydroxide: [OH]-, Nitrate: [NO3]-, Carbonate: [CO3]2-, Ammonium: [NH4]+
For non metals, their ions are always negatively charged and you can use its position on the periodic table to work the charge out using the distance it is away from the noble gas, e.g. Group 7 has 1- charge.
For metals, their overall positive charge corresponds with their group number.
Nature of covalent and dative covalent bonds
A single covalent bond shares a pair of electrons, and multiple bonds contain multiple pairs of shared electrons.
Methane: CH4, Ethene: C2H4, Ammonia: NH3
Covalent bonding is the electrostatic attraction between shared pairs of electrons and the nucleus.
Dative bonding is a type of covalent bonding, where both electrons are donated from one atom. This atom therefore must have a lone pair of electrons.
Metallic bonding involves attraction between delocalised
electrons and positive ions arranged in a lattice.