1. Essential Questions: What are the three types of bonds? What are the characteristics of each type? Types of Bonds:  Metallic  Ionic  Covalent

2. Metallic Bond: Results from the force of attraction of valence e-s for an atom positively charged kernel (nucleus). Often described as a sea of e-s Valence electrons- does not belong to a single atom, but rather are shared between the entire molecule.

3. Metallic Bond

5. Essential question: What are the characteristics of a metal? Characteristics of a Metal: Crystallization in structure Arranged in very compact and orderly patterns Have luster, malleability, ductility Conduct electricity (because of mobile e-s)

6. Alloy: Mixtures composed of two or more elements with one being a metal. Ex. Brass-Cu + Zn Sterling silver—-Ag + Cu Steel—Fe + C (B, Cr, Mn) Product is superior to reactants.

7. Covalent Bonds: Covalent Bonds: Sharing of e-s between two atoms to form a stable octet. Sharing of e-s between two atoms to form a stable octet. Usually between 2 non-metals exs. C1 2,SO 2 Usually between 2 non-metals exs. C1 2,SO 2 Covalent bonds usually occur in groups 4a, 5a, 6a, and 7a. Covalent bonds usually occur in groups 4a, 5a, 6a, and 7a. Molecule- neutral group of atoms joined by covalent bonds. Ex :O 2,F 2

8. Molecular Compounds: Formed by covalently bonded elements. Have low B.P and M.P Tend to be a liquid or gas at room temperature. Ex. H 2 O, SO 2, CO 2

9. Essential question: What is a diatomic molecule?

10. Monatomic- single atoms Exs. Ne, Ar. Diatomic- Molecule consisting of 2 atoms Exs. F 2, 0 2, H 2, Cl 2, Br 2, N 2,I 2

11. Molecular Formula: Chemical formula of a molecular compound. Chemical formula of a molecular compound. Shows the number and kinds of atoms Exs. H 2 O, O 2, C 6 H 12 O 6. Shows the number and kinds of atoms Exs. H 2 O, O 2, C 6 H 12 O 6.

12. Types of Covalent Bonding: Single Covalent Double Covalent Triple Covalent Coordinate Covalent Non-polar Covalent/Polar Covalent

14. Essential Questions: What are the types of covalent bonding? What are some examples of each type? Single Covalent- bond formed be sharing a pair of e-s Ex.H 2 H-H, H:H

15. Structural Formula: Represents the covalent bonds by dashes Shows arrangement of atoms Exs. H 2 Molecular H-H Structural H 2 0

16. Double Covalent Bonds- Sharing 2 pair of e-s. O 2 oxygen

18. Triple Covalent Bonds: Shares 3 pair of e-s. Ex: N 2

19. Coordinate Covalent- Bond in which the shared e-pair comes from 1 of the bonding atoms. Exs. NH 4 (ammonium ion) CO ( carbon monoxide)

21. Essential Questions: What is a nonpolar and a polar bond? How do you determine what type of bond is in a compound? Non-Polar Covalent Bond: Pair of e-s shared equally between 2 atoms. Formed between 2 atoms of equal electro -negativity values Table S Usually formed between diatomic molecules. Exs.: O 2, F 2

22. Polar Covalent Bond: Pair of e-s shared unequally between 2 atoms. Pair of e-s shared unequally between 2 atoms. Formed between 2 atoms of unequal electronegativity values Formed between 2 atoms of unequal electronegativity values Exs: HCl, Exs: HCl, CCl4 CCl4

23. Essential Questions: What is electronegativity? What is its significance to bonding? Electronegativity: Base on a scale of 0-4. Base on a scale of 0-4. It is how much an atom want to get an electron from another atom. It is how much an atom want to get an electron from another atom. The higher the number the stronger the need to take an electron. The higher the number the stronger the need to take an electron. Fluorine(4) is most electronegative atom. Fluorine(4) is most electronegative atom. Noble gas has no electronegativity values because they are stable and have no need to take and electron. Have 8 valence electrons. Noble gas has no electronegativity values because they are stable and have no need to take and electron. Have 8 valence electrons. Table S Table S

24. Bond Dissociation Energy (Bond Energy) Energy needed to break a bond in a covalently bonded atom. Energy needed to break a bond in a covalently bonded atom. Ex: H 2 The larger the amount of energy needed to break the bond, the stronger the bond is The larger the amount of energy needed to break the bond, the stronger the bond is Compounds with high bond energy tend to be unreactive. Compounds with high bond energy tend to be unreactive. Ex. :CH 4 (methane )

