The Shape of Covalent Molecules

1. 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

2. 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.

3. Different ways to draw covalent bond

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

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

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

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

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

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

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

11. Trigonal pyramidal Tetrahedral V-shaped

12. Trigonal bipyramidal Unsymmetrical tetrahedral T-shaped Linear

13. Square pyramidal Square planar Octahedral

14. Predict the Shapes of molecules withoutmultiple 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

15. 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

16. 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

17. 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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