1. Chapter 6 Ionic Bonds and Some Main-Group Chemistry

2. Ionic Bonding Occurs in ionic compound Results from transferring electron Created a strong attraction among the closely pack compound

3. Ions and their configuration Na Na + Mg Mg 2+ Cl Cl - O O 2-

4. Ions and Their Electron Configurations - 2 e - - 3 e - Fe: [Ar] 4s 2 3d 6 Fe 2+ : Fe 3+ : [Ar] 3d 6 [Ar] 3d 5 IonsAtoms Mn: Mn 2+

5. Ionic Radii

7. Ionization Energy Ionization Energy ( E i ) : The amount of energy necessary to remove the highest-energy electron from an isolated neutral atom in the gaseous state.

8. Ionization Energy Increasing Ei Decreasing Ei

9. Ionization Energy Boron has a lower E i due to a smaller Z eff (shielding by the 2 s electrons)

10. Ionization Energy Oxygen has a lower E i since the first electron is removed from a filled orbital

11. Higher Ionization Energies M 3+ + e - M 2+ + energy M 1+ + e - M + energy M 2+ + e - M 1+ + energy

12. Electron Affinity Electron Affinity ( E ea ) : The energy released when a neutral atom gains an electron to form an anion.

13. Ionic Bonds and the Formation of Ionic Solids Na 1+ Na 1 s 2 2 s 2 2 p 6 Cl+ 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3s13s1 3p53p5 1 s 2 2 s 2 2 p 6 1 s 2 2 s 2 2 p 6 3 s 2 Cl 1-

15. The Octet Rule Octet Rule : Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons.

16. The Octet Rule

17. Chapter 7 Covalent Bond and Molecular Structure

18. Why do atoms share electrons? Bonds: a force that holds groups of two or more atoms together and makes them function as a unit Required 2 e- to make a bond Single bond, Double bond and triple bond Bond energy: amount of energy required to form or to break the bond Octet Rule : Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons.

19. The Covalent Bond Covalent Bond : A bond that results from the sharing of electrons between atoms.

21. Polar Covalent Bonds: Electronegativity Electronegativity : The ability of an atom in a molecule to attract the shared electrons in a covalent bond.

22. Polar Covalent Bonds: Electronegativity NaCl HCl Cl 2

23. Polar Covalent Bonds: Electronegativity Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond Electrons are not completely transferred More electronegative atom: δ-. (δ represents the partial negative charge formed) Less electronegative atom: δ+

24. Relationship Between Electronegativity and Bond Type Predicting bond polarity ◦ Atoms with similar electronegativity ( Δ EN <0.4) – form nonpolar bond ◦ Atoms whose electronegativity differ by more than two ( Δ EN > 2) – form ionic bonds ◦ Atoms whose electronegativity differ by less than two ( Δ EN < 2) – form polar covalent bonds

25. Examples For each of the following pairs of bonds, choose the bond that will be more polar a.H-P, H-Cb.N-O, S-O

26. Lewis Structures represents how an atom’s valence electrons are distributed in a molecule Show the bonding involves (the maximum bonds can be made) Try to achieve the noble gas configuration Bonding pair: two of which are shared with other atoms Lone pair or nonbonding pair: those that are not used for bonding

27. Examples Draw Dot Lewis structure for the following atoms: ◦ Na ◦ Mg ◦ C ◦ S ◦ Co

28. Rules for multiple atoms Duet Rule: sharing of 2 electrons ◦ E.g H 2  H : H Octet Rule: sharing of 8 electrons ◦ Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule ◦ E.g F 2, O 2

29. Electron-Dot Structures

31. Rules for Wring Dot Lewis structure Step 1: Calculate the total number of valence electrons of all atoms in the molecule Step 2: Create a skeletal structure using the following rules: ◦ Hydrogen atoms (if present) are always on the “outside” of the structure. They form only one bond ◦ The central atom is usually least electronegative. It is also often unique (i.e,. the only one atom of the element in the molecule). Remember, there might be no “central” atom. ◦ Connect bonded atoms by line (2-electron, covalent bonds

32. Rules Step 3: Place lone pairs around outer atoms (except hydrogen) so that each atom has an octet Step 4: Calculate the number of electrons you haven’t used. ◦ Subtract the number of electrons used so far, including electrons in lone pair and bonding pairs, from the total in Step 1. Assign any remaining electrons to the central atom as lone pair Step 5: If the central atom is B (boron) or Be (beryllium), skip this step If the central atom has an octet after step 4, skip this step ◦ If the central atom has only 6 electrons, move a lone pair from an outer atom to form a double bond between outer atom and the central atom ◦ If the central atom has only 4 electrons, do Step 5a to two different outer atoms (i.e, form two double bonds) or twice to one outer atom (i.e., form one triple bond)

33. Electron-Dot Structures of Polyatomic Molecules Total valence electrons Step 4: Dot Lewis structure Step 1: Step 2: Draw an electron-dot structure for CF 4. Draw its skeletal

34. Examples Give the Lewis structure for the following ◦ H 2 O ◦ CO 2 ◦ BH 3 ◦ NH 4 + ◦ NO 3 -,