26. Bond Length: Distance between nuclei of 2 bonded atoms Measured in picometers 1pm=10 12 m Shorter the bond, the larger the bond energy

27. Essential Questions: What are network solids. What are the characteristics of network solids? Network Solids: High MP. Form hard brittle crystals Do not dissolve in Polar/Non-polar solvents Poor conductors of electricity in liquid and solid phase. Covalently bonded, to form a continuous pattern Ex. SiC, BeO, Mg 2 Si diamond

28. Essential Questions: What are ionic bonds? What are ionic substances? Ionic Bonds: Electrostatic forces that hold ions together. Electrostatic forces that hold ions together. Formed between 1 atom with low ionization energy and 1 with high electronegativity. Ex: NaCl Formed between 1 atom with low ionization energy and 1 with high electronegativity. Ex: NaCl Formed between a metal (cation) and a nonmetal (anion) Formed between a metal (cation) and a nonmetal (anion) Resulting compound is electrically neutral. Resulting compound is electrically neutral.. If electronegativity difference between 2 atoms is 1.7 or greater it is generally ionic. Exs: MgCl, KF. If electronegativity difference between 2 atoms is 1.7 or greater it is generally ionic. Exs: MgCl, KF

29. Ionic Solids: Ionic Solids: Are crystalline Are crystalline High MP High MP Conduct electricity and in a melted solution Conduct electricity and in a melted solution Hard, Brittle Hard, Brittle Dissolve in polar solvents Dissolve in polar solvents Ions are arranged in repeating patterns.Ex. NaCl Each Na is surrounded by 6 Cl and each Cl has 6 Na. Ions are arranged in repeating patterns.Ex. NaCl Each Na is surrounded by 6 Cl and each Cl has 6 Na. Strong, attractive forces therefore are very stable. Strong, attractive forces therefore are very stable. Formed between atoms with electronegativity difference greater than 1.7 Formed between atoms with electronegativity difference greater than 1.7 Exs. Cl, CsBr, CF 2 Exs. Cl, CsBr, CF 2

32. Essential Questions: What are polar and non polar molecules? How can you determine if a compound is polar or nonpolar?

33. Polar Molecules (Dipoles): Unbalanced charge distribution along bond. Ex: HF Non-Polar Molecule: Balanced charge distribution. Exs:CO 2,CCl 4 + - _ + _ _ _ _+ _

34. Polarity and Molecular Symmetry: Whether a molecule is polar depends on the symmetry of its (e-s) charge distribution.

35. Essential Question: What are intermolecular forces? Intermolecular Forces: Attractions between molecules Ex. Sucrose + H 2 O (Sucrose dissolves in H 2 0) Weaker than ionic or covalent bonds Includes Dipole forces, H bonds and London dispersion

36. Essential Question: What are dipole forces? Dipole Forces: Attraction between + and - ends of another molecule. Attraction between + and - ends of another molecule. Occurs in a polar molecule Occurs in a polar molecule Stronger the dipoles, the higher BP Stronger the dipoles, the higher BP Larger the molecule, the stronger the forces thus, higher BP. Larger the molecule, the stronger the forces thus, higher BP.

37. Dipole Forces in HCl and H 2 0

38. Essential Questions: What are Hydrogen bonds and Dispersion forces? H Bonds: Special case of dipole attraction Special case of dipole attraction Occurs only with a small highly electronegative element. Occurs only with a small highly electronegative element. Ex. O, N, F Are strongest between molecules in which hydrogen is covalent bonded to an element with high electronegative and a small atomic radius. Not a type of bond, but an attraction between molecules. Not a type of bond, but an attraction between molecules. Exs.: H 2 0,HF

41. London Dispersion Forces: Attraction between 2 non-polar molecules. Attraction between 2 non-polar molecules. Non-polar molecules can become dipoles for a short time thus leading to attractions. Non-polar molecules can become dipoles for a short time thus leading to attractions. Force decreases with decreasing molecular mass. Force decreases with decreasing molecular mass. Force decreases with increase in the distance between molecules Force decreases with increase in the distance between molecules

