1. 1 Bonding II: Types of substances 1) Ionic compounds = continuous Ionic bonding between weak loser metal and strong gainer nonmetal ex:NaCl Sodium Chloride

2. 2 Honors: Traditional system for Ionic Latin suffix with endings -ic (higher charge) or –ous (lower Charge) Ex: copper +1 OR +2 oxidation states Copper (I) oxideCopper (II) oxide Cu 2 OCuO Cuprous oxidecupric oxide Others: Ferrous / FerricFe = iron Stannous / StannicSn = tin plumbous / plumbicPb = lead What is the formula for ferrous chloride? FeCl 2

3. 3 Bonding II: Types of substances 1) Ionic compounds = continuous Ionic bonding between weak (loser) metal and strong (gainer) nonmetal 2) molecular compounds = covalent bonding inside molecules = IM forces between molecules between strong nonmetals ex: SO 2 Sulfur (IV) oxide Organics – with Carbon ex: CH 4 methane Acids – with H as the loser ex: HCl hydrogen chloride

4. 4 Organic molecular compounds (vs. inorganic) - Based on chain of carbon with H and other nonmetals attached. -Many molecular formulas can be reduced (simplified) to empirical (ratio) formulas Ex: glucose molecular formula: C 6 H 12 O 6 empirical formula: 6:12:6 = 1:2:1 ratio 6 CH 2 O

5. 5 Problem: What is the empirical formula for the molecular formula C 6 H 8 ? 6:8 = 3:4 ratio 2 C3H4C3H4

6. 6 Hydrogen compounds H + forms molecular (covalent) compounds with other nonmetals. named like ionic compounds when in their pure state ex: HClhydrogen chloride others: 1. H 2 S 2. H 2 SO 4 3. H 2 O 4. H 2 O 2 Hydrogen sulfide Hydrogen sulfate Hydrogen Peroxide [ O 2 ] 2- polyatomic ion Water (Sodium peroxide = Na 2 O 2 ) H - in back are hydrides Ex: Metal hydrides KH potassium hydride

7. 7 Acids: H + compounds (aq) in water Has properties of both ionic and molecular Made of molecules, but can break into ions in water Ex: HCl (g) → HCl (aq) = ( Separate H + & Cl - ions ) Hydrogenchloride H2OH2OH2OH2O Hydrochloricacid Unlike typical molecular compounds: acids are Gases and liquids that are soluble in water Solutions conduct electricity like ionic

8. 8 Acids w/ H in front Or COOH at end

9. 9 Bonding II: Types of substances 1) Ionic compounds = continuous Ionic bonding 2) molecular compounds = covalent bonding inside molecules Organic – with Carbon Acids – turn ionic in water 3) Network substances = continuous covalent bonding - C and Si Crystals - Elements that form 4 bonds (bonds in all directions)

10. 10 Network Solids: Continuous covalent bonds Hard, very high melting points Elements which form 4 bonds Ex: diamond (C), Sand (SiO 2 ) | | | --- C --- C --- C --- | | | --- C --- C --- C --- | | | Molecular: DiscreteParticles Ex: H 2 O Weak forces Between molecules = Low MP

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12. 12 Bonding II: Types of substances 1) Ionic compounds = continuous Ionic bonding 2) molecular compounds = covalent bonding inside molecules 3) Network substances = continuous covalent bonding 4) Metallic (pure and mixtures – alloys) continuous metallic bonds

13. 13 Learning Check: What are the four types of pure “substances” How can bonding type be surmised from the chemical formula of a substance? What identifies a compound as organic? Is the formula H 2 O an empirical or molecular formula? What identifies a substance as an acid? How is it different from other molecular substances? What kind of bonding is in a network structure? How is it different from molecular structures?

14. 14 Polarity in covalent (sharing) bonds Bond Polarity: Balance of charge in bonding is due to the difference in pull on electrons (atoms don’t share equally) No polarity Polar bond Atoms of same element BrINClHOF’s Bond is nonpolar Atoms of different elements Electrons pulled toward stronger atom. Bond is “polar”

15. 15 Problems 1.Which contain molecules with polar covalent bonds? a.H 2(g) b.HCl (g) c.CCl 4(l) d.NaCl (s) e.Fe (s) Different nonmetals Same = nonpolar Ionic, not covalent Metallic bonds

16. 16 Electronegativity (pulling strength) Ability of an atom to attract a pair of bonding electrons to itself Nonmetals Strong F is Most Electro- negative increasing decreasing Metals Weak Fr is Least Electro- Negative (“electro- positive”)

