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Topic 4 Bonding - Coggle Diagram
Topic 4 Bonding
Covalent bonding
can be represented using Lewis structure
Show all electrons
Formal charges can be used to distinguish between different Lewis structures for a molecule/ion
Formal charge (FC)=#of valence-0.5#of bonding-#of nonbonding
The preferred Lewis structure is the one in which the formal charges are closest to zero.
Resonance structures differ in the position of double bonds
Covalent bonding
Bond strength: Single < Double < Triple
Giant covalent compounds
Have high melting points because covalent bonds are broken when melting.
Examples
Silicon dioxide
Diamond
Allotropes of carbon
C60 is also an allotrope of carbon, but is covalent molecular.
Graphene
Graphite
Layered structure
Covalent molecular compounds
Non-polar molecules
Polar molecules
Caused by
Difference in electronegativity: Electrons lie more towards one atom than another
Shape of the molecule: if the dipole moments of individual bonds cancel, the molecule is not polar.
Only intermolecular forces are overcome when melting
types of intermolecular force
Diole-dipole forces
For the same Mr, a polar molecule will have a higher melting point than a non-polar molecule.
London forces, caused by instantaneous and induced dipoles
Stronger for higher Mr.
Hydrogen bonds
when H is bonded to N, O or F
Strongest intermolecular force
Use VSEPR to predict the shapes of molecules
Pairs of valence electrons (electron domains) repel each other and will take up positions to minimise these repulsions.
Standard shapes and bond angles
electron domains
3 Trigonal planar 120
4 tetrahedral 109.5
2 Linear 180
Electron domains
5 trigonal bipyramidal 90/120
6 octahedral 90
Lone pairs repel more than bonding pairs
Example shapes and bond angles
NH3 trigonal pyramidal 107
H2O bent 104.5
pairs of electrons are shared by different atoms
In a coordinate covalent bond both electrons come from the same atom
achieved by overlap of orbitals
types
σ bond – head-on overlap of 2 orbitals; electron density lies along the internuclear axis
π bond – side-on overlap of 2 p orbitals; electron density lies above and below the internuclear axis
bonds
A double bond consists of 1 σ and 1 π bond.
A triple bond consists of 1 σ and 2 π bonds.
Hybridisation
3 electron domains around central atom - basic shape is trigonal planar - sp2 hybridisation
4 electron domains around central atom - basic shape is tetrahedral - sp3 hybridisation
2 electron domains around central atom - basic shape is linear - sp hybridisation
Occurs when there is a small difference in electronegativity
In delocalization, a pair of electrons is shared between three or more atoms
The oxygen bond is weaker in O3 than in O2, so O3 absorbs longer wavelength UV radiation
The destruction of ozone can be catalysed by Cl radicals or NOx
The bond length is intermediate between that of a single and that of a double bond
electrostatic attraction between nuclei and shared pair of electrons
metallic bonding
Alloys are homogeneous mixtures of two or more metals or a metal with a non-metal.
Alloys have enhanced properties
Lattice of positive ions surrounded by delocalized electrons
Electrostatic attraction between positive ions and delocalized electrons
ionic bonding
giant lattice structure made up of positive metal ions and negative non-metal ions
electrostatic attraction between oppositely charged ions
characteristics
insoluble in non-polar solvents
conduct electricity when molten or dissolved in water
often soluble in water
high melting point
complete transfer of electrons — large difference in electronegativity