1. 1 Adventures of Oxygen Clip Chapters 7,8

2. 1. Compare and contrast types of chemical bonds (i.e. ionic, covalent). 2. Predict formulas for stable ionic compounds (binary and tertiary) based on balance of charges.GOALS 2 3. Use IUPAC nomenclature for both chemical names and formulas: Ionic compounds (Binary and tertiary) Covalent compounds (Binary and tertiary) 4. Apply concepts of the mole and Avogadro’s number to conceptualize and calculate empirical/molecular formulas, mass, moles and molecules relationships. 2 5. Identify substances based on chemical and physical properties

3. Why do Atoms Form Compounds? Stability. What makes an atom stable? Full outer energy level. Eight. 3

4. A Chemical Bond holds atoms together in a compound. Two basic typesTwo basic types: 1-Ionic 2-Covalent 4

5. Ionic Bonding T r a n s f e r o f e l e c t r o n s f r o m o n e a t o m t o a n o t h e r a t o m. O c c u r s b e t w e e n m e t a l s & n o n m e t a l s. Electrically neutral C a l l e d c o m p o u n d s. C o m p o u n d c o m p o s e d o f c a t i o n s a n d a n i o n s. 5

6. OPPOSITS ATTRACT! 6

7. CLIP 9 7

8. Properties of Ionic Compounds Crystalline solids at room temperature. Arranged in repeating three- dimensional patterns Have high melting points Can conduct electricity when melted or dissolved in water 8

9. Covalent Bonding T h e s h a r i n g o f e l e c t r o n s b e t w e e n a t o m s. E a c h a t o m a t t e m p t s t o f i l l t h e i r v a l e n c e s h e l l. O c c u r s b e t w e e n n o n m e t a l s a n d n o n m e t a l s. C a l l e d M o l e c u l e s : N e u t r a l g r o u p o f a t o m s j o i n e d b y a c o v a l e n t b o n d 9

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11. Hydrogen and Fluorine Hydrogen and Chlorine 11

12. Single, Double, Triple 12

13. Single Covalent Bonds (2e-) Structural Formula: dashes Unshared pair 13

14. Double and Triple Covalent Bonds Double bond- 2 pairs (for a total of 4) Triple bond- 3 pair (for a total 6) 14

15. 15 Clip

16. 16 The element that has a greater electronegativity attract the electrons more So, the electronegativity difference between two atoms tells you what kinds of bond is likely to form Unequal Sharing of Electrons Polar molecules happen when one atom has a greater positive charge Called Polar Molecules

17. Unequal Sharing of Electrons δ+δ+ Called Polar Molecules δ_δ_ 17

18. Animation

19. The shape may affect the polarity of an entire molecule Ex CO 2 (2 polar bonds cancel each other) The presence of a polar bond in a molecule often makes the entire molecules polar. (Water molecule) A molecule that has 2 poles is called a dipolar molecules, or dipole.

20. Properties of Covalent Molecules Many are gases or liquids at room temperature Composed of two nonmetals. Have low melting and boiling points 18

21. Ionic and Covalent Bonding Review ClipClip 19

22. Properties of Ionic and Covalent Compounds/Molecules

23. 1.CO 2 2.H 2 O 3.NaCl 4.MgCl 2 5.NO 2 6.Li 2 S 7.NaF 9.BeO 10.HCl 11.NaF 12.KCl 13.H 2 O 2 14.N 2 15.Cl 2 Covalent or Ionic? 20

24. Metallic Bonds Valence electrons (1-3) can be thought of as a sea of electrons. They are “mobile” and can easily drift freely from one part of the metal to another. Metallic bonds consist of the attraction of the free-floating valence electrons for positively charges metal ions. 21

25. Other Atomic Attractions Intermolecular attractions are weaker than either ionic or covalent bonds. Van der Waals Forces –Weak attraction consisting of dipole interactions and dispersion forces –Dipole interactions: when polar molecules are attracted to another. –Dispersion Forces: weakest of all interactions. Caused by motion of electrons. Occurs between nonpolar molecules. Temporary polarity. 22

