1. Chapter 6.1 Periodicity Ox Quiz#1, Ox Quiz#2 now! tomorrow! By now, you know that the Periodic Table is organized by increasing atomic #, by orbitals, and by the number of valence electrons. But… There’s still more! The organization of the periodic table can tell you several properties of elements, if you know what you’re looking for.

2. Periodic Law: When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties Elements within the same column, or group, have similar properties. Elements within the same row, or period, change from left to right across the PT, but repeat in the next period.

4. Metals From everyday life, you should be familiar with the word “metal” and should know of several metals… But just how many of the elements on the periodic table are metals?

5. Metals Roughly 80% of elements are metals.

6. Generally, elements can be classified as being metal, nonmetal, or metalloid. Each can be broken down into sub-categories, which can be seen in the periodic table in your book.

7. Metals: - Good conductors of heat/electricity. - Ductile (can be made into a wire) - Most are malleable (can be molded) - Tend to be solid at room temp. Who am I? 13p 14n

8. Metals: - Good conductors of heat/electricity. - Ductile (can be made into a wire) - Most are malleable (can be molded) - Tend to be solid at room temp. Al Outer energy levels overlap 13p 14n 13p 14n

9. Metals: - Good conductors of heat/electricity. - Ductile (can be made into a wire) - Most are malleable (can be molded) - Tend to be solid at room temp. Nonmetals: - Mostly poor conductors of heat/electricity. - Tend to be gaseous at room temp. - Not malleable forbidden gap 16p 16n 16p 16n

10. Metals: - Good conductors of heat/electricity. - Ductile (can be made into a wire) - Most are malleable (can be molded) - Tend to be solid at room temp. Nonmetals: - Mostly poor conductors of heat/electricity. - Tend to be gaseous at room temp. - Not malleable Metalloids: - Have properties of both metals and nonmetals. - Behavior often controlled by changing conditions

11. Sub-categories of metals and nonmetals: Ch6 HW#1

12. Chapter 6 HW#1 1. List your favorite element from the halogen family, then draw its dot diagram. 2. List your favorite element from the alkali metals, then draw its dot diagram. 3. Make a statement about the reactivity of the alkali metals and the halogens.

13. 4. Who am I? I am the only element that falls along the metal/nonmetal dividing line that is strictly a metal. Engineers like to make lightweight structures out of me and my alloys. I carry 3 valence electrons… 5. Who am I? I can sometimes be classified as a metal or a metalloid, and chemically, my valence electron configuration is similar to oxygen, except I got a whole bunch more energy levels. Did you know that Rutherford used one of my radioisotopes in his famous Gold Foil Experiment.

14. Ch6.2 Ions and Atomic SizeOx Quiz#2 today! Ox Quiz#3 tom! By looking at the periodic table, how does the atomic mass of a halogen compare to that of an alkali metal in the same period?

15. - Answer: It is greater You should already have known this from previous chapters, but it does lead to some interesting discussion… Now, how do you think the overall size of the atom compares for atoms of the same two elements?

16. -Answer: It is smaller! It may seem strange, but despite the fact that atomic mass increases from left to right across the periodic table, the atomic size, or atomic radius, actually decreases. This is because the increased number of protons (positive charges) in the nucleus creates a greater electromagnetic pull on the electrons (negative charges) orbiting around the outside.

17. Atomic radius does increase as you move down a group due to the fact that there are more energy levels.

18. Ions We know that atoms have a neutral charge because the charges of protons and electrons cancel each other out. But, when atoms react with their surroundings, they tend to either gain or lose an electron, which leaves it with either a positive or negative charge. An atom that has gone through this kind of reaction is called an ion.

19. Cation: an atom that has lost an electron and is left with an overall positive charge. Al: Anion: an atom that has gained an electron and is left with an overall negative charge. 13p 14n

20. 13p 14n 13p 14n 16p 16n Cation: an atom that has lost an electron and is left with an overall positive charge. Al: Al 3+ 10 electrons 13 protons = Anion: an atom that has gained an electron and is left with an overall negative charge. S:

21. Cation: an atom that has lost an electron and is left with an overall positive charge. Al: Al 3+ 10 electrons 13 protons = Anion: an atom that has gained an electron and is left with an overall negative charge. S: S 2– 18 electrons 16 protons = 13p 14n 13p 14n 16p 16n 16p 16n

22. Cation: Lost an electron Positive charge Denoted by a + after the symbol (Na +, Mg 2+, etc.) Tends to occur in metals. Decreases the atomic radius of the atom. Anion: Gained an electron Negative charge Denoted by a – after the symbol (F –, S 2–, etc.) Tends to occur in nonmetals. Increases the atomic radius of the atom.