17. 17 Bond “Character” NaF HF OF 2 F-F Depends on difference in electronegativity between bonded atoms 4.0 - 4.0 = 0 No polarity in the bond 4.0 – 3.5 = 0.5 some polarity in the bond 4.0 – 2.1 = 1.9 high polarity in bond 4.0 – 0.9 = 3.1 Highest polarity is ionic bond 100% covalent character Most ionic character Further apart = more polarity = more ionic character

18. 18 Bond “Character” NaF HF OF 2 F-F 4.0 - 0.9 4.0 - 2.1 4.0 - 3.5 4.0 - 4.0 3.1 1.9 0.5 0.0  --------------------------------------------------------  Ionic charactercovalent character Difference in electronegativity between bonded atoms most polar fairly polar less polar no polarity  ------------- 1.65 ------------------ 0.7----------------------  Ionic polar covalent nonpolar covalent

19. 19 Practice problems 1. Which bond is most polar? a. H-H b. H-F c. H-O d. H-I 2. Which substance has the greatest ionic character? a. C-Cl b. H-Cl c. Cl-Cl d. Na-Cl 2.1- 2.1 = 0.0 4.0- 2.1 = 1.9 1.4 0.6 3.2 - 2.6 = 0.6 1.1 0.0 2.3 Only ionic compound! Strongest gainer With H (Greatest polarity)

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21. 21 Molecule Polarity Due to symmetry of bonds, molecules can be polar or nonpolar A Polar molecule: Is called a Dipole”  it has negative and positive poles (ends) Is Asymmetrical (not symmetrical) in molecule shape and not symmetrical in electron arrangement Negative Pole Positive Pole Ex: H --- Cl Electrons pulled toward stronger Cl end of molecule Dipolesare“sticky” molecules with different ends are Polar or “dipoles”

22. 22 If it has a front, back, top or bottom its polar. If its symmetrical its Nonpolar Polar or nonpolar? Nonpolar = symmetrical, “balanced” Polar = asymmetrical, “unbalanced” Nonpolar NonpolarNonpolar

23. 23 Predicting Polarity (predicting shape) VSEPR Theory - used to predict the shape and polarity of molecules from Lewis electron dot structure (Valence Shell Electron Pair Repulsion) Four Electron pairs repel each other to the 4 corners of a tetrahedron arrangement Bonded atoms get arranged into different shapes Ex: H 2 O angular Vesper Electron Pairs Act as Bonded atoms Ne Tetrahedron

24. 24 Polar shapes: asymmetrical charge shape and charge distribution Linear Polar Ex: HF, HI, etc. Two different elements IH x (+) H --- I (--) Angular (bent) polar Ex: H 2 O, OF 2 (+) H --- O (--) | H (+) H x H Ox F x F O x Angular Or straight? Why?

25. 25 Pyramid (Trigonal) Ex: NH 3 X X H x H Nx H Vesper electron pair pushes other Atoms into a pyramid Shape!

26. 26 Nonpolar molecule shapes Symmetrical shape and distribution of electron charge Look for C or Si in compounds Linear nonpolar Ex: BrINClHOF’s CO 2, CS 2 O=C=O O x Cx O xx Notice: nonpolar compounds usually contain C All C’s electrons are in the bond No Vesper electron pair to push

27. 27 Tetrahedral molecules CH 4 or CCl 4 Structural formulaStructural formula: each line = 2 electrons Notice: CX 4

28. 28 Less common shapes: Trigonal Planer BF 3 Nonpolar Polar CH 2 O

29. 29 Learning check: Predict the shape of each molecule and if it is Polar or Nonpolar: 1. O 2 7. He 2. CCl 4 8. N 2 3. HCl9. OF 2 4. H 2 O10. CH 2 F 2 5. NF 3 11. PCl 3 6. CS 2 12. NH 4 +

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31. 31 Intermolecular (“sticky”) forces Called Vanderwaal’s forces (“IM” forces) Weak Physical force between adjacent molecules (ex: determines how high melting points will be) Affected by polarity (shape) of molecules (Negative pole attracted to positive pole) Allows molecular substances to exist as liquids and low melting point solids (instead of gases) Two types: dipole attractions and dispersion forces

32. 32 Polar molecules (dipoles) are attracted to each other. Depends on extent of molecules polarity (ex: HCl vs. HF) Hydrogen bonds special type H bonded to small, strong atom: F, O, or N Found between molecules in HF, H 2 O, NH 3 Accounts for higher than expected MP and BP H is a naked proton HF is more polar = stickier “Dipole Attractions”

33. 33 What kind of question is this: Notice all have H bonded to another nonmetal Hydrogen bonding!

