1. Outline for Review 1) The Atom (Nuclear, Electron Config)The Atom 2) Matter (Phases, Types, Changes)Matter 3) Bonding (Periodic Table, Ionic, Covalent)Bonding 4) Compounds (Formulas, Reactions, IMAF’s)Compounds 5) Math of Chemistry (Formula Mass, Gas Laws, Neutralization, etc.) Math of Chemistry 8) Oxidation and Reduction (Half Reactions, Cells, etc.)Oxidation and Reduction

2. The Atom 1) NucleonsNucleons 2) IsotopesIsotopes 3) Natural RadioactivityNatural Radioactivity 4) Half-LifeHalf-Life 5) Nuclear PowerNuclear Power 6) Electron ConfiguationElectron Configuation 7) Development of the Atomic ModelDevelopment of the Atomic Model

3. (c) 2006, Mark Rosengarten Nucleons 4 Protons: +1 each, determines identity of element, mass of 1 amu, determined using atomic number, nuclear charge 4 Neutrons: no charge, determines identity of isotope of an element, 1 amu, determined using mass number - atomic number (amu = atomic mass unit) 4 32 16 S and 33 16 S are both isotopes of S 4 S-32 has 16 protons and 16 neutrons 4 S-33 has 16 protons and 17 neutrons 4 All atoms of S have a nuclear charge of +16 due to the 16 protons.

4. Isotopes 4 Atoms of the same element MUST contain the same number of protons. 4 Atoms of the same element can vary in their numbers of neutrons, therefore many different atomic masses can exist for any one element. These are called isotopes. 4 The atomic mass on the Periodic Table is the weight- average atomic mass, taking into account the different isotope masses and their relative abundance.weight- average atomic mass 4 Rounding off the atomic mass on the Periodic Table will tell you what the most common isotope of that element is.common isotope of that element

5. Weight-Average Atomic Mass 4 WAM = ((% A of A/100) X Mass of A) + ((% A of B/100) X Mass of B) + … 4 What is the WAM of an element if its isotope masses and abundances are: –X-200: Mass = 200.0 amu, % abundance = 20.0 % –X-204: Mass = 204.0 amu, % abundance = 80.0% –amu = atomic mass unit ( 1.66 × 10 -27 kilograms/amu)

6. (c) 2006, Mark Rosengarten Most Common Isotope 4 The weight-average atomic mass of Zinc is 65.39 amu. What is the most common isotope of Zinc? Zn-65! 4 What are the most common isotopes of: –CoAg –SPb 4 FACT: one atomic mass unit (1.66 × 10 -27 kilograms) is defined as 1/12 of the mass of an atom of C-12. 4 This method doesn’t always work, but it usually does. Use it for the Regents exam.

7. Electron Configuration 4 Basic Configuration Basic Configuration 4 Valence Electrons Valence Electrons 4 Electron-Dot (Lewis Dot) Diagrams Electron-Dot (Lewis Dot) Diagrams 4 Excited vs. Ground State Excited vs. Ground State 4 What is Light? What is Light?

8. Basic Configuration 4 The number of electrons is determined from the atomic number. 4 Look up the basic configuration below the atomic number on the periodic table. (PEL: principal energy level = shell) 4 He: 2 (2 e - in the 1st PEL) 4 Na: 2-8-1 (2 e - in the 1st PEL, 8 in the 2nd and 1 in the 3rd) 4 Br: 2-8-18-7 (2 e - in the 1st PEL, 8 in the 2nd, 18 in the 3rd and 7 in the 4th)

9. Valence Electrons 4 The valence electrons are responsible for all chemical bonding. 4 The valence electrons are the electrons in the outermost PEL (shell). 4 He: 2 (2 valence electrons) 4 Na: 2-8-1 (1 valence electron) 4 Br: 2-8-18-7 (7 valence electrons) 4 The maximum number of valence electrons an atom can have is EIGHT, called a STABLE OCTET.

10. Electron-Dot Diagrams 4 The number of dots equals the number of valence electrons. 4 The number of unpaired valence electrons in a nonmetal tells you how many covalent bonds that atom can form with other nonmetals or how many electrons it wants to gain from metals to form an ion. 4 The number of valence electrons in a metal tells you how many electrons the metal will lose to nonmetals to form an ion. Caution: May not work with transition metals. 4 EXAMPLE DOT DIAGRAMS EXAMPLE DOT DIAGRAMS

11. (c) 2006, Mark Rosengarten Example Dot Diagrams Carbon can also have this dot diagram, which it has when it bonds to 4 other atoms.