23. Cation: Lost an electron Positive charge Denoted by a + after the symbol (Na +, Mg 2+, etc.) Tends to occur in metals. Decreases the atomic radius of the atom. Anion: Gained an electron Negative charge Denoted by a - after the symbol (F -, S 2-, etc.) Tends to occur in nonmetals. Increases the atomic radius of the atom. Cations and Anions are both results of an ionic reaction… For example, in sodium chloride (table salt): lose gain cationanion

24. See textbook or online for ionization/atomic radius chart. Ch6 HW#2 + Lab6.1 Pre-lab Questions

25. Chapter 6 HW#2 1. Explain why fluorine has a smaller atomic radius than both oxygen and chlorine. 2. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum and sodium. Does your arrangement demonstrate a periodic or group trend? 3. Which particle has the largest radius in each atom /ion pair? a. Na, Na + b. S, S -2 c. I, I -1 d. Al, Al +3

26. 4. Chlorine, selenium, and bromine are located near each other on the period table. Which of these elements is the smallest atom? 5. Phosphorus, sulfur, and selenium are located near each other on the periodic table. Which of these elements is the largest atom? 6. Scandium, yttrium, and lanthanum are located near each other in the periodic table. Which of these elements is the largest atom? 7. Which of the following atoms is smallest: vanadium, chromium or tungsten? 8. Which of the following atoms is smallest: nitrogen, phosphorus or arsenic?

27. Ch6.3 Ionization Energy Ox Quiz#3 Today! Ox Quiz#4 Tomorrow! We know that, as electrons absorb energy, they can move up to higher energy levels. If they absorb enough energy, they can leave the atom entirely. Ionization Energy: The energy required to remove an electron from an atom.

28. The first ionization energy is the energy required to remove one electron, leaving a +1 charge behind. The second ionization energy is the energy to remove a second electron (once the first has already been removed), leaving a +2 charge behind. And so on, and so on… See text or online for a list of first, second and third ionization energies for the first 20 elements.

29. In general, the ionization energy increases from left to right across a period and from bottom to top in a group on the PT. This figure shows the trend of the first ionization energies of the representative elements (those in the s and p blocks). (Table 1 for HW)

30. Graphically, we can also see the trends in ionization energies. Notice that the ionization energies of the alkali metals (Li, Na, K, etc.) decrease with atomic number, as well as the noble gases (He, Ne, Ar, etc.), just as they would in their respective rows. (Table 2 for HW)

31. Ex1) Lithium has a very low 1 st ionization energy. But there is a tremendous jump to its 2 nd IE. Explain. Ch6 HW#3 1 – 6 3p 4n (Table 3 for HW)

32. Chapter 6 HW#3 1. Distinguish between the first and second ionization energies of an atom. 2. Refer to Table 1, indicate which element in each of the following pairs has the greatest first ionization energy. a. lithium, boron b. magnesium, strontium c. cesium, aluminum 3. Chlorine, selenium, and bromine are located near each other on the period table. Which of these elements has the highest ionization energy? 4. Phosphorus, sulfur, and selenium are located near each other on the periodic table. Which of these elements has the highest ionization energy?

33. 5. Refer to table 2, there is a mini-spike in the graph just before atomic # 50. What element might this be? Justify your choice.

34. 6. Refer to table3, why does carbon have such a significant jump in ionization energy from I 4 to I 5 ?

35. Ch6.4 Electronegativity - the ability of an atom to attract electrons when the atom is in a compound. This property also tends to increase from left to right, and from bottom to top. (Noble gases are excluded from the effects of electronegativity because they have full valence shells and rarely form compounds)… (EN table, refer for HW)

36. Atoms desire to have either empty or full valence shells … In other words, they like to have either 0 or 8 electrons in their outer shells. Elements with low numbers of valence electrons (alkali metals, alkaline metals, etc) Tend to have low ionization energies/low electronegativity (easier to achieve 0 by taking electrons away) Elements with high numbers of valence electrons (halogens, etc) Tend to have high ionization energies/high electrogenativity (easier to achieve 8 by gaining electrons) (Use this electronegativity table in for your homework.)

37. Covalent bonding: Occurs when there is electron sharing between nonmetals and/or metaloids (high electronegativity). Ionic Bonding: Occurs when there is electron transfer from a metal (low ionization energy) to a nonmetal (high electrogenativity), forming both + and – ions.

