1. This is one of the chapters you must read…. chapter 6…bonding

2. Student will learn: 1. three types of bonding ionic, covalent, metallic 2. two categories of bonding polar, non-polar 3. how to draw Lewis structure 4. how to calculate electronegative 5. bond characteristics 6. VSPER theory

3. Chemical Bonds ch6 p.161 ?What holds chemicals together? Chemical Bonds: electrical attraction between +nuclei and –valence electron of different atoms

4. By bonding together the atoms are more stable, and have a lower level of energy arrangement

5. 3 types of Bonding: l. Ionic Bonding: lose or gain –e metal + Non-metal = ionic bonding = makes ions 2. Covalent Bonding : share –e Non-metal + non-metal=covalent bonds = makes molecules 3. Metallic Bonding: -e flow free in a sea of –e Transition metals Pick the bonding NaCl, CH 4, HCl, K 2 S, FeSO 4, LiF, H 2 0, Cu, Zn, Mg(OH) 2

6. 2 catagories for the bonding polar non-polar Unequal equal attraction for electrons balanced attraction Ionic bonding (Metals+nonMetals) is always polar Covalent (nonMetals +nonMetals) maybe either polar/nonpolar 2 ways to figure out Draw Lewis structure calculate and use chart

7. Lewis Structure for iodomethane CH 3 I 1. Lewis Dot for each element CHICHI 2. Arrange to form skeleton If a carbon then always in middle If no carbon then least electronegative atom in middle Hydrogen never in middle See that it is lopsided…… polar covalent molecule How about individual bonds?

8. Lewis for ammonia NH 3 1.Draw Lewis dot for each element NHNH 2. Do skeleton : hydrogen never in middle Does it look lopsided…. polar covalent molecule

9. Lewis for formaldehyde CH 2 0 l.Lewis dot for eachCHO 2. Do skeleton: carbon always in middle Notice left out –e….move to make a double bond Notice lopsided: polar……..covalent Single bonds Double bonds Triple bonds

10. Lewis struture for : CCl 4 l.Draw Lewis dot for each: C CL 2.Draw skeleton: carbon always in middle Does it look lopsided?........ No….. non-polar covalent molecule What about each bond?

13. Use table on p.151 and chart on page 162

14. Calculating Polarity of Bonds: using differences in electronegativity Remember: electronegativity = ability to gain electrons Bonding is rarely purely ionic or covalent………most of time somewhere in between Use table on p.151 and chart on page 162 (overhead 31) Subtract the two electronegativity numbers then ? is it less than 1.7 = polar covalent? Calculate bond type and polarity KCl, MgCl 2,H 2,H 2 S Cs 2 S, SCL 2,

16. Comparing Characteristics Ionic bonds (vs) Covalent bonds Metals + nonmetals Gain or lose electrons so…. + or – ends ……very polar Will form a crystalline lattice, look on page 177 (model) Stronger bonds Most are solids Higher melting point Higher boiling point Many dissolve in water, +ion and -ion break apart in water so will conduct electricity in water. Some do not dissolve because the pull between the charges are greater than the attraction of H 2 O molecule Hard but brittle----why? A shift of one row of ions causes a large build up of repulsive forces. And --do not like-- so if one layer moves that forces the other layers to move so they are brittle. Non-metals + Non-metals Share electrons Exist as individual molecules Weaker bonds Most are gases, some liquids Very low melting point Very low boiling point Will evaporate at room temperature Overhead 70

17. Ionic compounds form Crystalline lattice

18. How ionic compounds dissolve

19. Why Ionic Compounds are brittle

20. Comparing Characteristics Ionic bonds (vs) Covalent bonds Metals + nonmetals Gain or lose electrons so…. + or – ends ……very polar Will form a crystalline lattice, look on page 177 (model) Stronger bonds Most are solids Higher melting point Higher boiling point Many dissolve in water, +ion and -ion break apart in water so will conduct electricity in water. Some do not dissolve because the pull between the charges are greater than the attraction of H 2 O molecule Hard but brittle----why? A shift of one row of ions causes a large build up of repulsive forces. And --do not like-- so if one layer moves that forces the other layers to move so they are brittle. Non-metals + Non-metals Share electrons Exist as individual molecules Weaker bonds Most are gases, some liquids Very low melting point Very low boiling point Will evaporate at room temperature Overhead 70

21. Why most covalents are liquids or gases and evaporate easy.

25. Metallic Bonding p.181 Transitional Metals: vacant outer p orbitals because filling up d orbitals first. 4s 2, 3d 10, 4p … they overlap This overlapping lets –e roam freely about the metal network of empty atomic orbitals. These mobile –e form a sea of electrons which are packed in a lattice form. Overhead 68

26. Characteristics of Metallic bonding“Cu,Au,Ag, Fe” l.Conduct electricity Conduct heat ::::::due to the “sea of electrons” ability to move freely 2.Reflect light, Shiny, Polish :::::: Contain many orbitals (d 10 ) separated by extremely small energy differences, metals can absorb a wide range of light frequencies. This absorption of light energy accounts for the ability to reflect light and be shiny. …..p. 181 3.Malleable: hammer into a thin sheet. ::::::possible because the metallic bonding is same in all directions throughout the solid because of “sea of electrons”

27. Ductile: ability to be drawn into a thin wire. ::::::Because the metallic bonding is same in all directions throughout the solid because of “sea of electrons”

28. VSEPR THEORY V alence S hell E lectron P air R epulsion Theory : replusion between Valence Shell Electrons Pairs surrounding an atom causes these sets to be oriented as “far apart as possible”. “AS FAR APART AS POSSIBLE” Lewis dot, VSEPR TO PREDICT GEOMETRY OF MOLECULE,

34. Intermolecular force : the attraction between molecules 3 types: dipole-dipole hydrogen bonding London dispersion forces Dipole-dipole: strongest intermolecular force created when by equal but opposite charges of the molecule come within close distance of each other.

35. Hydrogen bonding: hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a near-by molecule. “this is why water expands when frozen”

36. London Dispersion force : results from the constant motion of electrons and the creation of instantaneous dipoles. Because London dispersion force depends on the motion of electrons, the strength increases with increasing atomic masses.