Organizing the Periodic Table

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2 Organizing the Periodic Table How did chemists begin to organize the known elements? Chemists used the properties of the elements to sort them into groups.

3 The Organizers JW Dobereiner grouped the elements into triads. Problem: Not all of the elements could be put into triads.

4 The Organizers Dmitri Mendeleev arranged the elements by increasing atomic mass. Problem: the chemical properties of the elements did not line up correctly.

5 Pg. 156 Figure 6.3

6 The Organizers Lother Meyer also created a periodic table similar to Mendeleev s, however Mendeleev published his table first and could explain his table better.

7 Development of Atomic Theory The structure of the atom was determined through years of experimentation. Once the structure of the atom was determined by Dalton, Thomson, Rutherford, Chadwick, Bohr, and others, Henry Moseley determined the atomic number for the elements.

8 Organization of Modern Periodic Table The elements are arranged in order of increasing atomic number.

9 Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

10 Three classes of elements: Metals Nonmetals Metalloids Across a period, the properties of elements become less metallic and more nonmetallic.

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12 Properties of Metals Solids (except for Hg) at room temperature Make up 80% of the periodic table Good conductors of heat and electricity Shiny Ductile (can be made into thin wire) Malleable (can be hammered into sheets)

13 Properties of Nonmetals Properties vary among the nonmetals. At room temperature, some are gases, some are solids, and 1 is a liquid. Poor conductors Brittle Not ductile Not malleable

14 Metalloids (Semi-metals) B, Si, Ge, As, Sb, Te, At Behave like metals under certain conditions. Behave like nonmetals under different conditions.

15 Which of these sets of elements have similar physical and chemical properties? a. oxygen, nitrogen, carbon, boron b. strontium, magnesium, calcium, beryllium c. nitrogen, neon, nickel, niobium

16 Identify each element as a metal, metalloid, or nonmetal. a. gold b. silicon c. sulfur d. barium

17 Name two elements that have properties similar to those of the element sodium.

18 What pattern is revealed when the elements are arranged in a periodic table in order of increasing atomic number?

19 Identify each property below as more characteristic of a metal or a nonmetal. a. a gas at room temperature b. brittle c. malleable d. poor conductor of electric current e. shiny

20 In general, how are metalloids different from metals and nonmetals?

21 Representative Elements Elements in columns 1, 2, 13, 14, 15, 16, 17, 18 OR Elements in columns 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A OR Elements in blocks s and p

22 Representative Elements Called REPRESENTATIVE elements because they display a wide range of chemical and physical properties.

23 Which of the following are symbols for representative elements: Na, Mg, Fe, Ni, Cl?

24 Transition Metals Elements in columns 3, 4, 5, 6, 7, 8, 9, 10, 11, 12 OR Elements in columns 3B, 4B, 5B, 6B, 7B, 8B, 1B, 2B OR Elements in the d block

25 Transition Metals Characterized by electrons filling the d orbitals When transition metal ions form compounds, the compounds often exhibit various colors. Iron oxide Cadmium sulfate Copper(II) chloride Red Orange Blue-green

26 Inner Transition Metals (Rare earth elements) Elements in the f block Consists of the Lanthanide (La) series and the Actinide (Ac) series.

27 Inner Transition Metals (Rare earth elements) Contains some elements that occur naturally and some that are synthetic. Some are radioactive, some are not.

28 Groups in the Periodic Table Alkali metals column 1 (or 1A) Alkaline earth metals column 2 (or 2A) Halogens (salt-formers) column 17 (7A) Noble gases column 18 (or 8A)

29 Why do the elements potassium and sodium have similar chemical properties?

30 Classify each element as a representative element, transition metal, or noble gas. a. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 b. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 c. 1s 2 2s 2 2p 6 3s 2 3p 2

31 Which of the following elements are transition metals: Cu, Sr, Cd, Au, Al, Ge, Co?

32 How many electrons are in the highest occupied energy level of a Group 5A element?

33 Trends in the Periodic Table

34 Terms to Know Shielding is the result of the core electrons blocking the pull of the nucleus on the outermost electrons. Nuclear charge is the charge in the nucleus.

35 Terms to Know Atomic Radius is one half the distance between the nuclei of two atoms of the same element when the atoms are joined. Pg. 170 Figure 6.13

36 Atomic Radius From left to right across a period, the atomic radius generally decreases. Why? As you go across a row, there is an increase in the number of protons and the number of electrons. However, the electrons being added are added to the same principle energy level, so the shielding remains the same. Therefore, the increased nuclear charge pulls the outermost electrons in closer to the nucleus. The result is a smaller atom. (See pg. 171, 3 rd paragraph.)

37 Atomic Radius From top to bottom down a group, the atomic radius generally increases. Why? As you go down a column, there is an increase in the number of protons and the number of occupied energy levels. The increase in positive charge pulls the electrons closer to the nucleus, but the increase in the number of occupied orbitals shields the nucleus more. The shielding effect is greater than the increased nuclear charge. The result is a larger atom. (See pg. 171, 1 st and 2 nd paragraph.)