38. Covalent bonding: Occurs when there is electron sharing between nonmetals and/or metaloids (high electronegativity). Ionic Bonding: Occurs when there is electron transfer from a metal (low ionization energy) to a nonmetal (high electrogenativity), forming both + and – ions. Electronegativity Difference – subtract the #’s from the EN Table. 1.67 is the magic #. Ionic Bonds > 1.67Covalent Bonds < 1.67 Ex1) Decide if the following elements bond covalently or ionic: a. Na and Cl: b. Si and O: c. O and O: Ch6 HW#4

39. Lab 6.1 Qualitative Analysis In the laboratory, an important activity of the chemist is to determine what elements are present in a given sample. Specific reagents are not available to identify each type of ion or atom. Therefore, an orderly plan is required to separate the ions and correctly identify each. In part A of this experiment, you will follow an orderly plan for identifying the ions of the halogen family. You will then attempt to identify an unknown substance that contains one of the ions from the halogen family.

40. Ch6 HW#4

41. 1. Which of these elements has a greater electronegativity? a. Na or K b. Mg or Pc. Fr or F 2. Use the Electronegativity Table to decide whether each pair of atoms will form an ionic or covalent bond: a. Na and Brb. Ca and Clc. C and O d. Al and Oe. Mg and S

42. Ch9.1 Naming Ions Review: Elements can be classified as metals, nonmetals, and metalloids. Metals and nonmetals tend to have vastly different properties, while metalloids can have properties similar to both, depending upon conditions. (For now, let’s just focus on metals and nonmetals)

43. Ions can form in two ways: Cation: Formed when a metal gives up electrons so that its valence shell is empty. (Positively Charged)

44. Ions can form in two ways: Cation: Formed when a metal gives up electrons so that its valence shell is empty. (Positively Charged) When metals in groups 1A, 2A, and 3A lose electrons, they form cations with positive charges equal to their group number.

45. Ions can form in two ways: Cation: Formed when a metal gives up electrons so that its valence shell is empty. (Positively Charged) When metals in groups 1A, 2A, and 3A lose electrons, they form cations with positive charges equal to their group number. Exs: K , Be , B  These ions are named the same as their neutral atom.

46. Ions can form in two ways: Anion: Formed when a nonmetal gains electrons so that its valence shell is full. (Negatively Charged) The charge of any ion of group A nonmetal is determined by subtracting 8 from the group number. In other words, the number of columns you must count over to reach an outer shell with 8 valence electrons.

47. Ions can form in two ways: Anion: Formed when a nonmetal gains electrons so that its valence shell is full. (Negatively Charged) The charge of any ion of group A nonmetal is determined by subtracting 8 from the group number. In other words, the number of columns you must count over to reach an outer shell with 8 valence electrons. Exs: N , Se , I  Named by taking the name of the neutral atom and and changing the ending to “-ide”

48. Elements in groups 4A and 8A generally do not form ions.

49. Transition metals (d and f blocks, or Group B elements) can each form multiple ions, depending on how many electrons it loses. Refer to your oxidation sheet for common transition metal ions formed. (We will talk about all those crazy (–) ions tomorrow.) Ch9 HW#1

50. Polyatomic Ions Polyatomic ions are ions that are composed of more than one atom. Combinations of atoms, which form their own ions, can act as one single ion. The most common polyatomic ion from a group is named with “-ate” Exs: SO 4 –2 NO 3 –

51. Most polyatomic ions end in “-ate” or “-ite” Common ions with both –ite and –ate suffixes for their names… Can you notice any patterns by comparing them? -ite:-ate: SO 3 –2 SO 4 –2 NO 2 – NO 3 – ClO 2 – ClO 3 – (Leave 1 line of space here…) Other endings include “-ium” (like ammonium, NH 3 1+ ) and the occasional “-ide” (like cyanide, CN – and hydroxide, OH – )

52. Common ions with both –ite and –ate suffixes for their names… -ite:-ate: SO 3 –2 SO 4 –2 NO 2 – NO 3 – ClO 2 – ClO 3 – The most common polyatomic ion from a group is named with “-ate” Each “–ite” contains one less oxygen than its corresponding “–ate” Can there be other ions in each group?

