1. Introduction to Bonding Essential Question: What are the characteristics of an ionic bond and how are they formed?

2. What is a bond? A bond can be thought of as a force that holds groups of two or more atoms together and makes them function as a single unit Example : water O HH Bonds require energy to break and release energy when made

3. Types of bonds Ionic bonds - typically formed between metals and nonmetals Covalent bonds - typically formed between nonmetals Metallic bonds - formed between metals

4. Ionic Bonds Ionic bonding results from the transfer of electrons. Then, the opposite charges attract each other. Ionic bonds are strong Ion: Atoms that have lost or gained electrons and have gained a charge Oxidation number: Charge of the ion (positive or negative)

5. Ionic Bonds Metals lose electrons and become positive (cations) Nonmetals gain electrons and become negative (anions) Either way, atoms become ions and gain a charge

6. 0 Predicting Ionic Charges 1+ 2+ 3+3- 2- 1- 4+/-

7. Ionic Bonds Na and Cl –Na is a metal and likes to lose one electron and form a +1 ion. –Cl is a nonmetal and likes to gain one electron and form a -1 ion. –the final ionic compounds is NaCl Na + Cl - + NaCl The electrostatic interaction keeps them together!

8. Ionic Bonds They do this to achieve an octet!

9. Atomic Stability Noble gases have a filled and stable outer energy level (8 electrons), except He which has 2 This is known as the octet rule Argon Valence electrons = 8 1 1 2 3 4 5 6 7 8

10. Ionic Bonds and Compounds Writing Chemical Formulas for Ionic Compounds

11. E.Q.: How do you write the chemical formulas for ionic compounds?

12. Ionic Compounds Consist of cations (positive ions) and anions (negative ions) Usually composed of metals (cations) and nonmetals (anions) Oxidation numbers (charges of ions) are important

13. Types of Ions Monatomic ions: Consist of a single atom (Ex.: Na +, Cl -, Mg +2 ) Mon = Single or one Can often be determined from the Periodic Table

14. Types of Ions Polyatomic ions: Groups of atoms that behave as a unit and carry a charge (Ex.: NO 3 -, OH -, SO 4 -2 ) Poly = Many or several You will need a list of polyatomic ions to determine the name and formula

15. nitrateNO 3 -1 chromateCrO 4 -2 chlorateClO 3 -1 dichromateCr 2 O 7 -2 sulfateSO 4 -2 phosphatePO 4 -3 carbonateCO 3 -2 acetateC 2 H 3 O 2 -1 hydroxideOH -1 cyanideCN -1 ammoniumNH 4 +1

16. Binary Ionic compounds Writing Formulas Made up of two monatomic ions (metal and nonmetal) Ex.: Potassium and Chlorine Ex.: Calcium and Bromine

17. Binary ionic compounds Writing Formulas Transition Metals (Groups 3 through 12) and some metals in Groups 3A and 4A (except aluminum, cadmium, silver, and zinc) can have several oxidation numbers Ex.: Iron (III) and Oxygen Ex.: Copper (II) and Oxygen

18. Ionic Compounds with Polyatomic ions Writing Formulas Ions made up of more than one atom Act as individual ions Rules used for binary compounds still apply Use parentheses when more than one polyatomic ion is needed and use the appropriate subscripts outside of the parentheses

19. nitrateNO 3 -1 chromateCrO 4 -2 chlorateClO 3 -1 dichromateCr 2 O 7 -2 sulfateSO 4 -2 phosphatePO 4 -3 carbonateCO 3 -2 acetateC 2 H 3 O 2 -1 hydroxideOH -1 cyanideCN -1 ammoniumNH 4 +1

20. Ionic Compounds with Polyatomic ions Writing Formulas Example: Ammonium ion and chloride ion *IWB Example: Calcium ion and phosphate ion *IWB

21. 1. sodium chloride 2. calcium oxide 3. potassium hydroxide 4. magnesium sulfide 5. copper(II) carbonate 6. aluminum oxide 7. iron(III) oxide 8. sodium carbonate 9. aluminum hydroxide 10. ammonium nitrate 11. zinc nitrate 12. magnesium carbonate Practice Problems

22. Ionic Bonds and Compounds Naming Ionic Compounds

23. E.Q.: How do you write the chemical name for an ionic compound?

24. Naming ionic compounds: Rules for naming ionic compounds 1.Name cation first and then the anion CsBr Cs Br cation anion +-

25. Naming ionic compounds: Rules for naming ionic compounds 2.Monatomic cations use the element name Cs = cesium 3.Monatomic anions use root of the element and end with – ide Br = bromide CsBr = cesium bromide +1

27. Naming ionic compounds: Rules for naming ionic compounds 4.This applies to transition metals and some metals in groups 13 (3A) and 14 (4A) (more than one oxidation number): Identify the oxidation number for the transition metal in the name of the compound Use a Roman numeral in parentheses after name of cation

