1. Formation of Compounds from atoms
Chemical Bonds
Formation of Compounds from atoms
Preparation for College Chemistry
Columbia University
Department of Chemistry

2. Trends in the Periodic Table
Lewis Structures
VSEPR Model

5. Atomic and Ionic Radii
Atomic Radii decrease going across a period from left to right, increases going down group

6. Ionization Energy X+(g) + e- X(g) ∆E = IE1 X2+(g) + e- ∆E = IE2 X+(g)
Minimum energy necessary to remove an electron from a neutral gaseous atom in its ground state (IE > 0, ground state stable system)
X+(g) + e-
X(g)
∆E = IE1
X2+(g) + e-
∆E = IE2
X+(g)

7. First Ionization Energies

8. Electron Affinity EA X(g) + e- X-(g) EA
Electron attachment energy. Energy released when an atom in ground state gains a single electron.
X(g) + e-
X-(g)
> 0 EA
< 0

9. F 2s2p5 F P 2s2p3 P Lewis Structures of Atoms
Gilbert Lewis. American Chemist F
2s2p5 F P
2s2p3 P H H H H + H F F F F F +

10. The Octect Rule F F H + H F H H O O H O + 2 H
In H2O and HF, as in most molecules and polyatomic ions, nonmetal atoms except H are surrounded by 8 electrons (an octet). Each atom has a noble gas electronic configuration (ns2p6 )

11. F- (2s2p6) F - Na + (2s2p6 ) Na + F - Na +
Ionic Bond. Electron Transfer F-
(2s2p6) F -
Na +
(2s2p6 ) Na + F - Na +

12. Ionic Bond. Electron TransferCl - Cl Mg
[Mg]2+ Cl - Cl O 2-
[Mg]2+ O Mg O Al
[Al]3+ 2-

13. Ionicity vs. Covalency

14. Covalent Bond. Sharing e-+ H

15. Collinear orbitals form s bond+
Two p AO = Two s MO
Only bonding MO shown

16. Coplanar orbitals form p bond+ O
Four p AO = Two s MO, and two p MO
Only bonding MO shown

17. Linus Pauling Electronegativity Cl H +d -d Dipole Cl H Cl H + Cl H +dF
Na+
Cl-
3.3

18. Lewis Structures of Compounds
Count valence electrons available.
number of valence electrons contributed by nonmetal atom is equal to the last digit of its group number in the periodic table.
(H = 1)
Add electrons to take into account negative charge.
Ex.
OCl– ion: 6 (O) + 7 (Cl) + 1 (charge) = 14 valence e–
CH3OH molecule: 4 (C) + 4(H) + 6 (O) = 14 valence e–

19. Lewis Structures of Compounds
Draw skeleton structure using single bonds
Note that carbon always forms four bonds.
Central atom is written first in formula.
Terminal atoms are most often H, O, or a halogen.
Ex. H
H — C — O— H
O — Cl -

20. Lewis Structures of Compounds
Subtract two electrons for each single bond
O-Cl– ion: 14 – 2 = 12 valence e- left
CH3 OH molecule: 14 – 10 = 4 valence e- left
Distribute remaining electrons to give each atom a noble gas structure (if possible). H
H — C — O— H
O — Cl -

21. Lewis Structures of Compounds
Too Few Electrons?
Form multiple bonds
Ex. What is the structure of the NO3 – ion?
valence e– = 5(N) + 18 (3O) + 1(charge) = 24 e–
Skeleton: O O N O O

22. Nitrate Ion (cont.) Resonance Structures
valence e– left = (3 single bonds) = 18 e–
Adding a double bond and rearranging:
Resonance Structures

23. Molecular Geometry VESPR principle:
electron pairs around a central atom tend to be oriented so as to be as far apart as possible.
BeF2 linear (2 pairs of e-)
BF3 trigonal planar (3 pairs of e-)
CF4 tetrahedral (4 pairs of e-)
PF5 triangular bipyramid (5 pairs of e-)
SF6 octahedral (6 pairs of e-)

24. Molecular Geometry

25. VSEPR model Molecule Lewis Str. Pairs of e- electron arrangem.
Molecular
Shape
H2S 4
tetrahedral
bent
H S H
CCl4 C Cl