The Shape of Covalent Molecules 1. VSEPR Theory 2. Different ways to draw covalent bond 3. Different shapes of molecules 4. Shapes of molecules with lone pair of electrons in the central atom 5. Predict the Shapes of molecules without multiple bonds 6. Shapes of Molecules with Multiple Bonds

VSEPR Theory Q. What is the charge of an electron carries? A. -ve Q. What will happen if two bond pairs of electrons are placed around a central atom? A. They will stay as far apart as possible to minimize the electronic repulsion. This is the key concept of VSEPR Theory. If you know the no. of electron pairs around the central atom, you can predict the shape of the molecule.

Different ways to draw covalent bond

Different shapes of molecules If there are 2 electron pairs, the shape of the molecule is ________ linear.

If there are 3 electron pairs, the shape of the molecule is ____________ trigonal planar

If there are 4 electron pairs, the shape of the molecule is ____________ tetrahedral

If there are 5 electron pairs, the shape of the molecule is ___________________ trigonal bipyramidal

If there are 6 electron pairs, the shape of the molecule is ____________ octahedral.

Shapes of molecules with lone pair of electrons in the central atom

Change of molecular shape 3 valence pairs of electrons Trigonal planar V-shaped

Trigonal pyramidal Tetrahedral V-shaped

Trigonal bipyramidal Unsymmetrical tetrahedral T-shaped Linear

Square pyramidal Square planar Octahedral

Predict the Shapes of molecules without multiple bonds 1. Count the no. of outermost e- in the central atom. 2. Add one if the particle has one negative charge or subtract one if it has one positive charge. 3. Add one for each bonding atom. 4. no. of pairs of e- = total / 2 5. no. of lone pair = no. of pairs of e- - no. of bonded atoms

e.g. 1, PCl4+ 1. Count the no. of outermost e- in the central atom. 5 ( P is group 5) 2. Add one if the particle has one negative charge or subtract one if it has one positive charge. 5 - 1 = 4 (the particle has 1 +ve charge) 3. Add one for each bonding atom. 4 + 4 = 8 (there are 4 Cl atoms) 4. no. of pairs of e- = total / 2 8 / 2 = 4 5. no. of lone pair = no. of pairs of e- - no. of bonded atoms 4 - 4 = 0  tetrahedral

e.g. 2, XeF2 1. Count the no. of outermost e- in the central atom. 8 ( Xe is group 0) 2. Add one if the particle has one negative charge or subtract one if it has one positive charge. 8 (the particle does not have charge) 3. Add one for each bonding atom. 8 + 2 = 10 (there are 2 F atoms) 4. no. of pairs of e- = total / 2 10 / 2 = 5 5. no. of lone pair = no. of pairs of e- - no. of bonded atoms 5 - 2 = 3  linear

Shapes of Molecules with multiple bonds  Both e- pairs must stay together in a double bond or triple bond.  We can treat them as single bonds. e.g. CO2 There are 2 double bonds and no lone pair.  linear

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