12. Excited vs. Ground State 4 Configurations on the Periodic Table are ground state configurations. 4 If electrons are given energy, they rise to higher energy levels (excited state). 4 If the total number of electrons matches in the configuration, but the configuration doesn’t match, the atom is in the excited state. 4 Na (ground, on table): 2-8-1 4 Example of excited states: 2-7-2, 2-8-0-1, 2-6-3

13. What Is Light? 4 Light is formed when electrons drop from the excited state to the ground state. 4 The lines on a bright-line spectrum come from specific energy level drops and are unique to each element. 4 EXAMPLE SPECTRUM EXAMPLE SPECTRUM

14. This is the bright-line spectrum of hydrogen. The top numbers represent the PEL (shell) change that produces the light with that color and the bottom number is the wavelength of the light (in nanometers, or 10 -9 m). No other element has the same bright-line spectrum as hydrogen, so these spectra can be used to identify elements or mixtures of elements.

15. Development of the Atomic Model 4 Thompson Model Thompson Model 4 Rutherford Gold Foil Experiment and Model Rutherford Gold Foil Experiment and Model 4 Bohr Model Bohr Model 4 Quantum-Mechanical Model Quantum-Mechanical Model

16. (c) 2006, Mark Rosengarten Thompson Model 4 The atom is a positively charged diffuse mass with negatively charged electrons stuck in it.

17. (c) 2006, Mark Rosengarten Rutherford Model 4 The atom is made of a small, dense, positively charged nucleus with electrons at a distance, the vast majority of the volume of the atom is empty space. Alpha particles shot at a thin sheet of gold foil: most go through (empty space). Some deflect or bounce off (small + charged nucleus).

18. (c) 2006, Mark Rosengarten Bohr Model 4 Electrons orbit around the nucleus in energy levels (shells). Atomic bright-line spectra was the clue.

19. (c) 2006, Mark Rosengarten Quantum-Mechanical Model 4 Electron energy levels are wave functions. 4 Electrons are found in orbitals, regions of space where an electron is most likely to be found. 4 You can’t know both where the electron is and where it is going at the same time. 4 Electrons buzz around the nucleus like gnats buzzing around your head.

20. Bonding 1) The Periodic TableThe Periodic Table 2) Ions Ions 3) Ionic BondingIonic Bonding 4) Covalent BondingCovalent Bonding 5) Metallic BondingMetallic Bonding

21. (c The Periodic Table 4 Metals Metals 4 Nonmetals Nonmetals 4 Metalloids Metalloids 4 Chemistry of Groups Chemistry of Groups 4 Electronegativity Electronegativity 4 Ionization Energy Ionization Energy

22. Metals 4 Have luster, are malleable and ductile, good conductors of heat and electricity 4 Lose electrons to nonmetal atoms to form positively charged ions in ionic bonds Lose electrons to nonmetal atoms to form positively charged ionsionic bonds 4 Large atomic radii compared to nonmetal atoms 4 Low electronegativity and ionization energyelectronegativityionization energy 4 Left side of the periodic table (except H)

23. Nonmetals 4 Are dull and brittle, poor conductors 4 Gain electrons from metal atoms to form negatively charged ions in ionic bonds Gain electrons from metal atoms to form negatively charged ions 4 Share unpaired valence electrons with other nonmetal atoms to form covalent bonds and moleculescovalent bonds 4 Small atomic radii compared to metal atoms 4 High electronegativity and ionization energyelectronegativityionization energy 4 Right side of the periodic table (except Group 18)

24. Metalloids 4 Found lying on the jagged line between metals and nonmetals flatly touching the line (except Al and Po). 4 Share properties of metals and nonmetals (Si is shiny like a metal, brittle like a nonmetal and is a semiconductor).