26. Hydrogen bonding Found in many biological molecules Important in the properties of water. Attraction between hydrogen (when bonded to a very electronegative element) and another molecule. About 5% the strength of an average covalent bond. 23

27. Goals revisited 24

28. Ionic Bonding- Formula Units AA formula unit is the lowest whole-number ratio of the ions in an ionic compound. AA chemical Formula shows the kinds and numbers of atoms in the smallest representative unit of a substance. HHow do you figure out the “Chemical Formula?” 25

29. Writing chemical formulas is a shorthand way of indicating what a substance is made of. These formulas also let you know how many atoms of each type are found in a molecule. The chemical formula for water is H 2 O. Carbon Dioxide is CO 2. Why does oxygen combine in different ratios, in different compounds? The chemical formula for table salt is NaCl. Calcium Chloride is CaCl 2. Why does chlorine combine in different ratios, in different compounds? 26

30. The simplest compounds are ones with only two elements These are called binary KI, CO, H 2 O, NaCl 27

31. +1 +2 -2 -3+3 +4 -4 0 Oxidation numbers Tell you how many electrons an atom must gain, lose or share to become stable. 28

32. We can predict the ratio of atoms in ionic compounds based on their oxidation numbers Oxidation numbers K Cl +1 KCl Tells you how many electrons an atom must gain, lose or share to become stable. 1 valence electron 7 valence electron All compounds are neutral That means the overall charge is ZERO! 29

33. Subscripts show the number of atoms of that kind in the compound Na Br +1 NaBr Ca Br +2 CaBr 2 To make it ZERO, you need 1 Ca & 2 Br. 30

34. Some elements have more than one oxidation number Fe O +3 -2 Fe 2 O 3 Fe O +2-2 FeO We call these elements- Multivalent Elements 31

35. Now You Try writing Binary Ionic formulas 1.K + Br 2.Mg + Cl 3.Ca + I 4.K + O 5.K + I 6.Sr + Br 7.Na + O 8.Ga + Br 9.Fe +2 + O 10.Fe +3 + O 11.Cu +2 + F 12.Cr +3 + O 13.Mg + O 14.Al + P 32

36. Cations:ammonium, NH 4 + Anions: nitrate, NO 3 - sulfate, SO 4 2- hydroxide, OH - phosphate, PO 4 3- carbonate, CO 3 2- chlorate, ClO 3 - permanganate, MnO 4 - chromate, CrO 4 2- Polyatomic Ions: -a tightly bound group of covalently bonded atoms that has a positive or negative charge and behaves AS A UNIT. 33

37. Polyatomic Ions -Compounds containing polyatomic ions include both ionic and covalent bonding Writing Formulas Examples: Sodium and Nitrate Magnesium and Chlorate Ammonium and Sulfate 34

38. Try these 1. Na + SO 4 2. Mg + PO 4 3. Ca + CO 3 4. Na + OH 5. Mg + OH 6. NH 4 + OH 7. K + PO 4 8. NH 4 + NO 3 9. H + SO 4 10. Ca + SO 4 11. K + NO 3 12. Na + PO 4 35

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40. Naming Binary Compounds and Molecules Steps: –I–If it is Binary- 1.Decide if it is an ionic or covalent bond. –M–Metal- nonmetal….. »I»Ionic –N–Nonmetal- nonmetal…. »C»Covalent Example: NaCl 37

41. If ionic ……. 2.O nly 2 elements 3.C heck to see if any elements are multivalent. 4.I f all single valent, write the name of the positive ion first. 5.W rite the root of the negative ion and add –ide. Examples: 1.NaCl 2.K 2 O 3.AlCl 3 4.BaF 2 5.KI 6.Li 2 O 3838

