1. Molecules and Compounds

2. Valance Electrons and valance shell - Valance electrons are the electrons that are involved in chemical bonds Stable electrons have 8 valance electrons - this is called an octet of electrons Octet Rule - In order to achieve an octet atoms are willing to give up, take or share their own electrons o Exception is Helium and Hydrogen

3. Energy levels Rule: the higher the energy level, the more energy is needed for the electron to stay in that part of the electron cloud Matter tends to exist in its lowest energy state More stable in lowest energy Less likely to undergo change

4. Chemical Stability The noble gases are stable because they each have a complete outer energy level. Notice that eight dots surround Kr, Ne, Xe, Ar, and Rn, and two dots surround He.

5. Energy Levels and Other Elements When you look at the elements in Groups 13 through 17, you see that each of them falls short of having a stable energy level.

6. Outer Levels —Getting Their Fill How does hydrogen, or any other element, trying to become stable, gain or lose its outer electrons? They do this by combining with other atoms that also have partially complete outer energy levels. As a result, each achieves stability.

7. New Properties The compound formed when elements combine often has properties that aren’t anything like those of the individual elements. When atoms gain, lose, or share electrons, an attraction forms between the atoms, pulling them together to form a compound.

8. Formulas A chemical formula tells what elements a compound contains and the exact number of the atoms of each element in a unit of that compound.

9. Formulas The compound that you are probably most familiar with is H 2 O, more commonly known as water. A subscript written after a symbol tells how many atoms of that element are in a unit of the compound. No subscript means one atom is present

10. Atomic Stability The electric forces between oppositely charged electrons and protons hold atoms and molecules together, and thus are the forces that cause compounds to form. This attraction is called a chemical bond. A chemical bond is the force that holds atoms together in a compound.

11. Chemical Bonds —when atoms combine chemically —attractive force that holds atoms/ions together —electrons involved in bonding –fill energy levels in an orderly way –# valence e- determines how it bonds w/ other atoms —all elements with unfilled valence shells can form chemical bonds –**noble gases don't readily form bonds b/c valence shells are full —atoms seek to become more stable by obtaining full valence shells = more stable –electrons to fill these shells come from other atoms –atoms interact with each other in order to obtain needed electrons –these interactions cause chemical reactions

12. Gain or Loss of Electrons An atom that has lost or gained electrons is called an ion. An ion is a charged particle because it now has either more or fewer electrons than protons. It becomes charged because it loses or gains electrons It is the electric forces between oppositely charged particles, hold compounds together.

13. Ions positive ion: atom that releases e- negative ion: atom that's accepted e- **two ions attracted b/c of opposite charges —size of ions *lose electrons, decrease in size *gain electrons, increase in size

14. IONS —metals tend to lose electrons —nonmetals tend to gain electrons

15. 3 different kinds of chemical bonds ionic: electrons transferred from atom to atom, electrons are lost or gained. Forms between metals and nonmetals covalent: electrons shared between atoms. Forms between nonmetals metallic: electron sharing among many atoms. Forms between metals. Creates a “sea” of electrons that flow freely through a piece of metal.

16. Ionic Bonds Sodium - soft, silvery highly reactive metal -1 valence electron Chlorine - yellow-green gas - 7 valence electrons What happens? – Sodium loses one electron –Chlorine gains sodium's lost electron –NaCl is formed - a white crystal –Sodium and chlorine become ions –Na+ and Cl- –Join because opposites attract + attracts – Ionic Bonds – form mostly solids

17. Sodium Chloride

18. Zero Net Charge The result of this bond is a neutral compound. The compound as a whole is neutral because the sum of the charges on the ions is zero. Ionic bonds are a result of electronegativity Electronegativity - attraction that an atom has for the shared electrons (Trend – increases as you move up and to the right)

19. Covalent Bond - Sharing Electrons The attraction that forms between atoms when they share electrons is known as a covalent bond. A neutral particle that forms as a result of electron sharing is called a molecule.

20. Covalent bonds Electrons can be shared unequally Can form between the same or different types of atoms -form between two or more non-metals -form molecules with definite size -molecules can exist as solids, liquids, or gases -much more common than ionic bonding

21. Diatomic Molecules (8 total) when two atoms of the same type share electrons These are elements that are not found alone in nature Examples: Cl ², H ², O ², F ², Br ², I ², At ², N 2

22. COVALENTLY BONDED IONS -some molecules formed by covalent bonding tend to lose/gain electrons as a unit **polyatomic ion --can form ionic bonds just like ions

23. Unequal Sharing Part of the strength of attraction has to do with how far away from the nucleus the electron being shared is. The other part of the strength of attraction has to do with the size of the positive charge in the nucleus. Electrons are not always shared equally between atoms in a covalent bond.

24. Unequal Sharing Chlorine atoms have a stronger attraction for electrons than hydrogen atoms do. Results in a polar molecule One example of this unequal sharing is found in a molecule of hydrogen chloride, HCl. Click image to view movie

25. Polar or Nonpolar? A polar molecule is one that has a slightly positive end and a slightly negative end although the overall molecule is neutral. Water is an example of a polar molecule. A nonpolar molecule is one in which electrons are shared equally in bonds.