29. Naming ionic compounds: Rules for naming ionic compounds Examples: Fe and O = = 2- 2+ 2- 3+ FeO Iron (II) oxide Fe 2 O 3 Iron (III) oxide

30. Naming ionic compounds: Rules for naming ionic compounds 5.If compound has a polyatomic ion, simply use the name of the ion NaOH = Na OH = sodium hydroxide (NH ) S = NH S = ammonium sulfide +1 4 2 4 +1 -2

31. Practice Problems 1. MgCl 2 2. LiOH 3. ZnCO 3 4. K 2 S 5. FePO 4 6. Ag 3 N 7. Mn(CN) 2 8. AgC 2 H 3 O 2 9. BaI 2 10. PbS 2 Magnesium chloride Lithium hydroxide Zinc carbonate Potassium sulfide Iron (III) phosphate Silver nitride Manganese (II) cyanide Silver acetate Barium iodide Lead (IV) sulfide

32. COVALENT BONDING E.Q.: WHAT ARE COVALENT BONDS AND HOW DO ELEMENTS FORM THESE BONDS?

33. Covalent Bonds –Exist between nonmetals bonded together –Form when atoms of nonmetals share electrons –There are no ions involved in covalent bonding since no element is losing or gaining electrons –Electrons can be shared equally or unequally –The unequal sharing results in polar molecules

34. NAMING COVALENT COMPOUNDS AND WRITING THEIR FORMULAS E.Q.: WHAT ARE THE RULES TO NAMING COVALENT COMPOUNDS AND HOW ARE THE FORMULAS WRITTEN?

35. Naming Covalent Compounds Names are usually composed of two words - First is the name of the first element in the formula - Second is the name of the second element in the formula, but changing the ending to –Ide Example:HF = hydrogen fluoride

36. Naming Covalent Compounds If there is more than one atom of an element in the molecule, then we need to use prefixes to tell us how many there are… Number of AtomsPrefix 1 Mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

37. Prefix mono- This prefix can be used for any element, as long as the element is the second one in the formula….. Correct name… HI = Hydrogen monoiodide Incorrect name… HI = Monohydrogen monoiodide

38. Naming Covalent Compounds Some common names to some very important covalent compounds FormulaCommon Name Molecular Compound Name H Owater dihydrogen monoxide NHAmmonia nitrogen trihydride N Onitrous oxide (laughing gas) dinitrogen monoxide NOnitric oxide nitrogen monoxide CHmethane carbon tetrahydride 2 3 2 4

39. Naming Covalent Compounds Practice…… P 2 O 5 = diphosphorus pentaoxide CO = carbonmonoxide CF 4 = carbon tetrafluoride IF 5 = Iodinepentafluoride

40. Writing Formulas for Covalent Compounds Practice…… SbBr 3 Antimony tribromide = B 6 Si Hexaboron silicide = N2O3N2O3 Dinitrogen trioxide = PI 3 Phosphorus triiodide =

41. Metallic Bonds Metallic bonds exist between metals Occur when two metals, usually the same metal, are bonded together “sea of electrons” “delocalized electrons” ttp://www.youtube.com/watch?v=XAnTCYZP JsE&feature=bf_next&list=PLBFE28832E57 7A62Bttp://www.youtube.com/watch?v=XAnTCYZP JsE&feature=bf_next&list=PLBFE28832E57 7A62B

42. Regents Chemistry Electronegativity

43. How can we tell really tell which type of bond we have? Electronegativity – is the relative ability of an atom in a molecule to attract shared electrons to itself This tells us what type of bond we have; –Covalent, polar covalent or ionic Electronegativity values are determined by measuring the polarities of bonds between various elements to determine a specific value for each element

44. Electronegativity Electronegativity values for each element are obtained by using the Periodic Table In fact, there is a general trend in electronegativity we observe in the Periodic Table Electronegativity values increase across and up the Periodic Table –See table on pg. 332

45. We take the difference between the electronegativity values to determine exactly what type of bond exists, in essence the polarity of the bond See table 12.1 Electronegativity

46. Determining Bond Polarity If the difference between the electronegativity values is: –0.0 – 0.5: covalent bond (equal sharing) –0.6 – 1.6: polar covalent bond (unequal sharing) –1.7 – up: ionic bond (transferring electrons)

47. Examples Use your Reference Tables to determine the difference in electronegativity values and the type of bond for each of the following: –H-H –H-Cl –H-O –H-S H-F NaCl O 2 KBr Worksheet

48. Regents Chemistry Intro to valence electrons

49. Electrons in an atom Electrons surround the nucleus of an atom in specific energy levels or shells Each level can hold only a certain amount of electrons It is an atoms ability to the lose, gain or share electrons from its outer shell that determine its reactivity