38 Which element in each pair has atoms with a larger atomic radius? a. sodium, lithium b. strontium, magnesium c. carbon, germanium d. selenium, oxygen

39 Explain why fluorine has a smaller atomic radius than both oxygen and chlorine.

40 Terms to Know Ion is a neutral atom that has either lost or gained an electron. Atoms that have lost electron(s) are called Positive Ions or Cations Atoms that have gained electrons are called Negative Ions or Anions

41 Cation is an ion that has lost one or more electrons resulting in a net positive charge.

42 Anion is an ion that has gained one or more electrons resulting in a net negative charge.

43 Ion Size Cations are always SMALLER than the atoms from which they form. Why? When a metal atom loses an electron, the attraction between the remaining electrons and the nucleus is increased. The electrons are drawn closer to the nucleus. (See pg. 176, 2 nd paragraph)

44 Ion Size Anions are always LARGER than the atoms from which they form. Why? As the number of electrons increases, the attraction of the nucleus for any one electron decreases. (See pg. 176, 3 rd paragraph)

45 Li Be B C N O F Li 1+ Be 2+ B 3+ C N 3- O 2- F 1-

46 Ion Trends From left to right across a period, the cation size decreases. From left to right across a period, the anion size decreases. (See figure 6.20 on pg. 176)

47 Pg. 176 Figure 6.20

48 Ion Trends From top to bottom down a group, the ionic size increases. (Pg. 172)

49 Which particle has the larger radius in each atom/ion pair? a. Na, Na + b. S, S 2- c. I, I - d. Al, Al 3+

50 In each pair, which ion is larger? a. Ca 2+, Mg 2+ b. Cl -, P 3- c. Cu +, Cu 2+

51 The atomic radius from left to right across the periodic table. The atomic radius from top to bottom on the period table. Aside from the noble gases, what is the smallest atom on the periodic table? Aside from the noble gases, what is the largest atom on the periodic table?

52 The atom is than its cation. The atom is than its anion. Cations are than the anions that are in the same row. The cation trend from left to right is that the ionic radius. The anion trend from left to right is that the ionic radius.

53 List the following from smallest to largest according atomic radius. a. Mg, S, Na, Si b. Be, Ba, Ca, Ra c. F, As, Br, Ga, Cl List the following from smallest to largest according ionic radius. a. Mg 2+, S 2-, Na +, Cl - b. Be 2+, Ba 2+, Ca 2+, Ra 2+ c. F -, As 3-, Br -, Ga 3+, Cl -

54 Ionization Energy is the energy required to remove an electron from an atom. (Pg ) Terms to Know

55 Explain the difference between the terms first ionization energy and second ionization energy of an element. The first ionization energy is the energy needed to remove a first electron from an atom. The second ionization energy is the energy needed to remove a second electron.

56 Why is there a large increase between the first and second ionization energies of the alkali metals? It is relatively easy to remove the first electron from an alkali metal atom because they have only one valence electron; it is much more difficult to remove the second because it is being removed from a full energy level.

57 There is a large jump between the second and third ionization energies of magnesium. There is a large jump between the third and fourth ionization energies of aluminum. Explain these observations.

58 It is relatively easy to remove two electrons from magnesium because it has 2 valence electrons; it is much more difficult to remove a third electron because it comes from a full energy level. It is relatively easy to remove three electrons from aluminum because it has 3 valence electrons; it is much more difficult to remove a fourth electron because it comes from a full energy level.

59 Ionization Energy From left to right across a period, the ionization energy generally increases. Why? As you go across a row, there is an increase in the nuclear charge, but the shielding effect remains the same. There is an increase in the attraction of the nucleus for an electron. Thus it takes more energy to remove an electron from the atom. (See pg. 174, 2 nd paragraph.)

60 Ionization Energy From top to bottom down a group, the ionization energy generally decreases. Why? As the size of the atom increases and the amount of shielding electrons increases, nuclear charge has a smaller effect on the electrons in the outermost energy level. Therefore, less energy is required to remove an electron from the outermost energy level. (See pg. 174, 1 st paragraph.)

61 Which element in each pair has a greater first ionization energy? a. lithium, boron b. magnesium, strontium c. cesium, aluminum

62 Would you expect metals or nonmetals in the same period to have higher ionization energies? Give a reason for your answer.

63 Explain each of the following comparisons. Calcium has a smaller second ionization energy than does potassium.

64 Explain each of the following comparison. Lithium has a larger first ionization energy than does cesium.

65 Explain the following comparison. Magnesium has a larger third ionization energy than does aluminum.

66 Electronegativity Electronegativity is the ability of an atom to attract electrons when the atom is in a compound. (See pg. 177, 1 st paragraph)

67 Electronegativity For the representative elements: From left to right across a period, the electronegativity values increase. From top to bottom down a group, the electronegativity values decrease. (See pg. 177, 3 rd paragraph)

68 Electronegativity Fluorine is the most electronegative element on the periodic table. Why are the noble gases NOT included in electronegativity table? Because they are UNREACTIVE and do NOT want to bond with any atoms.

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70 Which element in each pair has a higher electronegativity value? a. Cl, F b. C, N c. Mg, Ne d. As, Ca

71 When the elements in each pair are chemically combined, which element in each pair has a greater attraction for electrons? a. Ca or O b. O or F c. H or O d. K or S

72 Pg. 171 Figure 6.14

73 Pg. 174 Figure 6.17

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