25. Chemistry of Groups 4 Group 1: Alkali Metals Group 1: Alkali Metals 4 Group 2: Alkaline Earth Metals Group 2: Alkaline Earth Metals 4 Groups 3-11: Transition Elements Groups 3-11: Transition Elements 4 Group 17: Halogens Group 17: Halogens 4 Group 18: Noble Gases Group 18: Noble Gases 4 Diatomic Molecules Diatomic Molecules

26. Group 1: Alkali Metals 4 Most active metals, only found in compounds in nature 4 React violently with water to form hydrogen gas and a strong base: 2 Na (s) + H 2 O (l)  2 NaOH (aq) + H 2 (g) 4 1 valence electron 4 Form +1 ion by losing that valence electron 4 Form oxides like Na 2 O, Li 2 O, K 2 O

27. Group 2: Alkaline Earth Metals 4 Very active metals, only found in compounds in nature 4 React strongly with water to form hydrogen gas and a base: –Ca (s) + 2 H 2 O (l)  Ca(OH) 2 (aq) + H 2 (g) 4 2 valence electrons 4 Form +2 ion by losing those valence electrons 4 Form oxides like CaO, MgO, BaO

28. (c) 2006, Mark Rosengarten Groups 3-11: Transition Metals 4 Many can form different possible charges of ions 4 If there is more than one ion listed, give the charge as a Roman numeral after the name 4 Cu +1 = copper (I) Cu +2 = copper (II) 4 Compounds containing these metals can be colored.

29. (c Group 17: Halogens 4 Most reactive nonmetals 4 React violently with metal atoms to form halide compounds: 2 Na + Cl 2  2 NaCl 4 Only found in compounds in nature 4 Have 7 valence electrons 4 Gain 1 valence electron from a metal to form -1 ions 4 Share 1 valence electron with another nonmetal atom to form one covalent bond.

30. Group 18: Noble Gases 4 Are completely nonreactive since they have eight valence electrons, making a stable octet. 4 Kr and Xe can be forced, in the laboratory, to give up some valence electrons to react with fluorine. 4 Since noble gases do not naturally bond to any other elements, one atom of noble gas is considered to be a molecule of noble gas. This is called a monatomic molecule. Ne represents an atom of Ne and a molecule of Ne.

31. Diatomic Molecules 4 Br, I, N, Cl, H, O and F are so reactive that they exist in a more chemically stable state when they covalently bond with another atom of their own element to make two-atom, or diatomic molecules. 4 Br 2, I 2, N 2, Cl 2, H 2, O 2 and F 2 4 The decomposition of water : 2 H 2 O  2 H 2 + O 2

32. (c) 2006, Mark Rosengarten Electronegativity 4 An atom’s attraction to electrons in a chemical bond. 4 F has the highest, at 4.0 4 Fr has the lowest, at 0.7 4 If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion. 4 If the two atoms have an END of less than 1.7, they will share their unpaired valence electrons…covalent bond!

33. Ionization Energy 4 The energy required to remove the most loosely held valence electron from an atom in the gas phase. 4 High electronegativity means high ionization energy because if an atom is more attracted to electrons, it will take more energy to remove those electrons. 4 Metals have low ionization energy. They lose electrons easily to form (+) charged ions. 4 Nonmetals have high ionization energy but high electronegativity. They gain electrons easily to form (-) charged ions when reacted with metals, or share unpaired valence electrons with other nonmetal atoms.

34. Ions 4 Ions are charged particles formed by the gain or loss of electrons. –Metals lose electrons (oxidation) to form (+) charged cations.Metals lose electrons –Nonmetals gain electrons (reduction) to form (-) charged anions.Nonmetals gain electrons 4 Atoms will gain or lose electrons in such a way that they end up with 8 valence electrons (stable octet). –The exceptions to this are H, Li, Be and B, which are not large enough to support 8 valence electrons. They must be satisfied with 2 (Li, Be, B) or 0 (H).

35. Metal Ions (Cations) 4 Na: 2-8-1 4 Na +1 : 2-8 4 Ca: 2-8-8-2 4 Ca +2 : 2-8-8 4 Al: 2-8-3 4 Al +3 : 2-8 Note that when the atom loses its valence electron, the next lower PEL becomes the valence PEL. Notice how the dot diagrams for metal ions lack dots! Place brackets around the element symbol and put the charge on the upper right outside!

