Chemical Bonding
B - Polarity
D - Metallic bonding
C - Dot and Cross Diagrams
E - Giant Ionic and covalent structures.
A - Valance Shell Electron Pair Repulsion Theory
F - Simple molecular Structures
Simple molecular structures are substances composed of covalently bonded molecules which are held together by weak intermolecular forces.
A molecule is 2 or more atoms chemically bonded together, either from the same element or from multiple
The have low melting and boiling points because of the weak intermolecular forces. They require little energy to break the bonds between molecules, thus it will be liquid or gas at room temperature.
It states that we can predict the shape of a molecule using the fact that electrons repel all other electrons. So the electrons will take up positions further away from each other. These can be shared pair or lone pair.
Every molecule has Van der Waal forces. Molecules that contain Hydrogen and either Florine, Oxygen or Nitrogen can have Hydrogen bonds.
Dipole Dipole bonds are created depending on the shape of the molecule and if the bonds it has are polar or non polar.
In solutions, atoms may or may not be able to exist independently and they have to form molecules.
Ions can exist independently in solutions.
Generally 2 atoms from the right side of the periodic table bonding together would be covalent.
Non-metals that has a electronegativity difference of less than 0.4 normally bond with covalent bonds.
A giant covalent structure will have many covalent bonds that forms a regular pattern. In graphite, carbon atoms form hexagons which fit into layers that can slide across each other.
Examples -
Graphite
Diamond
Silicon dioxide
Positive metallic ions are in a general structure. They are a in a sea of delocalised electrons .
Giant covalent structures has many strong covalent bonds between the atoms. This means it requires lots of energy to break all those bonds therefore they have high boiling points.
Giant ionic structures are lattice shaped. The ions arrange themselves so that no cation will ever be next to each other as they repel each other due to them having the same charge. So cations will attract anions and so on.
Giant ionic structures can only conduct electricity when molten or aqueous because when in their solid state, they do not have free moving ions that can carry a charge.
Because the metallic ions are arranged in a general structure, they have layers of metallic ions on top of more layers which can easily slide across another.
They have a sea of delocalised electrons that can move freely and carry a charge. Therefore they can conduct electricity.
Oxygen has 2 lone pairs of electrons so they form a double bond to fill up the outer shell making it more stable. They also have enough space to form the bonds.
Normally non-metals and metals react together to form ionic bonds. They must have an electronegativity difference of above 2
Aluminium(II) ion = +2 charge
Aluminium(III) ion = +3 charge
Nitride ion = -3 charge
0 - 0.4 electronegativity difference means its covalent
0.4 - 2 electronegativity difference means its polar covalent.
2 - 4 electronegativity difference means its ionic.
Polar covalent bonds happens when 1 atom is more electronegative than the other so it tends to pull the charged electron closer to its nucleus making it delta negative leaving the other atom as delta positive.
In non-polar covalent bonds, the differences isn't high so they attract the electron at the equal strength so no one is slightly positive or negative.
Ionic bonds happens when 1 atoms is so much more electronegative that it comepletely removes the electron from the other atom making it self negative and the other positive.
Dative covalent bonds is when an atom bonds to a lone pair of electrons on another atom.
Since there are less electrons the the cation, it means that the positive nucleus can pull harder on the existing electrons making the radius smaller. This means it is more polarising which distorts the electron field around the anion. Pulling electrons closer to the cation
Hydrogen bonds are the strongest of the intermolecular forces. The more hydrogen bond there are the higher the melting and boiling points. Van der Waal are the weakest but the more van der waals forces there are the higher the boiling points but it wont affect it as much as hydrogen bonds.
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A molecule with six electron pairs could be SF6. The shape of the molecule is a Octehedron so its bond angles of 90 degrees from the central atom to all the other atoms. There are no lone pairs of electrons.
A molecule with 5 electron pairs could be PCl5. Its shape is Triagonal Biyramidal which has bond angles of 90 and 120 degrees. There are no lone pairs of electrons.
A molecule with 4 electron pairs could be CH4. Its shape is Tetrahedral with bond angles of 109.5 degrees. There are no lone pairs of electrons.
AX3 - 2 types of structures with 3 electron bond pairs.
Trigonal pyramidal has bond angles of 107 degrees. It has 1 lone pair so that cause repulsion forcing the 3 electron pairs to change angles.
Example - NH3
Trigonal planar 1 of 2 structures with 3 electron bond pairs. It is flat as there are no lone pairs of electrons. so the angles are 120 degrees.
Example - BF3
AX2 - 2 types of structures with 2 electron bond pairs
V-shaped or Bent has bond angles of 104.5 degrees. It has 2 lone pairs of electrons which repels the other atoms changing the bond angles.
Example - H20
Linear, Bond angle is 180 with 2 electron bond pairs. There are no lone pairs of electrons.
Example - CO2