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Chemistry - Module 2 - Electronic structure - Part 2 - Coggle Diagram
Chemistry - Module 2 - Electronic structure - Part 2
molecular shapes and VSEPR principles
electron pair repulsion theory
electron pairs around central atom repel and arrange to minimise repulsion
lone pairs repel more strongly than bonding pairs, reducing bond angles (each pair by 2.5 degrees)
predicting shape
Steps: Identify central atom, count outer-shell electrons, count bonding electrons from molecular formula, divide total electrons by 2 to get electron pairs, deduce lone pairs and bond arrangement
shapes depend on numbers of bonding and lone pairs
Common shapes include: linear (180), trigonal planar (120), tetrahedral (109.5), trigonal pyramidal (107), bent (104.5)
Polarity and electronegativity
electronegativity concept
atoms differ in ability to attract bonding electrons; fluorine most electronegative, followed by O, N, Cl
Measured on Pauling scale
bond and molecular polarity
polar bond forms when atoms have different electronegativities generating a permanent dipole
molecular polarity depends on vector sum of bond dipoles; symmetrical molecules can be non-polar if dipoles cancel
intermolecular forces and their effects
types of intermolecular forces
induced dipole-dipole (London dispersion) exists in all atoms/molecules due to instantaneous dipoles
larger electron clouds yield stronger london forces
permanent dipole-dipole occur between polar molecules in addition to London forces
hydrogen bonding occurs when H is bonded to F, N or O and interacts with lone pairs on these atoms in neighbouring molecules
consequences for properties
hydrogen bonding raises boiling/freezing points and solubility in water
main determinant of boiling point for non-hydrogen-bonding substances is strength of induced dipole forces
simple covalent substances have low melting/boiling points and generally do not conduct electricity
polar molecules and those able to hydrogen-bond tend to be soluble in water; non-polar molecules often insoluble in water