42. If ionic ……. 5.Check to see if any elements are multivalent. 6.If multivalent ions, determine the oxidation number of the element. 7.Use Roman numerals in parentheses after the name of the element. 8.Write the root of the negative ion and add –ide. Examples: 1.FeO 2.Fe 2 O 3 3.CuO 4.Cu 2 O 5.PbCl 4 6.PbI 2 39

43. If Covalent... (Molecular Formula) 2.Use Greek prefix to indicate how many atoms of each element are in the molecule 3.Add -ide to the more electronegative element Greek Prefixes 1- mono- 2- di- 3- tri- 4- tetra- 5- penta- 6- hexa- 7- hepta- 8- octa- Example: NO Nitrogen Monoxide PCl 3 Phosphorous trichloride 40

44. If it contains a polyatomic ion... 2.Write the name of the positive ion. 3.Write the name of the polyatomic ion. Examples: 1.NaCO 3 2.KNO 3 3.NaC 2 H 3 O 2 Example: KOH Potassium Hydroxide CaCO 3 Calcium Carbonate 41

45. Name the following: 1.KBr 2.HCl 3.MgO 4.CaCl 2 5.H 2 O 6.NO 2 7.CuSO 4 8.CaSO 4 9.NH 4 OH 10.CaCO 3 11.Cu(ClO 3 ) 2 12.Cr 2 O 3 33 42

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47. Drawing Lewis Structures Step #1: Add up the number of valence electrons that should be included in the Lewis Structure. (TVE) Step #2: Calculate # of bonds. –Determine TOE: Theoretical Octet Electrons –TOE- TVE from step1 –Divide by 2 ( 2 electrons for each bond) Step #3: Draw the “skeleton structure” with the central atoms and the other atoms, each connected with a single bond. Step #4: Any “leftover” electrons so that all elements meet octet rule (or full outer energy). NH 3 1.5 + 3(1) = 8 (nitrogen has five; each hydrogen has one) 2.. N-8, H (2 each x 3=) 6… –so TOE=14 –14-8= 6 –6/2= 3 bonds 3.. 4..

48. Double, triple bonds. Same as last except… Step #4: If there are no electrons left, move electrons from a different atom to form another bond…double Side note: When more than one Lewis structure can be drawn, the molecule or ion is said to have resonance. CO 3 2- Drawing Lewis Structures

49. Try these… 1.CCl 4 2.NF 3 3.SH 2 4.H 2 O 5.CH 4 6.CO 2 7.BF 3 8.F 2 O 9.SO 2 10. SO 3 11.NF 3 12.N 2 13.NH 4 + (notice the + charge) 14.NO 3 - (notice the - charge)

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51. Molecules Have Shapes. VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of atoms about a central atom in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom. The number of electron pairs around the central atom can be determined by writing the Lewis structure for the molecule. The geometry of the molecule depends on the number of bonding groups (pairs of electrons) and the number of nonbonding electrons on the central atom.

52. Molecular Shapes VSEPR Theory: (Valence electron-pair repulsion theory) The repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible Lone pairs have more repulsive force than do shared electron pairs, and thus they force the shared pairs to squeeze more closely together. The repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible Lone pairs have more repulsive force than do shared electron pairs, and thus they force the shared pairs to squeeze more closely together. Linear Tetrahedral PyrimidalTrigonal Planar Practice: Go back to Lewis Structure Practice, and predict shapes. Bent

53. Shapes and Polarity Molecules can be polar, and when they are polar, they are called dipoles. Dipoles are molecules that have a slightly positive charge on one end and a slightly negative charge on the other Shape can help determine polarity Molecules that are symmetrical tend to be nonpolar. Molecules that are asymmetrical tend to be polar 45 Practice: Go back to Lewis Structure Practice, and predict polarity.

54. H 2 O- Bent-Polar SO 3 -trigonal planar- nonpolar. BF 3 - trigonal planar- nonpolar. SO 2 -bent-polar Molecules in Motion Website

56. Starter 7