2. 1.1 Inside the Atom

3. Inside the Atom 4 The atom is the ________ part of any element that retains all of the _________ of that ____________. 4 The atom is made up of: protons (____) electrons (____) neutrons (____) 4 The central core of the atom is called the _________, containing _______ and _________. (Nucleus charge = +)

4. Inside the Atom Atomic Mass Unit (AMU): The unit for ___________ of an _______. 1 Proton = _____, 1 Neutron = _____, 1 Electron = ______ Mass Number (AKA: ): Total ______ of the entire atom, which is the sum or the _________ and __________. Atomic Number: The total number of _________ in an atom, which is _________ for each _________ and never changes.

5. PROTONS (________, 1 AMU) # protons = _______ # ELECTRONS (negative, _______) # electrons = # _________ NEUTRONS (_________, 1 AMU) # neutrons = ___________- atomic #

6. ProtonsNeutronsElectrons In these notes, determine the number of protons, electrons, and neutrons for the following elements: 4 Oxygen 4 Argon 4 Magnesium 4 Neon 4 Nitrogen 4 Chlorine

7. 1.2 Modern Atomic Theory

8. Bohr Model  Created by _______ Bohr  Identified _________ in orbits at _______ distances from the _______.  The ______ are called ______ levels or electron shells.

9.  The ________ are located in ________levels around the nucleus.  The energy levels ___________ to the nucleus are the ___________ energy.  As you move ________ from nucleus energy __________. Atomic Theory and Energy Levels Energy Level 1 (E 1 ) = Energy Level 2 (E 2 ) = Energy Level 3 (E 3 ) = Energy Level 4 (E 4 ) =

10.  Electrons can jump from ______ levels of energy to _______ levels of energy if they _______ the necessary energy.  Electrons can jump from ________ levels of energy to _______ levels of energy and the will _______energy, as light.

11. Electron Fill Rule (_______ rule) 4 Electrons fill energy levels from ___________ to ___________ as follows

12. Valence Electrons 4 The ___________ energy level contains valence electrons. 4 These electrons determine the ___________ of an atom. 4 All atoms desire a ______ outer energy level.

13. 2.1 How Elements Are Organized

14. Periodic Table Developed by _________________ by listing all the ______________ on cards and arranging them in _____________ that fit best. He saw a repeating pattern start to appear so he called it the periodic table (Periodic meaning __________). The columns were called __________ or families The rows were called ____________.

15. Periodic Box 7 N 14.0067 Element Symbol Atomic Number: The number of _______ in element Atomic Mass: the total _______ of one atom of this element

16. GROUPS: Vertical ___________. Elements in groups have _______ chemical properties. (____ total) Periods: Horizontal _____. Each element in a period has ________ properties. (____ total)

17. Periodic Table 4 The table was not complete will all ________ elements like today. 4 Most of elements were not discovered yet, but with Mendeleev’s table he could ________ the missing elements and each elements characteristics. 4 He predicted ______________correctly years before it was identified.

18. Writing Symbols MGorMgormg All symbols are written with a ____________ letter first. 4 If it is a one letter symbol it is capital 4 If it is a two letter symbol first is capital, second letter is lowercase MgONiCoNCsKBe

19. 2.2 Classes of Elements

21. Metals, Nonmetals, and Metalloids 4 Metals want to ______ valence electrons to fill their outermost energy level. This makes them __________ ions. 4 Nonmetals want to _______ valence electrons to fill their outermost energy level. This would make them _________ ions. MetalsNonmetals

22. Metals, Nonmetals, and Metalloids 4 Metalloids could either ______ or ______ electrons to fill their outermost energy level, therefore they tend ________ to be reactive. 4 Some Nonmetals have full outer energy levels and don’t want to give or receive any electrons so they don’t bond. These are called _____________

23. 3.1 Introduction to Chemical Bonds

24. Chemical bonding is the _______ of atoms to form new substances.

25. A valence electron is an electron in the ____________ energy level of an atom, which causes bonding. n p + 1st 3rd2nd 8 e 2 e valence

26. Elements in the same have the same electron number Elements in the same have the same number of energy _________

27. How do we determine Valence Electrons? Take out your structure grid and fill in the shortcut. 1. Find out group # of atom 2. If the group # is 1 or 2, that is the # of valence electrons, if the group # is greater than 12, you must subtract 10. The remaining # is the # of valence electrons.