53. Common ions with both –ite and –ate suffixes for their names… -ite:-ate: SO 3 –2 SO 4 –2 NO 2 – NO 3 – ClO – ClO 2 – ClO 3 – ClO 4 – The most common polyatomic ion from a group is named with “-ate” Each “–ite” contains one less oxygen than its corresponding “–ate” Can there be other ions in each group? One less oxygen than “-ite” gets a prefix “hypo-” One more oxygen than “-ate” gets the prefix “per-”

54. Common ions with both –ite and –ate suffixes for their names… -ite:-ate: SO 3 –2 SO 4 –2 NO 2 – NO 3 – ClO – ClO 2 – ClO 3 – ClO 4 – hypochlorite chlorite chlorate perchlorate The most common polyatomic ion from a group is named with “-ate” Each “–ite” contains one less oxygen than its corresponding “–ate” Can there be other ions in each group? One less oxygen than “-ite” gets a prefix “hypo-” One more oxygen than “-ate” gets the prefix “per-” Ch9 HW#1

55. 1. Identify the following elements as metals, metalloids, or nonmetals. a. goldb. siliconc. manganesed. sulfure. barium 2. Describe two ways in which an ion forms from an atom. 3. Complete this table. 4. The most common polyatomic ion of bromine is bromate, BrO 3 –. Name and write the formulas of the bromine ions with 4 oxygens and also 1 oxygen. SymbolChange in electronsFormula of IonName of Ion a. Al3 electrons lost b. S S –2 c. Ca +2 d. Br1 electron gained e.P –3 f.Xe g.K

56. Ch9.2 Chemical Formulas and Naming Some substances have common names Ex: H 2 O: Just about all others we will name systematically: CO 2 : NaCl: Why the difference?

57. Ionic compounds - as a general rule: metal and a nonmetal - cation listed before the anion, anion changed to “-ide.” Ex1) Li and F Ca and S Na and I

58. Ex1) Li and F  LiF  Lithium Fluoride Ca and S  CaS  Calcium Sulfide Na and I  NaI  Sodium Iodide (Tomorrow we can learn how to write formulas.) - the total charge of the compound is zero. The ions always combine in ratios to cancel each other out.

59. molecular compounds - usually consist on nonmetals bonded together - share electrons in a covalent bond - usually neutral in charge (polyatomic ions are covalently bonded elements with an overall charge.) - still end in “-ide” - naming based on prefixes: Ex2) SO 2 N 2 O 4 BF 3 What would be another name for water? PrefixNumber Mono-*1 Di-2 Tri-3 Tetra-4 Penta-5 Hexa-6 Hepta-7 Octa-8 Nona-9 Deca-10

60. Here are some more properties of compounds formed from ionic and covalent bonds: Ch9 HW#2

61. 5. If ionic compounds are composed of charged particles (ions), why isn’t every ionic compound either positively or negatively charged? 6. The melting point of a compound is 1240°C. Do you predict this compound to be an ionic or molecular compound? 7. Identify the number and kinds of atoms present in a molecule of each compound: a. ascorbic acid (vitamin C) C 6 H 8 0 6 b. monosodium glutamate (MSG) C 5 H 8 O 4 NNa c. sucrose (table sugar) C 12 H 22 O 11 d. trinitrotoluene (TNT) C 7 H 5 N 3 O 6 e. ammonium nitrate (fertilizer) NH 4 NO 3 8. Would you expect the following pairs of atoms to combine chemically to give an ionic or a molecular compound? a. Li and Sb. O and Sc. Al and O d. F and Cle. I and Kf. H and N

62. 9. Name these molecular compounds: a. CS 2 b. Cl 2 O 7 c. N 2 O 5 d. CCl 4 e. CrCl 3 10. Name these ionic compounds: a. KBr b. CaCl 2 c. MgO d. LiCl e. Fe 2 O 3

63. Ch9.3 Chemical Formulas Binary compounds – made up of only two kinds of atoms. Ternary compounds – made up of more than two kinds of atoms. Ex1) Write chemical formulas for the following: sodium fluoride calcium sulfide magnesium chloride lithium nitride aluminum oxide

64. Ex2) Write chemical formulas for the following: sodium nitride sodium nitrate calcium carbide calcium carbonate magnesium sulfide magnesium sulfite magnesium sulfate There will be quizzes on formula writing ALL year! Ch9 HW#3Might as well learn it now! Do you HW!