36. (c) 2006, Mark Rosengarten Nonmetal Ions (Anions) 4 F: 2-7 4 F -1 : 2-8 4 O: 2-6 4 O -2 : 2-8 4 N: 2-5 4 N -3 : 2-8 Note how the ions all have 8 valence electrons. Also note the gained electrons as red dots. Nonmetal ion dot diagrams show 8 dots, with brackets around the dot diagram and the charge of the ion written to the upper right side outside the brackets.

37. Ionic Bonding 4 If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion. 4 The oppositely charged ions attract to form the bond. It is a surface bond that can be broken by melting or dissolving in water. 4 Ionic bonding forms ionic crystal lattices, not molecules.

38. Example of Ionic Bonding

39. (c) 2006, Mark Rosengarten Covalent Bonding 4 If two nonmetal atoms have an END of 1.7 or less, they will share their unpaired valence electrons to form a covalent bond. 4 A particle made of covalently bonded nonmetal atoms is called a molecule. 4 If the END is between 0 and 0.4, the sharing of electrons is equal, so there are no charged ends. This is NONPOLAR covalent bonding.  If the END is between 0.5 and 1.7, the sharing of electrons is unequal. The atom with the higher EN will be  - and the one with the lower EN will be  + charged. This is a POLAR covalent bonding. (  means “partial”)

40. (c) 2006, Mark Rosengarten Examples of Covalent Bonding

41. Metallic Bonding 4 Metal atoms of the same element bond with each other by sharing valence electrons that they lose to each other. 4 This is a lot like an atomic game of “hot potato”, where metal kernals (the atom inside the valence electrons) sit in a crystal lattice, passing valence electrons back and forth between each other). 4 Since electrons can be forced to travel in a certain direction within the metal, metals are very good at conducting electricity in all phases.

42. Compounds 1) Types of CompoundsTypes of Compounds 2) Formula WritingFormula Writing 3) Formula NamingFormula Naming 4) Empirical FormulasEmpirical Formulas 5) Molecular FormulasMolecular Formulas 6) Types of Chemical ReactionsTypes of Chemical Reactions 7) Balancing Chemical ReactionsBalancing Chemical Reactions 8) Attractive ForcesAttractive Forces

43. (c) 2006, Mark Rosengarten Types of Compounds 4 Ionic: made of metal and nonmetal ions. Form an ionic crystal lattice when in the solid phase. Ions separate when melted or dissolved in water, allowing electrical conduction. Examples: NaCl, K 2 O, CaBr 2 Ionic 4 Molecular: made of nonmetal atoms bonded to form a distinct particle called a molecule. Bonds do not break upon melting or dissolving, so molecular substances do not conduct electricity. EXCEPTION: Acids [H + A - (aq)] ionize in water to form H 3 O + and A -, so they do conduct. Molecular 4 Network: made up of nonmetal atoms bonded in a seemingly endless matrix of covalent bonds with no distinguishable molecules. Very high m.p., don’t conduct. Network

44. (c) 2006, Mark Rosengarten Ionic Compounds

45. (c) 2006, Mark Rosengarten Molecular Compounds

46. Network Solids Network solids are made of nonmetal atoms covalently bonded together to form large crystal lattices. No individual molecules can be distinguished. Examples include C (diamond) and SiO 2 (quartz). Corundum (Al 2 O 3 ) also forms these, even though Al is considered a metal. Network solids are among the hardest materials known. They have extremely high melting points and do not conduct electricity.

47. Formula Writing 4 The charge of the (+) ion and the charge of the (-) ion must cancel out to make the formula. Use subscripts to indicate how many atoms of each element there are in the compound, no subscript if there is only one atom of that element. 4 Na +1 and Cl -1 = NaCl 4 Ca +2 and Br -1 = CaBr 2 4 Al +3 and O -2 = Al 2 O 3 4 Zn +2 and PO 4 -3 = Zn 3 (PO 4 ) 2 4 Try these problems! Try these problems!

48. ( Formulas to Write 4 Ba +2 and N -3 4 NH 4 +1 and SO 4 -2 4 Li +1 and S -2 4 Cu +2 and NO 3 -1 4 Al +3 and CO 3 -2 4 Fe +3 and Cl -1 4 Pb +4 and O -2 4 Pb +2 and O -2

49. (c) 2006, Mark Rosengarten Formula Naming 4 Compounds are named from the elements or polyatomic ions that form them. 4 KCl = potassium chloride 4 Na 2 SO 4 = sodium sulfate 4 (NH 4 ) 2 S = ammonium sulfide 4 AgNO 3 = silver nitrate 4 Notice all the metals listed here only have one charge listed? So what do you do if a metal has more than one charge listed? Take a peek! Take a peek!