34. 34 Dispersion forces In large and Nonpolar molecules Random Dispersion of electrons creates temporary (+) and (-) poles Stronger in molecules with more electrons (larger) and as molecules get closer Ex: liquefaction of gases only at high pressure (molecules forced together) Ex: wax, dry ice (CO 2(s) ) “Big and sticky force” What explains the pattern of phases in group 17?

35. 35 Low temp = weak IM forces Phase change / boiling point questions: Usually relate to IM (sticky) forces

36. 36 46 At STP, fluorine is a gas and iodine is a solid. This observation can be explained by the fact that fluorine has (1) weaker intermolecular forces of attraction than iodine (2) stronger intermolecular forces of attraction than iodine (3) lower average kinetic energy than iodine (4) higher average kinetic energy than iodine gas = weaker IM forces gas = weaker IM forces F 2 smaller than I 2 Average Kinetic energy  Ie. How fast molecules are moving = Temperature At STP, both at same temperature

37. 37 Propane has weaker IM forces than butane OR Weaker IM forces in smaller molecules

38. 38 Can you explain the trends?

39. 39 Practice problems: 1.Which liquid contains the strongest intermolecular forces: He (l), Ne (l), Ar (l), Or Kr (l) ? Why? 2. Which has the strongest Hydrogen bonds? H 2, HF, HCl, or HBr? 3. Which is expected to have the lowest Boiling point? F 2, Cl 2, Br 2, or I 2 ? 4.What two kind of bonds are found in a sample of H 2 O (l) ?

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41. 41 41 Based on intermolecular forces, which of these substances would have the highest boiling point? (1) He (3) CH 4 (2) O 2 (4) NH 3 23 Hydrogen bonding is a type of (1) strong covalent bond (2) weak ionic bond (3) strong intermolecular force (4) weak intermolecular force Polar / H-bonds = strongest IM forces, highest BP Oxymoron: Strongest of the weak forces! Boiling point questions: Usually relate to IM (sticky) forces

42. 42 Each molecule listed below is formed by sharing electrons between atoms when the atoms within the molecule are bonded together. Molecule A: Cl 2 Molecule B: CCl 4 Molecule C: NH 3 59 Explain why NH 3 has stronger intermolecular forces of attraction than Cl 2. [1] Molecules like NH 3 are polar, vs. Cl 2 is nonpolar.

43. 43 Distinguishing between substances: Ionic compounds = Ionic bonding molecular compounds = covalent bonding metallic – pure or alloys = metallic bonds Network substances = covalent bonding Based on properties: Melting or boiling point / phase Solubility in water Electrical conductivity

44. 44 Electrical conductivity What is required to conduct electricity? Metals: yes as solid or liquid Ionic salts: no when solid, yes when (aq) in H 2 O or when fused (melted) into liquid Molecular – no, except acids when (aq) in water Requires particles which are: (1) charged AND(2) mobile Mobile electrons Mobile ions Electrolyte: substance that conducts electricity when dissolved in water Includes ionic salts and acids Ex: NaCl, HCl

45. 45 Metals w/ mobile electrons Molecular have Neutral molecules Ionic w/ mobile Ions when dissolved

46. 46 Ions in water Allow acids to Conductelectricity Strong acids = lots of ions formed Weak acids = few ions formed

47. 47 Properties of solids Ionic molecular metallic network Strong Weak Continuous strong Continuous intermolecularflexible continuous Ionic bonds forces metallic bonds covalent bonds Brittle solids soft solids malleable very hard High MP low MP MP varies extremely high MP Usually soluble not usually soluble NONO In water in watersolubility solubility Conductors: as (l) or (aq) Poor conductor good conductorpoor conductor (mobile ions)(neutral particles) (mobile electrons) (nothing is mobile)

48. 48 Review problems What kind of Substance? 1.A white crystalline substance, conducts when dissolved in aqueous solution. 2.A soft solid, melts at 75 0 C. Nonconductor 3.Silvery grey solid. Good electrical conductor 4.Crystalline solid. Melts at 1000 0 C. Nonconductor. 5.White crystalline solid. Dissolves easily, but aqueous solution doesn’t conduct electricity

49. 49 Review problems: What kind of bond accounts for these properties? (Describe) 1.White crystalline solid, melts at 700 0 C. Conducts electricity as a when fused (melted). 2.A liquid, evaporates readily. 3.A malleable solid. 4.A sticky viscous liquid. 5.Water’s high surface tension.

50. 50 Sample questions H with small e- gainer High melting point = strong forces:

51. 51 Salts are soluble: ionic Conducts e - when dissolved: ionic Strong bonds:

52. 52 Methane is CH 4