65. Ch9 HW#3 11. What is the net ionic charge of every ionic compound? 12. Write formulas for compounds composed of these pairs of ions: a. Na +, O -2 b. Sn +4, S -2 c. K +, Cl - d. Mg +2, N -3

66. 13. Write the formulas for these compounds. a. potasium chloride b. potasium hypochlorite c. potasium chlorite d. potasium chlorate e. potasium perchlorate 14. Write names for these compounds. a. Al(OH) 3 b. NaI c. Na 2 CrO 4

67. Ch9.4 More Naming Transition metals can have more than one oxidation number. This is denoted with Roman Numerals Ex1) Write the chemical formula and name for: Sn 4+ and oxygen Fe 2+ and oxygen Fe 3+ and oxygen

68. Ex2) Write the chemical formula for: lead (I) carbonate copper (II) sulfate silver nitrate Ch9 HW#4

69. 15. How do you determine the charge of a transition metal cation from the formula of an ionic compound? 16. Write the names for these binary ionic compounds: a. AlF 3 b. FeS c. CaO d. Cu 2 Se e. BF 3 17. Write the formulas for: a. calcium carbonate b. lithium hypochlorite c. potassium phosphate d. tin (IV) dichromate

70. Ch4 Rev WS on Formulas and Naming (30 questions, we do 10, you do 2 for HW)

73. Ch9.5 Naming Acids and Bases (Discuss these in more detail later) Acids: A compound that contains one or more hydrogen atoms and produces hydrogen ions (H+) when dissolved in water. (The names of compounds can change when acids are present). The name of an acidic compound depends on the type of anion it is reacting with... For “–ide” anions, the prefix “hydro-” comes before the root word in the anion and the suffix “-ic” comes after. Ex: HCl For “-ite” anions, just the suffix “-ous” comes after the root word. Ex: HClO 2 For “-ate” anions, just add the suffix “-ic” after the root word. Ex: HClO 3 The word “acid” is then added after each of these names.

74. Bases: An ionic compound that produces hydroxide ions (OH-) when dissolved in water. Naming is similar to any ionic compound: Ex: NaOH Chemical formulas of Acids and Bases: calcium hydroxide nitric acid nitrous acid Ch9 HW#5

75. 18. The formula of a common acid will always contain what element? 19. Name these acids and bases. Please feel free to look on your oxidation sheet to find the name of the polyatomic ion. a. HFb. H 2 S c. H 2 SO 4 d. H 2 SO 3 e. KOH 20. Write the chemical formula of the following acids and bases. a. Hydrobromic acidb. Carbonic acid c. Phosphoric acidd. Magnesium hydroxide e. Aluminum hydroxide

76. Ch6 Review 1. Between lithium and potassium, which element has a. the largest atomic radius__________ b. the highest ionization energy__________ c. the highest electronegativity__________ 2. Between sodium and aluminum, which element has a. the largest atomic radius__________ b. the highest ionization energy__________ c. the highest electronegativity__________ 3. Between sulfur, arsenic and selenium, which element has a. the largest atomic radius__________ b. the highest ionization energy__________ c. the highest electronegativity__________ 4. Which has the larger size: a. calcium or its ion__________ b. aluminum or its ion__________ c. chlorine or its ion__________

77. 5. “Periodic” means that something repeats or goes through cycles. From the graph do you see any cyclical nature to the points? 6. List the elements that represent the peaks in your graph. How are these elements related on the periodic table? 7. List the elements that represent the lower peaks in your graph. How are these elements related on the periodic table? 8. List the elements that occur just before each peak on your graph. How are these elements related on the periodic table? 9. List the elements that occur just after each lower peak on your graph. How are these elements related on the periodic table?

78. 13p 14n 10. The following graph shows the successive ionization energies for period 3 on the periodic table. Referring to the diagram of aluminum, explain the first four ionization energies.

79. Ch9 Rev 1. Would you expect the following compounds to be ionic or molecular? a. COb. Na 2 Sc. C 2 H 6 d. N 2 O 3 2. Write the formulas for these compounds. a. sodium iodide b. stannous fluoride (Sn 2+ ) c. potassium sulfide d. calcium iodide

80. 3. Name these binary ionic compounds. a. ZnS b. KCl c. BaO d. CuBr 2

81. 4. Write formulas for the following binary molecular compounds. a. nitrogen trifluoride b. disulfur dichloride c. dinitrogen tetroxide d. phosphorus pentachloride 5. When are parentheses needed in a chemical formula? Use barium hydroxide to explain. 6. The formula of a common acid will always contain what element?

82. 7. Name these compounds. a. CaCO 3 b. KMnO 4 c. PbCrO 4 d. Mg 3 P 2 e. (NH 4 ) 2 SO 4 f. NaNO 2 8. Write the formulas for these compounds. a. tin(II) hydroxide b. barium fluoride c. calcium sulfide d. sodium phosphate