Chemical Bonding

1. Chemical Bonding 1 Adventures of Oxygen Clip

2. 2 GOALS 1. Compare & contrast ionic and covalentbonds in terms of electron position. 2. Determine the Types of ionsformed by representative elements 3. Develop and use models to predict formulas for stable, binary ionic compounds based on balance of charges. 5. Analyze and interpret data to predict properties of ionic and covalent compounds. (Properties like: types of bonds formed, elemental composition, melting point, boiling point, and conductivity.) 34 16 (IUPAC) International Union of Pure and Applied Chemistry

3. 3 Why do Atoms Form Compounds? • Stability. • What makes an atom stable? • Full outer energy level. • Eight. • They can either…… • 1) Gain electrons • 2) Lose electrons • 3) Share electrons

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. Remember: Atoms need a full outer energy level to be stable. EIGHT! Called compounds.

6. 6 Ionic Bonding Occurs between metals and nonmetals. Metals are electron donors. SO, they become POSITIVE Non-metals are electron accepters. SO, they become NEGATIVE. Opposites ATTRACT!

7. 7 When Atoms gain or lose electrons, they are called Ions. 3P 3P 3P Anion 3P 3P 3P Cation

8. 8 • Metals lose electrons to become stable. • Nonmetals gain electrons to become stable.

9. 9 Atoms can gain or lose electrons Ionization: requires energy Why do atoms lose and gain electrons? To become more stable. Stability=full outer energy level

11. 11 Animation Examples Opposites Attract!

12. 13 Ionic Bonding CLIP

13. 14 Covalent Bonding Occurs between nonmetals and nonmetals. The sharing of electrons between atoms. Called Molecules.

14. 15

15. 16 Hydrogen and Fluorine Hydrogen and Chlorine

16. 17 Single, Double, Triple 2 e- 4e- 6e-

17. 18 Clip

18. 19 Unequal Sharing Called Polar δ+ δ_ Polar molecules happen when one atom has a greater positive charge

19. CO2 NaCl H2O MgCl2 NO2 Li2S NaF BeCl2 BeO HCl KCl H2O2 N2 Cl2 AgCl2 21 Covalent or Ionic?(write the formula, then write “C” or “I”) clip

20. 12 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

21. 20 Properties of Covalent Molecules • Many are gases or liquids at room temperature • Composed of two nonmetals. • Have low melting and boiling points… covalent bonds are weaker than ionic bonds

22. Ionic and Covalent Bonding Review Clip

23. Do starter #1 on page 3 • a) for part II

24. Writing Chemical Formulas Goals revisited

25. 22 • 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? 

26. 23 The simplest compounds are ones with only two elements These are called binary KI, CO, H2O, NaCl

27. Oxidation numbers +4 -4 +1 0 -2 +2 +3 -3 -1 Tell you how many electrons an atom must gain, lose or share to become stable. 24

28. 25 Oxidation numbers We can predict the ratio of atoms in ionic compounds based on their oxidation numbers +1 -1 K Cl 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!

29. Smartboard examples

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 26

31. 27 K + Br Mg + Cl Ca + I K + O K + I Sr + Br Na + O Ga + Br Writing Binary Ionic formulas

32. Some elements have more than one oxidation number (Chart p588) -2 +2 -2 +3 Co O Co O Co2O3 CoO We call these elements- Multivalent Elements 28

33. Fe+2 + O Fe+3 + S Cu+2 + F Cr+3 + Br 29 Multivalent Practice

34. Naming Chemical Formulas

35. Naming Binary Compounds and Molecules 32 • Steps: • If it is Binary- • Decide if it is an ionic or covalent bond. • Metal- nonmetal….. • Ionic • Nonmetal- nonmetal…. • Covalent Example: 1-NaCl

36. 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: CaO K2S AlCl3 Ba3N2 Example: • NaCl • MgI2 If ionic …….

37. 33-Practice • KI • Cs2S • Fr3P • SrCl2

38. If ionic ……. NiCl Mn2S Cr3P2 Fe2O3 34 • 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.

39. Ionic-multivalent Examples: • FeO • Fe2O3 • CuO • Cu2O

40. 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- 36 Example 1-PCl3 If Covalent... • PCl3 • Phosphorous trichloride • NO • Nitrogen Monoxide

41. 36 ¾ Naming Covalent Practice • P4S5 • SF6 • N2O5 • H2O • NF3 • SiO2 • P2Br4 • SO3 • Tetraphosphorus Pentasulfide • Sulfur Hexafluoride • Dinitrogen Pentaoxide • Dihydrogen Monoxide • Nitrogen Trifluoride • Silicon Dioxide • Diphosphorus Tetrabromide • Sulfur Trioxide Greek Prefixes 1- mono- 2- di- 3- tri- 4- tetra- 5- penta- 6- hexa- 7- hepta- 8- octa-

42. Vocab not in book…. • Subscript: Small numbers after an element’s symbol that indicate how many atoms of that element are found in the molecule. • Chemically Stable: This term describes an atom whose outer energy level is full.

43. Do this. Put answers on p.3 in packet. Then do Starter #5 (p.4) KBr HCl MgO CaCl2 H2O NO2 CaS Cr2O3 FeO LiBr Name the following: Potassium Bromide Hydrogen Monochloride Magnesium Oxide Calcium Chloride Dihydrogen Monoxide Nitrogen Dioxide Calcium Sulfide Chromium (III) Oxide Iron (II) Oxide Lithium Bromide 37

44. End of Chemical Bonding Study Packet

45. 38 Writing & Balancing Chemical Equations Goals revisited

46. GOALS 1. Apply the Law of Conservation of Matter by balancing following types of chemical equations: • Synthesis • Decomposition • Single Replacement • Double Replacement 2. Demonstrate the Law of Conservation of Matter in a chemical reaction

47. 39 Chemical Reactions • A chemical reaction is a change in which one or more substances are converted into new substances. • Rearrangement of bonds in compounds and molecules. • Chemical Equations make it possible to see clearly what is happening during a chemical reaction

48. 40 Chemical equations are a shorthand way to show chemical reactions. Reactants Products H2 + O2 H2O Read as “yields” or “produces”

49. 41 Conservation of Mass The mass of the products always equals the mass of the reactants

50. 42 H2 + O2 H2O Does this meet the Conservation of Mass Law? Must Balance the Equation to show Conservation of Mass. 2 Hydrogen atoms 2 Oxygen atoms 2 Hydrogen atoms & one Oxygen atom