50. The outer shell The outer shell in an atom contains the valence electrons Valence electrons can be lost, gained or shared to have eight electrons in the outer shell Each group on the table tells the number of valence electrons

51. Periodic Table Groups 1, 2, 13, 14, 15, 16, 17, 18 have 1,2,3,4,5,6,7,8 valence electrons, respectively We will not consider the transition metals See periodic table

52. Sharing to reach the Octet Rule The octet rule states that an atom cannot have more than 8 electrons in its outer shell Valence electrons are lost, gained or shared with other atoms to attain 8 electrons in the outer shell Eight valence electrons means a filled and happy shell - like the noble gases

53. Nonmetals share Nonmetals share electrons to reach eight valence electrons Single, double and triple bonds can be formed by sharing electrons

54. Metals + non-metals = lose/gain e - metals and nonmetals interact by losing and gaining electrons to reach 8 electrons (filled outer shell) The oxidation states on the periodic table represent this desire to move electrons ex: K+ want to lose 1 electron to reach noble gas configuration of eight electrons

55. Lewis structures: your assignment The reading and problems focus on drawing Lewis structures Lewis structures are a means to represent bond formation between atoms Covalent bonded compounds have different Lewis structures than Ionic bonded compounds

56. Example of a Lewis Structure CHH H H C H CH4 Covalent bonds

57. Regents Chemistry Lewis Structures

58. b The Lewis Structure is a representation of a molecule that shows how the valence electrons are arranged among the atoms in a molecule b We used dots around the elemental symbol to represent the valence electrons C

59. Single Lewis Structure - Practice b Draw four lone electrons first (if necessary) them pair them up b Draw Lewis Structures for the following atoms NaBe Al Br

60. Lewis Structures for Ionic Compounds b For Lewis Structures of ionic bonds the atoms are not joined but draw next to each other [ Br ]K+K+ - Bromine gains an electron to achieve the noble gas configuration of Krypton Potassium loses an electron to achieve the noble gas configuration of Argon example: KBr

61. Lewis Structures – Covalent Bonds b Hydrogen forms stable molecules when it shares two electrons b Two electrons fill Hydrogen’s valence shell b Helium does not form bonds because its valence shell is already filled; it is a noble gas b Second row non-metals Carbon through Fluorine from stable molecules when surrounded by eight electrons – the Octet Rule

62. Lewis Structures – Covalent Bonds b Valence electrons in covalent bonds can either be bonding pairs, if involved directly in the bond or lone pairs if not involved in the bond

63. Writing Lewis Structures - Rules b Obtain the total sum of the valence electrons from all of the atoms b Use one pair of electrons to form a bond between each pair of bound atoms. For convenience, a line (instead of a pair of dots) can be used to indicate each pair of bonding electrons b Arrange the electrons to satisfy the duet rule for hydrogen and the octet rule for second row non metals non metals

64. Lewis Structures – Covalent Bonds b Examples PH 3 PH Step 1) 8 total valence e - total Step 2) Draw one pair of electrons per bond 8-6 = 2 left Step 3) Arrange the remaining electrons according to octet rule H H H l H– P –H

65. Lewis Structures – Covalent Bond Practice Examples HBr.. H:Br: ·· CF 4 Worksheet

66. Regents Chemistry –Ionic Lewis Structures –Multiple bonds in Lewis Structures –Polyatomic ion Lewis Structures and Resonance

67. Lewis Structures for Ionic Compounds b For Lewis Structures of ionic bonds the atoms are not joined but draw next to each other [ Br ]K+K+ - Bromine gains an electron to achieve the noble gas configuration of Krypton Potassium loses an electron to achieve the noble gas configuration of Argon example: KBr

68. Examples of Ionic Lewis Structures b Draw Lewis Structures for the following: NaClLiBrKI

69. Multiple Bonds and Lewis Structures…review first b We have seen how to draw Lewis Structures for molecules with single bonds For exampleFor example NH3 NH3 1.Sum the total valence e - 2.Subtract number of bonding e- 3.Place remaining valence e - 8 total valence e- 3 bonds x 2e- = 6 bonding 2 e - left over H HHN

70. Multiple Bonds b Between atoms of the same element b Example OxygenOxygen O O O O O = O Also a Lewis Structure Just O = O is called a structural model

71. Example of Multiple Bonds Nitrogen N We now meet the octet rule!

72. Multiple Bonds b Between atoms of different elements b CO 2 COO O = C = O We must use double bonds to meet the octet rule!

73. Lewis Structures for Polyatomic Ions and Resonance Structures b Read pg. 344 (bottom) to 349 and answer questions a-g in example 12.4 (pg. 347) and a-i in the Self Check exercise 12.4 (pg. 348) 12.4 (pg. 347) and a-i in the Self Check exercise 12.4 (pg. 348)