Chemical Bonding and Nomenclature

1. Chemical Bonding and Nomenclature 1 Chapters 7,8 Adventures of Oxygen Clip

2. 2 2 GOALS 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. 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. 5. Identify substances based on chemical and physical properties

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

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

5. Ionic Bonding 5 Transfer of electrons from one atom to another atom. Occurs between metals & nonmetals. Compound composed of cations and anions. Electrically neutral Called compounds.

6. 6 OPPOSITS ATTRACT!

7. 9 7 Ionic Bonding CLIP

8. 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

9. 9 Covalent Bonding Occurs between nonmetals and nonmetals. The sharing of electrons between atoms. Each atom attempts to fill their valence shell. Called Molecules: Neutral group of atoms joined by a covalent bond

10. 10

11. 11 Hydrogen and Fluorine Hydrogen and Chlorine

12. 12 Single, Double, Triple

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

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

15. 15 Clip

16. 16 Unequal Sharing of Electrons Called Polar Molecules • 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 Polar molecules happen when one atom has a greater positivecharge Clip

17. Unequal Sharing of Electrons 17 Called Polar Molecules δ+ δ_ Practice Quiz

18. Animation

19. The shape may affect the polarity of an entire molecule • Ex CO2 (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. 18 Properties of Covalent Molecules • Many are gases or liquids at room temperature • Composed of two nonmetals. • Have low melting and boiling points

21. 19 • Ionic and Covalent Bonding Review Clip

22. Properties of Ionic and Covalent Compounds/Molecules

23. CO2 H2O NaCl MgCl2 NO2 Li2S NaF BeO HCl NaF KCl H2O2 N2 Cl2 20 Covalent or Ionic?

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. 24 Writing Chemical Formulas Goals revisited

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

29. 26 • 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 H2O.  Carbon Dioxide is CO2.  Why does oxygen combine in different ratios, in different compounds?  The chemical formula for table salt is NaCl. Calcium Chloride is CaCl2. Why does chlorine combine in different ratios, in different compounds? 

30. 27 The simplest compounds are ones with only two elements These are called binary KI, CO, H2O, NaCl

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

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

33. 30 +1 -1 +2 -1 Na Br Ca Br To make it ZERO, you need 1 Ca & 2 Br. NaBr CaBr2 Subscripts show the number of atoms of that kind in the compound

34. Some elements have more than one oxidation number 31 +3 -2 +2 -2 Fe O Fe O Fe2O3 FeO We call these elements- Multivalent Elements

35. K + Br Mg + Cl Ca + I K + O K + I Sr + Br Na + O Ga + Br Fe+2 + O Fe+3 + O Cu+2 + F Cr+3 + O Mg + O Al + P 32 Now You Try writing Binary Ionic formulas

36. Polyatomic Ions: 33 -a tightly bound group of covalently bonded atoms that has a positive or negative charge and behaves AS A UNIT.

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

38. Na + SO4 Mg + PO4 Ca + CO3 Na + OH Mg + OH NH4 + OH K + PO4 NH4 + NO3 H + SO4 Ca + SO4 K + NO3 Na + PO4 35 Try these

39. 36 Naming Chemical Formulas

40. 37 Naming Binary Compounds and Molecules • Steps: • If it is Binary- • Decide if it is an ionic or covalent bond. • Metal- nonmetal….. • Ionic • Nonmetal- nonmetal…. • Covalent Example: • NaCl

41. Only 2 elements Check to see if any elements are multivalent. If all single valent, write the name of the positive ion first. Write the root of the negative ion and add –ide. Examples: NaCl K2O AlCl3 BaF2 KI Li2O 38 If ionic …….

42. If ionic ……. 39 Examples: • FeO • Fe2O3 • CuO • Cu2O • PbCl4 • PbI2 • Check to see if any elements are multivalent. • If multivalent ions, determine the oxidation number of the element. • Use Roman numerals in parentheses after the name of the element. • Write the root of the negative ion and add –ide.

43. Use Greek prefix to indicate how many atoms of each element are in the molecule 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- 40 If Covalent... (Molecular Formula) Example: • NO • Nitrogen Monoxide • PCl3 • Phosphorous trichloride

44. Write the name of the positive ion. Write the name of the polyatomic ion. Examples: NaCO3 KNO3 NaC2H3O2 41 If it contains a polyatomic ion... Example: • KOH • Potassium Hydroxide • CaCO3 • Calcium Carbonate

45. KBr HCl MgO CaCl2 H2O NO2 CuSO4 CaSO4 NH4OH CaCO3 Cu(ClO3) 2 Cr2O 3 33 42 Name the following:

46. 43 Lewis Structures

47. Drawing Lewis Structures NH3 5 + 3(1) = 8 (nitrogen has five; each hydrogen has one) . N-8, H (2 each x 3=) 6… so TOE=14 14-8= 6 6/2= 3 bonds . . • 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).

48. Drawing Lewis Structures Double, triple bonds. CO32- • 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.

49. Try these… • CCl4 • NF3 • SH2 • H2O • CH4 • CO2 • BF3 • F2O • SO2 • SO3 • NF3 • N2 • NH4+ (notice the + charge) • NO3- (notice the - charge)

50. 44 Molecular Shapes