1. Reactivity of Elements How many valence electrons do the noble gases have? 8 electrons = octet Compounds form when two or more elements come together, ideally so each can have a full octet Ionic compound: electron(s) transferred from one atom to another Covalent compound: electron(s) shared between two atoms Electron interactions generate chemical bonds, which hold together atoms in a compound 1

2. Metals –often lose e- –CATION Nonmetals –often gain e- –ANION –“-ide” Ions

3. Practice Predict the ion each of the following elements becomes: O I S Mg Cs Al N

4. Ionic Compounds 4 Composed of ions Transfer of electrons Neutral (no net charge) Charges must balance Subscripts may be used Binary ionic compounds –Two elements –Usually metal + nonmetal –Solids at room temp

5. Ionic Compound: Sodium Cloride 5

6. 6

7. Naming Binary Ionic Compounds Cation + Anion name of metal +root of nonmetal-ide Example: Strontium and Iodine strontium iodide Example: Nitrogen and Cesium cesium nitride

8. How would you name the following compound? –Calcium and Fluorine Calcium Fluoride –Lithium and Bromine Lithium Bromide 8 Naming Binary Ionic Compounds

9. Ionic Compounds Net ionic charge = 0 Examples: –Sodium chloride –Calcium bromide –Magnesium sulfide –Lithium oxide

10. Metals that form more than one ion Usually in B groups (transition metals) Ex. Iron Fe 2+ Fe 3+ FeCl 2 FeCl 3 Iron (II) chlorideIron (III) chloride chromiumcoppergoldironleadtin Cr 2+ Cu + Au + Fe 2+ Pb 2+ Sn 2+ Cr 3+ Cu 2+ Au 3+ Fe 3+ Pb 4+ Sn 4+

11. Polyatomic Ions Polyatomic ions: ions made up of more than one atom

12. Polyatomic Ions Nitrate: NO 3 charge = -1 Sulfate: SO 4 charge = -2 Examples: Potassium nitrate Calcium nitrate Potassium sulfate Calcium sulfate

13. 13

14. Covalent Compounds Atoms share electrons to form molecules 14 HH+HHHHH2H2 FF+FF FF F2F2 Lone pair electrons Covalent bonds usually between nonmetal + nonmetal Bonding electrons

15. CH+CHCH + H CH H CH HCHH H H Carbon atom: 4 valence electrons Wants to form 4 bonds C+4H

16. NH+ + H Nitrogen atom: 5 valence electrons 3 unpaired electrons Wants to form 3 bonds +3HNNHH H NHNH NH H NH H

17. H+ + H Oxygen atom: 6 valence electrons 2 unpaired electrons Wants to form 2 bonds +2H OOHOH OH H OH H O OH H

18. Multiple Bonds Unpaired electrons left! 18 NN+NNNN C++OOCOO COOCOO

19. Resonance Structures When two or more electron-dot formulas can be drawn, these are called resonance structures Example: O 3 19 OOO 18 electrons total OOOOOO OOO Each bond is a “one and a half” bond

20. Covalent Compounds Combination of two different elements (usually nonmetal + nonmetal) Rules for Naming -Element with lower group number is first word; higher group number element is second word. -Same group? element with higher period number first -Second element: “root”-ide -Add Greek numerical prefixes

21. Naming Covalent Compounds Example: CS 2 1.Which element is in lower group? 2.(Not in same group so doesn’t apply) 3.Second element is root-ide 4.Add prefix2 sulfurs Carbon Disulfide

22. Naming Covalent Compounds Example: N 2 O 1.Which element is in lower group? 2.(Not in same group so doesn’t apply) 3.Second element is root-ide 4.Add prefix2 nitrogens Dinitrogen oxide

23. 23

24. Electronegativity Electronegativity: ability of an atom to attract electrons 24 (EN)

25. Types of Bonding / Variations Ionic bond: large difference in electronegativity; polar Covalent bond: small difference in electronegativity Some difference in electronegativity: polar covalent No difference in electronegativity: nonpolar Polar bonds creates a dipole 25

26. Look up electronegativities (Fig. 4.5) Find the Difference between the two 26

27. Shapes of Molecules VSEPR: Valence Shell Electron Pair Repulsion Minimize electron repulsion Electron groups pushed as far apart as possible Electron groups: Bonding electrons (bonds) –Note: single, double, triple bonds all count as ONE electron group Lone pairs of electrons 27

28. Basic Molecular Shapes 2 electron groupslinear O=C=O 3 electron groups trigonal planar bent S OO BF F F

29. Basic Molecular Shapes: 4 electron groups tetrahedral bent O HH pyramidal C H H H H N H H H

30. Polarity of Molecules Polar bonds: differences in negativities of bonded atoms 30 H-HF-F Non-polar (atoms equal EN) Polar (F more EN than H; lone pairs around F) H-F O HH Add polar bond together to see if molecule is polar overall (head to tail) PolarO partially negative H partially positive

31. Polarity of Molecules 31 Non-polar Add polar bonds (head to tail) CHH H H Molecules with polar bonds may not be polar overall!