Chemical Bonding: Ionic, Covalent, Polar Covalent, and Metallic Bonds

1. Lesson 7 Chemical Bonding Anything in black letters = write it in your notes (‘knowts’)

2. Review… Elements in the same column have similar properties because… they have the same number of valence electrons electrons in the highest energy level

3. variable

4. Section 1 – Ionic & Covalent Bonding Lewis Symbol – show valence electrons as dots around chemical symbol C Na O Br

6. Octet Rule – Atoms will gain or lose electrons to have 8 valence electrons. metals lose e-; to form ca+ions nonmetals tend to gain e-; to form anions

7. Normal Ion Charges

8. Ionic bonds are formed by electron transfer. (between metal & nonmetal) Covalent bonds are formed by electron sharing. (between 2 nonmetals) The more electronegative element acquires most of the e- charge

9. Ionic compounds consist of cations and anions arranged in repeating patterns; NOT as molecules A formula unit is the lowest ratio of ions in an ionic compound.

10. Covalent Bond single bond double bond H2 O2 Each atom shares e- with another to complete the octet.

11. COMPOUNDS IONIC COVALENT transfer of e- (ions!) sharing of e- (COvalent) formula units molecules metal & nonmetal 2 nonmetals low melting points high melting points (most are gases or liquids at room temp) (solid crystals at room temp) conduct e- when molten or dissolved (aqueous)

12. FORMULA UNITS vs. MOLECULES Array of sodium ions and chloride ions Collection of water molecules Formula unit of sodium chloride Molecule of water NaCl Chemical formula Chemical formula H2O

13. Section 2 – Polar Covalent Bonding Covalent bonding involves a sharing of electrons between atoms but this sharing may not be equal… First, a review of electronegativity….

14. Electronegativity Review… B<H<C Noble gases do not have e-neg values

15. A nonpolar covalent bond occurs between 2 identical atoms. Since each atom in a nonpolar bond has the same electronegativity, the electrons are shared equally. 7 nonpolar (diatomic) molecules (H2, N2, O2, F2, Cl2, Br2, I2)

16. Polar Covalent Bond (polar bond) – covalent bond in which the electrons are shared unequally. The more electronegative atom attracts electrons more and has a slightly negative charge (δ-) δ+ δ- partial positive charge partial negative charge

17. A bond or molecule with +/- charged ends is also called a dipole. “die-pull” δ- δ+ H Cl partial negative charge partial positive charge H Cl

18. H—Cl The polar nature of a bond can also be shown by an arrow pointing to the more electronegative atom. • Identify the bonds between these elements as ionic, polar covalent or nonpolar covalent. • H – Br b) K – Cl • C – O d) Li – O • Cl – F f) Br – Br • H – O h) H – Br • Place a δ- symbol above the more electronegative atom in the bond.

19. If the polar bonds in a molecule cancel out, the molecule is nonpolar. When the polar bonds do NOT cancel out, the molecule is polar.

20. We will not cover molecular geometries and shapes except for that of water… H2O unshared pairs of e- (lone or nonbonding pair) A molecule of water has a bent shape due to the space needed by the lone pairs of e-

21. Section 3 – Bonding Between Molecules So far, we have talked about bonding between atoms (ionic, covalent) Now, we will talk about bonding between molecules The bonds between separate molecules are much weaker than ionic or covalent bonds, but without these forces there would be no liquids or solids.

22. Forces between separate molecules are called Intermolecular Forces (or Van der Waals Forces) 2 types of intermolecular forces that we will discuss, Hydrogen Bonding Dispersion Forces (London Forces)

23. Hydrogen Bond – intermolecular force between H in one molecule and an electronegative atom (N, O or F) in another nearby molecule. Hydrogen Bonds in H2O

24. hydrogen bond Opposites partial poles attract

26. Dispersion Forces – temporary attractive forces between molecules due to electron dispersion (motion) Without dispersion forces, nonpolar molecules could never be liquids or solids. F2 melts at 53 K, Boils at 85 K

27. Molecule Melt Pt Boil Pt F2 53 K 85 K Cl2171.6 K 239.1 K Br2 265.8 K 332.0 K I2 386.8 K 457.4 K Dispersion Force

28. Cohesion – attraction to same substance. Adhesion – attraction to different substance

29. A tale of two molecules… butane (C4H10) 58 amu acetone (C3H6O) 58 amu b.p. = -0.5°C gas @ room temp b.p. = 56.1°C liquid @ room temp 1. Why does acetone have a higher boiling point? 2. What would cause butane molecules to stick to each other to become a liquid

30. Why does acetone have a higher boiling point? the opposite partial charges on each molecule hold the molecules together. etc… etc… δ- δ+ δ+ etc… butane is a nonpolar molecule and does not have this type of intermolecular attraction etc…

31. 2. What would cause butane molecules to stick to each other to become a liquid? butane, like all molecules, has electrons that are randomly moving. This produces temporary poles within the molecule.

32. Section 4 – Metallic Bonding The valence electrons in metals are loosely held and are free to move. The properties of metals can be explained by the ‘sea of electrons’ model.

33. Alloy – mixture of metals. Steel – Stainless Steel – Bronze – Solder – Brass – Sterling Silver – Amalgam – Nichrome – Alloy – mixture of metals. Steel – Fe & C Stainless Steel – steel w/ Cr, Ni, or Mn Bronze – Cu & Sn Solder – Sn & Pb Brass – Cu & Zn Sterling Silver – Ag & usually Cu Amalgam – Hg w/ other metals (Ag, Sn, Cu) Nichrome – Ni & Cr