42. Essential Questions: What are cations and anions? What are the characteristics of each type? Cation: Ions with positive charge Ions with positive charge Usually metals Usually metals Groups la-3a, form cations with + charge = group number Groups la-3a, form cations with + charge = group number Exs:Cu 0  Cu+ 1 Cu 0  Cu +2 Cu 0  Cu +2 **Cations are named for the element and word ion Exs:Na +1 Sodium ion Ca +2 Calcium ion Ca +2 Calcium ion

43. Anions: ► Ion with negative charge ► Usually non-metals ► Charge can be determined by subtracting 8 from the group number Exs: Group 7a F 7-8 =-1 Exs: Group 7a F 7-8 =-1 Group 5a N 5-8=-3 Group 5a N 5-8=-3 **Anions are named for the stem of the element's name + ide Exs: Cl -1 Choride ion P -3 Phosphide ion **Majority of elements from group 4a and 8 do not form ions.

44. Transition Ions: Elements from group B Usually form more than 1 cation with different charges (oxidation numbers) Ex: Fe +2 or Fe +3 Transition elements form colorful solutions (pigments)

45. Essential Question: What are the methods of naming transition ions? 2 Methods of Naming Transition Ions:  Stock System  Root System Stock System: Preferred method Preferred method Uses a roman numeral in parenthesis after the name of the ion. Uses a roman numeral in parenthesis after the name of the ion.Exs: Fe +2 Iron (II) ion Fe +3 Iron (III) ion

46. Root System: Uses a root name for the element. Uses a root name for the element. Root name is derived from its Latin name. Root name is derived from its Latin name.Exs:Iron—Ferrum Fe +2 Ferrous ion (lower oxidation number) Fe +3 Ferric ion (higher oxidation number) Copper—cuprousTin—Stannous Doesn't not tell actual oxidation number Doesn't not tell actual oxidation number

47. Essential Question: What is a polyatomic ion?

48. Polyatomic Ions: Group of at least 2 atoms with electrical charges Behave as a single unit and carry a charge Most polyatomic anions end ite or ate. Exs: (NH 4 ) +1 Ammonium (CO 3 )-2 Carbonate Exs: ite SO 2 - Sulfite N0 2 - Nitrite ClO 2 - Chlorite Exs: ate SO 4 -2 Sulfate NO 3 - Nitrate ClO 3 - Chlorate Table E- Selected Polyatomic Ions

49. Essential Questions: What are binary compounds? How do you name binary compounds? Binary Compounds: Composed of 2 elements Composed of 2 elements Can be ionic or molecular Can be ionic or molecular

51. Rules For Naming Binary Ionic Compounds Put cation name is followed by anion name Use the proper name for the 1st element, the 2nd element, use stem and add the suffix ide. Exs. Chlorine + Magnesium Mg +1 Cl -1 O + Al O -2 Al +3 therefore Al +3 O -2 Aluminum Oxide Li +1 Br -1 Li -1 Br +1 Lithium Bromide Copper + Oxygen Cu +1 O -2 therefore Cu 2 O Cu +1 O -2 therefore Cu 2 O Copper (I) Oxide

52. Essential Question: How do you name compounds with polyatomic ions? Writing Formulas for Binary Ionic Compounds:  Write the symbol of the cation and then the anion.  Cross the oxidation number, therefore adding subscripts  Compound must = O.

53. Sodium Hydroxide Na +1 (OH )-1 NaOH Potassium Sulfate K +1 (SO 4 )-2 K 2 S0 4 Iron (III) Chlorate Fe +3 (ClO 3 )-1 Fe (C10 3 ) 3

54. Binary Molecular Compounds: Composed of 2 non-metals Composed of 2 non-metals Contains a covalent bond. Contains a covalent bond. Rules for Naming Binary Molecular Compounds: Rules for Naming Binary Molecular Compounds: Name elements in order listed in formula. Name elements in order listed in formula. Use prefixes- tell how many atoms of each element are present. Use prefixes- tell how many atoms of each element are present.

55. Prefixes: Mono-1 Hexta-6 Di-2 Hepta-7 Tri-3 Octa-8 Tetra-4 Nona-9 Penta -5 Deca-10 2nd element ends in ide. Exs: CO Carbon Monoxide SF 6 Sulfur Hexafluoride Cl 2 08 Bichloro Octoxide

56. IF 7 Iodine Heptafluoride. S 2 Cl 2 Disulfur Dichloride

57. Writing Binary Molecular Formulas: Use prefix in name to tell the subscript of each element in formula. Write correct symbols with appropriate subscripts. Ex. Dinitrogen Tetraoxide N 2 O 4