Unit 6 Chemical Bonding

1. Unit 6Chemical Bonding Chemistry I Mr. Patel SWHS

2. Topic Outline • MUST know all assigned ions and elements!!! • Review Ions and Octet Rule (7.1) • Ionic Bonding (7.2) • Naming Ionic Compounds (9.2) • Metallic Bonding (7.3) • Covalent Bonding (8.1, 8.2) • Polarity (8.4) • Naming Covalent Molecules (9.3)

3. Ionic Bonding Intro

4. Metallic Bonding Intro

5. Covalent Bonding Intro

6. Ions • Ion – charged species • Must show the sign and value of charge • Valence electrons – electrons in highest occupied energy level • How do we find the number of valence electrons? • For elements 1A-8A = Group Number (except He) • Depict using Lewis Dot Structures

7. EX: Consider the element Aluminum. • How many valence electrons does Al have? • Draw the Lewis Dot Structure for Al.

8. The Octet Rule • Octet Rule – Atoms try to have 8 valence electrons • Goal: Be like a noble gas = stable • Will lose or gain electrons • Results in ions…What do we call these ions? • Cations – Positive charge species (metals) • Anions – Negative charge species (nonmetals)

9. EX: Consider the element Phosphorus. • How many valence electrons does P have? • Draw the Lewis Dot Structure for P. • Draw the Lewis Dot Structure for the ion form of phosphorus. • Will it form a cation or anion? Name it.

10. Ionic Bonding • Bond between Metal and Nonmetal • Actually, it is between cations and anions • Metal always comes first • Ionic bonding is due to the transfer of electrons • Important: The compound is always neutral • Positive = Negative

11. Ionic Bonding CaCl2 • Consider sodium chloride, CaCl2 • Metal first then nonmetal • Subscript tells you number of ions • 1 calcium ion for 2 chloride ions • Repeated array of ions – crystal • Chemical Formula – shows # of ions and smallest unit • Formula unit – lowest whole-number ratio of ions NaCl

12. CaCl2

13. Ionic Bond Formation Ionic Bond Formation

14. Ionic Bond Formation • There are 4 steps to diagram ionic bonding • Draw neutral Lewis Dot Structures (one with dots and other with x) • Show transfer of electrons (follow Octet Rule) • Show Resulting Ions • Write Formula

15. Use x and dots or different colors to show differences in valence electrons (VE). Step 1:Draw Lewis Dot Structure Ex: Show the Ionic Bond formation between sodium and chlorine. Na Cl x Step 2:Transfer the Electron(s) - Metal lose e-, Nonmetal gains.- Metal must lose all VE- Nonmetal must have 8 VE Na Cl x x 1- Step 3:Resulting Ions Show resulting ions – must have all charges. Anion must show the transferred electron. 1+ Na Cl NaCl Step 4:Chemical Formula Only show element symbols and subscripts – no charges, dots

16. Ca F x Step 1:Draw Lewis Dot Structure Ex: Show the Ionic Bond formation between calcium and fluorine. x - Metal lose e-, Nonmetal gains.- Calcium must lose 2 VE- Fluorine has 7 VE, can only take 1 more = Problem Ca Step 2:Transfer the Electron(s) F F x x Step 3:Resulting Ions We need to add another fluorine atom to take other VE from Calcium = Solution 2 1- 2+ Ca F x CaF2 Step 4:Chemical Formula There is one calcium and two fluoride ions in this bond.

17. Ex: Show the Ionic Bond formation between elements X (Group 3A) and Z (Group 6A).

18. Properties of Ionic Compounds • Arranged into a crystal lattice • Large attractive forces = stable, strong structure • Solid at room temperature • High melting points • Poor conductor as a solid • Good conductor when molten or in solution • Overall exothermic

19. Covalent Bonding • Bond between Nonmetal and Nonmetal • Can also include semimetals • NO IONS (cations/anions) • Covalent bonding is due to the sharing of electrons • Molecule – group of neutral atoms held together by covalent bonds

20. Covalent Bonding • Covalent molecules are defined structures • No crystal lattice • Has a specific 3-D structure • Molecular Formula – shows how many atoms of each element are in a molecule • We do not reduce formulas like ionic compounds • Ex: H2O, CO, CH4, C6H12O6

21. Depictions of Covalent Molecules

22. Covalent Molecule Shapes • Sharing of electrons are caused by overlapping and hybridizing orbitals (electron location) • VSEPR Theory – Valence Shell Electron Pair Repulsion Theory • VSEPR helps explain and predict the shape of molecules • Theory states that shape of molecules based on minimizing the repulsion of valence electron pairs • Keep electrons as far apart as possible

24. Methane = CH4 - Tetrahedral

25. Properties of Covalent Molecules • Distinct groupings of atoms = molecule • Solid, liquid or gas at room temperature • Low melting points • Poor conductor • Polar or Nonpolar

26. Diatomic Molecules • There are 7 elements that can not be found as individual atoms – found in pairs • Diatomic molecule – two atoms • H2 , N2 , O2 , F2 , Cl2 , Br2 , I2 (Group 7 + HON)

27. Types of bonds • Covalent molecules share bonds to complete octets – octet rule still applies! • Three types of bonds: single, double, triple

28. Comparing Bonds Valence Electrons not participating in bonding are called non-bonding electrons or lone pairs.

29. Polarity • Covalent Bonding is sharing of electrons • Electrons can be shared equally or unequally depending on the strengths of the atoms • If electrons have different electronegativities, the molecule will be polar • Like dissolves Like

30. Polarity • Polar – electrons shared unequally • Align themselves with an electric field • Ex: Water • Nonpolar – electrons shared equally • All diatomics are nonpolar

31. Metallic Bonding • These are the forces that hold metals together • Valence electrons are a sea of electrons around nuclei • Excellent conductors • Metals atoms arranged in compact and orderly patterns.