50. (c) 2006, Mark Rosengarten The Stock System 4 CrCl 2 = chromium (II) chloride Try 4 CrCl 3 = chromium (III) chloride Co(NO 3 ) 2 and 4 CrCl 6 = chromium (VI) chloride Co(NO 3 ) 3 4 FeO = iron (II) oxideMnS = manganese (II) sulfide 4 Fe 2 O 3 = iron (III) oxideMnS 2 = manganese (IV) sulfide 4 The Roman numeral is the charge of the metal ion!

51. Empirical Formulas 4 Ionic formulas: represent the simplest whole number mole ratio of elements in a compound. 4 Ca 3 N 2 means a 3:2 ratio of Ca ions to N ions in the compound. 4 Many molecular formulas can be simplified to empirical formulas –Ethane (C 2 H 6 ) can be simplified to CH 3. This is the empirical formula…the ratio of C to H in the molecule. 4 All ionic compounds have empirical formulas.

52. (c) 2006, Mark Rosengarten Formula Mass 4 Gram Formula Mass = sum of atomic masses of all elements in the compound 4 Round given atomic masses to the nearest tenth 4 H 2 O: (2 X 1.0) + (1 X 16.0) = 18.0 grams/mole 4 Na 2 SO 4 : (2 X 23.0)+(1 X 32.1)+(4 X 16.0) = 142.1 g/mole 4 Now you try: –BaBr 2 –CaSO 4 –Al 2 (CO 3 ) 3

53. (c) 2006, Mark Rosengarten Percent Composition The mass of part is the number of atoms of that element in the compound. The mass of whole is the formula mass of the compound. Don’t forget to take atomic mass to the nearest tenth! This is a problem for you to try.This is a problem for you to try

54. Practice Percent Composition Problem 4 What is the percent by mass of each element in Li 2 SO 4 ?

55. Mole Problems 4 Grams Moles Grams Moles 4 Molecular Formula Molecular Formula 4 Stoichiometry Stoichiometry

56. Grams Moles 4 How many grams will 3.00 moles of NaOH (40.0 g/mol) weigh? 4 3.00 moles X 40.0 g/mol = 120. g 4 How many moles of NaOH (40.0 g/mol) are represented by 10.0 grams? 4 (10.0 g) / (40.0 g/mol) = 0.250 mol

57. Molecular Formula 4 Molecular Formula = (Molecular Mass/Empirical Mass) X Empirical Formula 4 What is the molecular formula of a compound with an empirical formula of CH 2 and a molecular mass of 70.0 grams/mole? 4 1) Find the Empirical Formula Mass: CH 2 = 14.0 4 2) Divide the MM/EM: 70.0/14.0 = 5 4 3) Multiply the molecular formula by the result: 5 (CH 2 ) = C 5 H 10

58. (c) 2006, Mark Rosengarten How many Sig Figs? 4 Start counting sig figs at the first non-zero. 4 All digits except place-holding zeroes are sig figs. Measurement# of Sig Figs 234 cm3 67000 cm2 _ 45000 cm 4 560. cm3 560.00 cm5 Measurement# of Sig Figs 0.115 cm3 0.00034 cm2 0.00304 cm3 0.0560 cm3 0.00070700 cm5

59. What Precision? 4 A number’s precision is determined by the furthest (smallest) place the number is recorded to. 4 6000 mL : thousands place 4 6000. mL : ones place 4 6000.0 mL : tenths place 4 5.30 mL : hundredths place 4 8.7 mL : tenths place 4 23.740 mL : thousandths place

60. Rounding with addition and subtraction 4 Answers are rounded to the least precise place.

61. Rounding with multiplication and division 4 Answers are rounded to the fewest number of significant figures.

62. Metric Conversions 4 Determine how many powers of ten difference there are between the two units (no prefix = 10 0 ) and create a conversion factor. Multiply or divide the given by the conversion factor. How many kg are in 38.2 cg? (38.2 cg) /(100000 cg/kg) = 0.000382 km How many mL in 0.988 dL? (0.988 dg) X (100 mL/dL) = 98.8 mL