Presentation on theme: "VSPER Molecular shapes Valence Shell Electron Repulsion Theory."— Presentation transcript:
VSPER Molecular shapes Valence Shell Electron Repulsion Theory
VSEPR Theory Covalent bonds consist of shared pairs of electrons These electrons must exist between the atoms sharing them Electron pairs are negatively charged and must repel each other These repulsions force the molecule to exist in a specific arrangement of atoms or shape
So… In Valence Shell Electron Pair Repulsion (VSEPR) theory, pairs of electrons that surround the central atom of a molecule or ion are arranged as far apart as possible to minimise electron-electron repulsion.
What does this mean? Essentially… Electrons have a mutually repulsive force- they repel each other due to their negative charge The shape of the charge cloud affects how much it repels other charge clouds Where electrons are involved in bonding they have a different repulsive effect than lone pairs- those not used in bonding Produces predictable shapes
So… Greatest angles between lone pairs- bonding pairs produce somewhat lower angles Presence or absence of lone pairs will affect overall shape of the molecule
This Idea Is Easily Seen With Balloons Tied Together The Shapes of 2, 3, and 4 Repulsions
In summary- Lone pair/lone pair bonds largest angle Lone pair/bonding pair next lowest Bonding pair/bonding pair lowest Simples…..? Lets start with bonded pairs
Linear Atoms in a straight line- examples Two bonding pairs- only- angle 180 o Beryllium Chloride (BeCl 2 ) Carbon Dioxide (CO 2 ) What's the difference?
Single bonds and double/triple bonds do not affect the overall shape of the molecule! Just as well as this is complex enough as it stands!
Trigonal planar Triangular/flat- three bonding pairs- angles 120 o Example- Boron Trichloride (BCl 3 )
Ethene (C 2 H 4 )is similar, but consists of two sets of three bonds- still trigonal planar
Tetrahedral Four bonding pairs- angles o Example methane (CH 4 )
Trigonal Bipyramidal Five bonding pairs- 90 o and 120 o Example- phosphorous fluoride (PF 3 )
Octahedral Six bonding pairs- all 90 o Example- Sulphur Hexafluoride (SF 6 )
What about lone pairs? Alters the angle- effectively 'bends' the molecule- non linear Example- planar- water (H 2 O) Bonding angle o - due to two lone pairs
Pyramidal Again, three bonding pairs Example- ammonia (NH 3 ) One lone pair- bonding angle 107 o
CommonBonding Geometries Common Bonding Geometries Atoms (B) which are bonded to a central atom (A): AB n AB 2 AB 3 AB 4
tetrahedron o trigonal planar 120 o linear180 o Each shape is associated with particular angles