34. Comparing Ionic and Covalent Bonding

35. NOMENCLATURERULES

36. Nomenclature: Type 1Ionic Compounds with Fixed Charges • Groups 1A-7A have fixed charges…memorize these charges (Skip 4A and 8A) 1A: 1+ 2A: 2+ 3A: 3+ 5A: 3- 6A: 2+ 7A: 1- • Must be able to go from formula to name AND name to formula

37. Nomenclature: Type 1Ionic Compounds with Fixed Charges • Rules for Formula  Name: • Write down full name of the cation • Write down name of the anion (-ide) • Ex: K2O = potassium oxide • Practice: • H2SLiF Al2O3 • hydrogen sulfide, lithium fluoride, aluminum oxide

38. Nomenclature: Type 1Ionic Compounds with Fixed Charges • Rules for Name  Formula: • Write symbol and charge of cation and anion • Use subscripts to make all positive = negative (cross charges and reduce) • EX: LithiumPhosphide = Li1+P3- Li3P • Practice: • Magnesiumbromide, bariumsulfide, Calciumnitride • MgBr2BaSCa3N2

39. Nomenclature: Type 2Ionic Compounds with Variable Charges • Groups 1A-7A have fixed charges- charge always the same (Skip 4A and 8A) • Other metals (transition metals) do not have fixed charges – multiple possibilities for charge • We must indicate the specific charge • Example: Mg – always Mg2+ • Example: Mn – can be Mn1+, Mn5+, Mn6+, Mn7+

40. Nomenclature: Type 2Ionic Compounds with Variable Charges • Rules for Formula  Name: • Write down full name of the cation and anion (-ide) • Find total negative charge = total positive charge • Find charge on each cation • Write charge as Roman Numeral between cation and anion in name • Ex: FeCl3 = iron(III) chloride • Each Cl is 1- charge = b/c there are 3 Cl there is total of 3- • This means there is a total of 3+ so Fe must be 3+ • Write charge of Fe as roman numeral in name

41. Nomenclature: Type 2Ionic Compounds with Variable Charges Practice Formula  Name: Tin(IV) sulfide Copper(I) oxide Iron(II) phosphide • SnS2 • Cu2O • Fe3P2

42. Nomenclature: Type 2Ionic Compounds with Variable Charges • Rules for Name  Formula: • Write symbol and charge of cation and anion • Charge of cation comes from Roman Numeral • Use subscripts to make all positive = negative (cross charges and reduce) • EX: Cobalt(II) nitride = Co2+N3- Li3N2 • Charge of cobalt came from roman numeral • Charge of anion came from periodic table • Cross charges (positive = negative)

43. Nomenclature: Type 2Ionic Compounds with Variable Charges Practice Name  Formula: MnCl2 Fe2O3 CuS • Manganese(II) chloride • Iron(III) oxide • Copper(II) sulfide

44. Nomenclature: Type 3Ionic Compounds with Polyatomic Ions • The compounds have more than two elements • Must know polyatomic ions (page 257) • Treat the polyatomic ion as a single unit that WILL NOT CHANGE • Nitrate = NO31- 2 nitrates = (NO31-)2 • Must be able to go from formula to name AND name to formula

45. Nomenclature: Type 3Ionic Compounds with Polyatomic Ions • Rules for Formula  Name: • Write down full name of the cation • Use Roman Numerals is cation is transition metal • Write down name of anion (-ide or polyatomic ion) • Ex: Ba(OH)2 = barium hydroxide • Ex: Pb3(PO4)2 = lead(II) phosphate • Practice: • Fe(CN)3 Li2SO4 NH4C2H3O2 • Iron(III) cyanide, lithium sulfate, ammonium acetate

46. Nomenclature: Type 3Ionic Compounds with Polyatomic Ions • Rules for Name  Formula: • Write symbol and charge of cation and anion • Use subscripts to make all positive = negative (cross charges and reduce) • EX: Tin(IV) sulfite= Sn4+(SO32-) Sn(SO3)2 • Practice: • Calcium hydroxide, copper(I) nitrite, ammonium phosphate • Ca(OH)2CuNO2(NH4)3PO4

47. Nomenclature: Type 4Covalent Molecules • The molecules do not contain metals. • Need to know Greek prefixes

48. Nomenclature: Type 4Covalent Molecules • Rules for Formula  Name: • Write down full name of the first element • Write down modified name of second element (-ide) • Place Greek prefixes before each element name to denote the number of atoms • No mono prefix on first element • Ex: CO2 = carbon dioxide • Practice: • N2O5NO3 XeF6 • Dinitrogenpentoxide, nitrogen trioxide, xenon hexafluoride

49. Nomenclature: Type 4Covalent Molecules • Rules for Name  Formula: • Write symbol of both elements • Use prefixes as subscripts • EX: phosphorus pentafluoride= PF5 • Practice: • Dihydrogenmonoxide, sulfur heptachloride • H2O SCl7