1. Unit 6 Chemical 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. a)How many valence electrons does Al have? b)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. a)How many valence electrons does P have? b)Draw the Lewis Dot Structure for P. c)Draw the Lewis Dot Structure for the ion form of phosphorus. d)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 Consider sodium chloride, CaCl 2 – 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 CaCl 2

13. Ionic Bond Formation

14. There are 4 steps to diagram ionic bonding 1.Draw neutral Lewis Dot Structures (one with dots and other with x) 2.Show transfer of electrons (follow Octet Rule) 3.Show Resulting Ions 4.Write Formula

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

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

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. Colligative Properties Ionic Compounds are known as salts Salts will ionize when in solution – Split into ions Colligative Properties – Properties that depend on the nature of the solute but not the quantity – Boiling Point Elevation – salts increase Boiling Point – Freezing Point Depression – salts decrease Freezing Point Number of ions determine the effect: CaCl 2 vs NaCl

20. 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

21. 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: H 2 O, CO, CH 4, C 6 H 12 O 6

22. Depictions of Covalent Molecules

23. 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

25. Methane = CH 4 - Tetrahedral

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

27. Diatomic Molecules There are 7 elements that can not be found as individual atoms – found in pairs Diatomic molecule – two atoms H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 (Group 7 + HON)

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

29. Comparing Bonds Bond Electrons Involved Bond Length Bond StrengthDiatomics Single Bond 2 electronsLongestWeakestH 2, F 2, Cl 2, Br 2, I 2 Double Bond 4 electronsModerateIntermediateO2O2 Triple Bond 6 electronsShortestStrongestN2N2 Valence Electrons not participating in bonding are called non-bonding electrons or lone pairs.

30. 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

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

32. 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.

35. Comparing Ionic and Covalent Bonding CharacteristicIonic BondingCovalent Bonding ElementsMetal and NonmetalNonmetal and Nonmetal Bond FormationTransfer electronShare electron Product of bondFormula Unit (salt)Molecule Physical StateSolidSolid, Liquid, Gas Melting PointHighLow ConductivityGood ConductorPoor to Non-conductor NomenclatureNo Prefixes Can Have Roman Numerals Always Prefixes No Roman Numerals Follow Octet RuleYES ForceIntramolecular

36. NOMENCLATURE RULES

37. Nomenclature: Type 1 Ionic 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

38. Nomenclature: Type 1 Ionic Compounds with Fixed Charges Rules for Formula  Name: – Write down full name of the cation – Write down name of the anion (-ide) – Ex: K 2 O = potassium oxide Practice: – H 2 S LiF Al 2 O 3 – hydrogen sulfide, lithium fluoride, aluminum oxide

39. Nomenclature: Type 1 Ionic 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: Lithium Phosphide = Li 1+ P 3-  Li 3 P Practice: – Magnesium bromide, barium sulfide, Calcium nitride – MgBr 2 BaS Ca 3 N 2

40. Nomenclature: Type 2 Ionic 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 Mg 2+ Example: Mn – can be Mn 1+, Mn 5+, Mn 6+, Mn 7+

41. Nomenclature: Type 2 Ionic 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: FeCl 3 = 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

42. Nomenclature: Type 2 Ionic Compounds with Variable Charges Practice Formula  Name: 1.SnS 2 2.Cu 2 O 3.Fe 3 P 2 1.Tin(IV) sulfide 2.Copper(I) oxide 3.Iron(II) phosphide

43. Nomenclature: Type 2 Ionic 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 = Co 2+ N 3-  Li 3 N 2 Charge of cobalt came from roman numeral Charge of anion came from periodic table Cross charges (positive = negative)

44. Nomenclature: Type 2 Ionic Compounds with Variable Charges Practice Name  Formula: 1.Manganese(II) chloride 2.Iron(III) oxide 3.Copper(II) sulfide 1.MnCl 2 2.Fe 2 O 3 3.CuS

45. Nomenclature: Type 3 Ionic 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 = NO 3 1- 2 nitrates = (NO 3 1- ) 2 Must be able to go from formula to name AND name to formula

46. Nomenclature: Type 3 Ionic 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: Pb 3 (PO 4 ) 2 = lead(II) phosphate Practice: – Fe(CN) 3 Li 2 SO 4 NH 4 C 2 H 3 O 2 – Iron(III) cyanide, lithium sulfate, ammonium acetate

47. Nomenclature: Type 3 Ionic 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= Sn 4+ (SO 3 2- )  Sn(SO 3 ) 2 Practice: – Calcium hydroxide, copper(I) nitrite, ammonium phosphate – Ca(OH) 2 CuNO 2 (NH 4 ) 3 PO 4

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

49. Nomenclature: Type 4 Covalent 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: CO 2 = carbon dioxide Practice: – N 2 O 5 NO 3 XeF 6 – Dinitrogen pentoxide, nitrogen trioxide, xenon hexafluoride

50. Nomenclature: Type 4 Covalent Molecules Rules for Name  Formula: – Write symbol of both elements – Use prefixes as subscripts – EX: phosphorus pentafluoride = PF 5 Practice: – Dihydrogen monoxide, sulfur heptachloride – H 2 O SCl 7