28. Electron Dot Diagrams 4 A shortcut to draw atomic structures to see valence electrons. 4 Determine the number of valence electrons for an atom. 4 Write out symbol and draw electrons around symbol in configuration below X 15 6 2 37 4 8 X is the symbol. The rectangular box around the symbol is imaginary Place the valence electron around the symbol in the order seen on the left.

29. Counting Atoms (Rule 1) The number of atoms is written as a subscript AlCl 3 = ____________________________ Na 2 O = ____________________________ 1 is never written and is always understood

30. Counting Atoms (Rule 2) 3Li 2 O = ________________________. A number in front must be distributed through the entire compound. 6C 6 H 12 O 6 =

31. Counting Atoms (Rule 3) Pb(OH) 2 = __________________________ In parenthesis the subscript indicates the number of everything in the parenthesis and therefore must be distributed through. Al 2 (SO 4 ) 3 =

32. Counting Atoms (Rule 4) NH 4 NO 3 = ______________________________ If an element appears in two different places it must be added together and totaled. 2H 3 C 6 H 5 O 7 =

33. Practice 1. KOH 2. Hg(OH) 3. FeCl 3 4. NH 4 OH 5. CCl 4 6. Pb(OH) 2 7. NH 4 NO 3 8. Al 2 (SO 4 ) 3 9. Li 2 O 10. Na 2 SO 3

34. 11. 2H 2 12. 2NaCl 13. 6H 2 O 14. 3C 2 H 6 15. 2Al 2 O 3 16. 4AgNO 3 17. 5C 2 H 6 O 18. H 3 C 6 H 5 O 7 19. 2Ga(OH) 3 20. 3Pb(NO 3 ) 2

35. 3.2 Ionic Bonds

36. In order to achieve stability: Every ____ wants a _____ outer shell !! Shell = Energy Level Full outer shell = ___ electrons (except for H and He which have a full outer shell of 2, first energy level)

37. Atoms with a ________ outermost energy level are the most _______ (not likely to bond). The Noble gases are inert, meaning they won’t bond, because they have full outer energy levels. He, Ne, Ar, Kr, Xe, Rn

38. Bonding that involves a ________ of _______ is called ionic bonding. Both _____ wish to ________ their outermost energy levels one atom (_________) gives up electrons (_______ ion) & one atom (_________) takes in electrons (_______ ion)

40. P All Ionic bonds are bonding _________ to _________. P Metals always _______ e- P Nonmetals always _______ e- P Metals are always written ___. P The final bonded compound must be stable with a total charge of ________

41. 3.3 Covalent Bonds

42. Covalent Bonds ___________ of electrons. All atoms still want a _____ outer shell. All formed molecules must have a stable charge ________. __________ to ___________. Hydrogen is a nonmetal and always forms a covalent bond because hydrogen has no shell underneath to give away its only electron.

43. The combination of atoms formed by a covalent bond is called a ____________.

44. Polyatomic ions (Found in parenthesis)

45. Diatomic Atoms: Those that easily bond with _________________. There are only _________ atoms that bond to themselves:_____________________ They make: ___________________________

46. 3.4 Metallic Bonds

48.  Metallic bonds: ________ & Metal Tips For Predicting Bond Type  Ionic bonds: _______ & Nonmetal Or _______ & Polyatomic ion  Covalent bonds: ________ & Nonmetal