26. Oxidation Number tells how many electrons are lost, gained or shared Rules sign (+ or-) always written after the number not before) loses electrons - becomes positive - positive oxidation gained electrons - becomes negative - negative oxidation

27. EXAMPLE OF HOW OXIDATION IS WRITTEN Sodium always loses an electron When it loses an electron it now has one more proton than electrons The atom becomes positively charged Oxidation number is 1+ Chlorine always gains an electron When it gains an electron it now has one more electron than protons The atom becomes negatively charged Oxidation number is 1-

28. Oxidation numbers and the periodic table Group 1 = 1+ Group 2 = 2+ Group 13= 3+ Group 14= 4+/4- Group 15 = 3- Group 16= 2- Group 17 = 1- Group 18= 0

29. How to Write and name Chemical Formulas. Ionic Bonds –Name to Formula –Formula to Name –Working with Polyatomic Ions Name to Formula Formula to Name Covalent Bonds (Molecular Compounds) –Name to Formula –Formula to Name

30. Writing formulas and naming compounds Naming Binary compounds A binary compound is one that is composed of two elements. Binary compounds are formed mostly between metals and nonmetals

31. Compounds Are Neutral A formula must have the right number of positive ions and the right number of negative ions so the charges balance. When complete: Formulas have to be neutral

32. When writing a binary ionic compound You need to know 1. Elements involved 2. If they lose, gain, or share electrons to become stable (Oxidation Number) Oxidation Number is the same as the charge on the ion 3. Do they have more than ONE OXIDATION NUMBER?

33. Oxidation Numbers The oxidation number is expressed in the name with a roman numeral. For example, the oxidation number of iron in iron (III) oxide is 3+.

34. Writing a Binary (2 atoms) Ionic Formula Monatomic ions (single atoms) 1. Write the symbol for the monatomic ion that has a positive charge first –If the element has more than one oxidation number use a Roman numeral to indicate the correct oxidation number 2. Write the symbol for the monatomic ion that has a negative symbol second 3. Add subscripts so the sum of the positive and negative oxidation numbers is equal to zero (CRISS CROSS METHOD)

35. Write an Ionic formula Iron (III) and oxygen 1. Find the oxidation numbers of the elements Iron is 3+ because of the Roman numeral and Oxygen is 2- 2. Determine the ratios of each element and write using the criss-cross method to do this Fe 3+ O 2- Fe 2 0 3

36. SEPARATE PAPER Writing Ionic Formula Practice 1. Strontium Sulfide 2. Cesium Oxide 3. Beryllium Bromide 4. Manganese (II) Oxide

37. 5. Sodium Oxide 6. Rubidium Oxide 7. Aluminum sulfide 8. Strontium Phosphide

38. Naming Ionic Compounds 1. Write the name of the first element Check to see if the positive ion is capable of forming more than one oxidation number. If so, use the oxidation # to get the Roman numeral. 2. Write the root name of the second element (chlor is the root of chlorine) 3. Add the ending -ide to the root name (chloride) **Subscripts do not become part of the name for ionic compounds.

39. Elements with more than on Oxidation #

40. Writing Ionic Names 1. CuBr 2. MgS 3. Fe 2 O 3 4. Ba 3 P 2

41. 5. BeS 6. KCl 7. CaI 2 8. CrF 2

42. Compounds with Complex Ions Not all compounds are binary. Baking soda has the formula NaHCO 3. This is an example of an ionic compound that is not binary. They contain polyatomic ions.

43. Compounds with Complex Ions A polyatomic ion is a positively or negatively charged, covalently bonded group of atoms. So the polyatomic ions as a whole contains two or more elements.

44. Examples

45. Writing Formulas Using Polyatomic Ions Rules 1. Write the formula and the oxidation number for the positive ion 2. Write the formula and number for the negative ion 3. If the numbers add up to zero, just write the formulas. If not, criss cross and write the subscripts. If you have to criss cross you MUST put POLY in parenthesis. Example: Aluminum and sulfate: Al 3+, SO 4 2- Al 2 (SO 4 ) 3

46. Writing Poly Formulas 1. Calcium Phosphate 2. Ammonium Sulfate 3. Vanadium (V) Phosphate 4. Lead (IV) Nitrite

47. 5. Zinc hydroxide 6. Magnesium chlorate 7. Calcium Acetate 8. lithium sulfate

48. Naming Polyatomic Ionic Compounds Polyatomic 1. Write the name of the positive ion 2. Write the name of the negative ion

49. Naming Poly Ions 1. CaSO 4 2. Al 2 (SO 4 ) 3 3. Ba(ClO 3 ) 2 4. Mg(PO 4 ) 2

50. 5. CuI 6. Ca(NO 3 ) 2 7. Cr 2 O 3 8. KBr

51. Naming Covalent compounds Covalent compounds are those formed between elements that are nonmetals. Scientists use the Greek prefixes to indicate how many atoms of each element are in a binary compound.

52. Naming Covalent Compounds 1. write the name of the first nonmetal 2. write the name of the second nonmetal with the ending changed to -ide 3. insert prefixes before the element name to reflect subscripts in the formula: –EXCEPTION - never start a name with mono-

53. PREFIX MEANING and Covalent Naming Mono1 Di2 Tri3 Tetra4 Penta5 Hexa6 Hepta7 Octa8 Nona9 Deca10

54. Examples N 2 S 5 dinitrogen pentasulfide NO 2 nitrogen dioxide S 2 Cl 2 disulfur dichloride N 2 O 4 dinitrogen tetroxide SF 6 sulfur hexafluoride N 2 O 5 dinitrogen pentoxides

55. Naming Covalent Compounds 1. NI 3 2. BN 3. N 2 S 5 4. CCl 4

56. 5. SiO 2 6. P 4 O 10 7. N 2 O 3 8. Na 2 S

57. Writing Covalent formulas 1. write the symbol for the first nonmetal –use the prefix to determine the subscript that will follow the symbol 2. write the symbol for the second nonmetal –use its prefix to determine the subscript that will follow the symbol

58. Writing Covalent compounds 1. silicon dioxide 2. antimony tribromide 3. hexaboron silicide 4. diboron monosilicide

59. 5. xenon trifluoride 6. osmium tetroxide 7. chlorine dioxide 8. tetraphosphorus decoxide