49. Predict what type of bond will form between the elements listed below: 1. Hydrogen & Oxygen11. Sodium & Hydroxide (OH) 2. Potassium & Chlorine12. Potassium & Iodine 3. Zinc & Chlorine13. Hydrogen & Hydrogen 4. Chlorine & Chlorine14. Iron & Hydroxide (OH) 5. Iron & Iron15. Manganese & Oxygen 6. Sulfur & Oxygen16. Chlorine & Chlorine 7. Oxygen & Oxygen17. Calcium & Sulfate (SO 4 ) 8. Iron & Oxygen18. Copper & Copper 9. Gold & Gold19. Nitrogen & Oxygen 10. Carbon & Oxygen20. Sodium & Carbonate (CO 3 )

50. 4.1 Introduction to Chemical Reactions

51. Physical vs. Chemical Rxn Physical Rxn: Start with ___ and end with ____ Chemical Rxn: Start with ____ and end with ____ 1. ____________________ 2. ____________________ 3. ____________________

52. Signs of Chemical Rxn 1. _______________________________________ 2. _______________________________________ 3. _______________________________________ 4. Formation of a __________________ (solid produced in solution) *These are indicators, but not a guarantee of a chemical rxn.

53. Nitric Acid and Penny Was a gas formed? Was Energy given off? Was there a color change? Did a precipitate form? H 2 SO 4 and Sugar (C 6 H 12 O 6 )

54. 4.2 Chemical Equations

55. Chemical Equations 6 CO 2 + 6 H 2 O C 6 H 12 O 6 + 6 O 2 ___________ ____________ (What you start with.)(What you finish with.) 1. ______________ 2. __________________ 3. __________________

56. Writing Equations 4 Reactants: _____ ______________. KI and Pb(NO 3 ) 2 4 Products: ______ ______________ KNO 3 and PbI 2

57. KI and Pb(NO 3 ) 2 1. Was a gas formed? 2. Was Energy given off? 3. Was there a color change? 4. Did a precipitate form?

58. Law of Conservation of Mass/Matter Mass/Matter ________ be _______ or _________ it can only be transformed. AKA=What you ______ with you must ________ with.

59. Balanced??? KI + Pb(NO 3 ) 2 KNO 3 + PbI 2 A number needs to go in front to change the number of molecules. We can bring in or create as many molecules as we want. WE CANNOT BREAK ANY BONDS!

60. Balancing Rules 1. Count the atoms on each side of reaction (If a polyatomic group leave as one group. i.e. (CO 3 ), count as one CO 3 ) 2. Pick one atom, or polyatomic group, and go to the side with fewer atoms. 3. Place a coefficient in front of the molecule containing that atom, or polyatomic group, to balance number of atoms. 4. Re-calculate 5. If not balanced, repeat step 2 for all atoms.

61. Balanced??? KI + Pb(NO 3 ) 2 KNO 3 + PbI 2

62. 4.3 Types of Chemical Reactions

65. 4.4 Chemical Reactions and Energy

66. Law of Conservation of Energy 4 Energy cannot be ________ or ____________, but can be changed from one _______ to another.

67. Exo vs. Endo Exothermic Reaction: Energy is ________ by __________ bonds Example: Sulfuric Acid and Sugar releases heat. Endothermic Reaction: Energy is ________ by _________ new bonds Example: Chemical Ice packs, Photosynthesis (H 2 O and CO 2 becomes O 2 and Sugar C 6 H 12 O 6 ) All Chemical Energy is ________ in chemical bonds

68. Activation Energy ____________ amount of _________ required to ________ a reaction Could be ______, another _______, _________, _________, etc

69. Energy Graphs ExothermicEndothermic Products have ________ energy in bonds than reactants Products have ___________ energy in bonds than reactants

70. Rates (Speed) of Reaction Affected by: 1. _____________ 2. _____________ 3. _____________ 4. _____________ 5. _____________

71. _______ Temperature = ______ reaction rate Temperature = Amount of _______ energy _______ Concentration = ______ reaction rate Concentration = Amount of ___________ in reaction Surface Area = Amount of _______ exposed to __________ _______ Surface Area = _______ reaction rate

72. Catalyst and Inhibitors = Separate __________ that are added to a reaction __________ Catalyst = __________ reaction rate __________ Inhibitors = __________ reaction rate

73. Catalysts _________ activation energy to start reaction _________.