The How and Why

2 History Elements known to ancients: C, Cu, Au, Fe, Pb, Hg, Ag, S, Sn Added before 1700: As, Sb, Bi, P, Zn Dobereiner, Johann ( ): arranged elements in triads (Ca, Sr, Ba; Cl, Br, I) John Newlands ( ): arranged elements in group of eight Properties repeat every 8 th elements: Li, Be, B, C, N, O, F, Na Na, Mg, Al, Si, P, S, Cl, K

3 History Dmitri Mendeleev ( ): used the masses of elements as most of the masses were determined in XIX century (1869) Arranged elements in order of increasing atomic masses Found a pattern of repeating properties

4 Mendeleev’s Table Grouped elements in columns by similar properties in order of increasing atomic mass. Found some inconsistencies - felt that the properties were more important than the mass, so switched order ( Te, I). Found gaps in the trends- maybe undiscovered elements. Predicted their properties before they were found ( eka boron – Sc; eka aluminum- Ga; eka silicon: Ge).

5 The Modern Table Elements are still grouped by properties. Similar properties are in the same column. Order is in increasing atomic number (Moseley, 1914). Added a column of elements Mendeleev didn’t know about (Noble Gases). The noble gases weren’t found because they didn’t react with anything.

6 Periodic Law u Mendeleev (1869): Properties of elements are a function of the atomic masses of the elements. u Modern periodic Law (Mosley, 1914) properties of elements are a periodic function of their atomic numbers ( # of protons in the nucleus). Explains Mendeelev’s irregularities.

7 u Horizontal rows are called periods u There are 7 periods

8 u Vertical columns are called groups or families u Elements are placed in columns by similar properties.

9 1A 2A3A4A5A6A 7A 8A 0 The elements in the A groups are called the representative elements (also numbered 1-18)

10 The group B are called the transition elements Inner Transition elements

11 Group 1A(1) are the alkali metals Group 2A(2) are the alkaline earth metals

12 Group 7A (17) is called the Halogens Group 8A (18) are the noble or inert gases

13 Why the similarities in Properties? The part of the atom another atom sees is the electron cloud. These are the outside or valence orbitals. The orbitals fill up in a regular pattern. The outside orbital electron configuration repeats. The properties of atoms repeat.

14 1s 1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87

15 He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86 1s 2 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6

16 Alkali metals all end in s 1 Alkaline earth metals all end in s 2 Have to include He but it fits better later. He has the properties of the noble gases. s2s2 s1s1 S- block

17 Transition Metals -d block d1d1 d2d2 d3d3 s1d5s1d5 d5d5 d6d6 d7d7 d8d8 s 1 d 10 d 10

18 The P-block p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

19 F - block u inner transition elements

20 Each row (or period) is the energy level for s and p orbitals

21 d orbitals fill up after previous energy level … first d is 3d even though it’s in row d

22 f orbitals start filling at 4f f 5f

Writing Electron configurations the easy way Review Notes

24 Electron Configurations Repeat u The shape of the periodic table is a representation of this repetition of electron configurations. u When we get to the end of the column the outermost energy level is full. u This is the basis for our shorthand.

25 Trends and Properties of Elements in the Periodic Table u The following properties will be examined: Radius of the atom Ionization energy Electron affinity Electron negativity

26 Effective Nuclear Charge u Effective nuclear charge is experienced by an outer electron at the outer edge of an atom u Z eff = Z – S Z is the atomic number S is the number of core electrons. See blackboard for examples.

27 Atomic Size (Radius) First problem where do you start measuring. The electron cloud doesn’t have a definite edge. Determined by measuring more than 1 atom at a time.

28 Atomic Size u Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

29 Trends in Atomic Size Influenced by two factors. Energy Level Higher energy level is further away from nucleus Charge on nucleus ( Zeffective) More charge pulls electrons in closer.

30 Group Trends As we go down a group Each atom has another energy level So the atoms get bigger. H Li Na K Rb

31 Periodic Trends As you go across a period the radius gets smaller. Filling the same energy level. ***More nuclear charge (higher Zeff). Outermost electrons are pulled closer. NaMgAlSiPSClAr

32 Atomic Radii in the PT

33 Overall Atomic Number Atomic Radius (nm) H Li Ne Ar 10 Na K Kr Rb

34 Ionization Energy (IE) The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a +1 ion. The energy required to remove the first e- is called the first ionization energy. A(g) + IE → A(g) +1

35 Ionization Energy The 2 nd ionization energy is the energy required to remove the second electron. 2 nd IE is always greater than 1 st IE. The 3 rd IE is the energy required to remove a third electron. 3 rd IE > 2 nd IE > 1 st IE

36 SymbolFirstSecond Third H He Li Be B C N O F Ne

37 SymbolFirstSecond Third H He Li Be B C N O F Ne

38 What Determines IE 1. The greater the nuclear charge the greater IE. 2. Distance form nucleus influences IE. 3. Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. 4. Shielding effect

39 Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus. It is shielded from the nucleus by all the inner electrons A second electron in the same energy level has the same shielding.

40 Group Trends  As you go down a group first IE decreases because The electron is further away More shielding by inner electrons as there are more energy levels.

41 Periodic trends All the atoms in the same period have the same energy level. Same shielding. Increasing nuclear charge increases the force of attraction between the nucleus and the electrons. IE generally increases from left to right. Exceptions at full and 1/2 fill orbitals.

42 First Ionization energy Atomic number He He >IE than H same shielding greater nuclear charge H

43 First Ionization energy Atomic number H He Li < IE than H more shielding further away (>n) outweighs greater nuclear charge Li

44 First Ionization energy Atomic number H He Be > IE than Li same shielding greater nuclear charge Li Be

45 First Ionization energy Atomic number H He B < IE than Be same shielding greater nuclear charge By removing an electron we make s-orbital half filled Li Be B

46 First Ionization energy Atomic number H He Li Be B C

47 First Ionization energy Atomic number H He Li Be B C N

48 First Ionization energy Atomic number H He Li Be B C N O Breaks the pattern because removing an electron gets to 1/2 filled p orbital

49 First Ionization energy Atomic number H He Li Be B C N O F

50 First Ionization energy Atomic number H He Li Be B C N O F Ne Ne < IE than He Both are full, Ne has more shielding Greater distance (>n)

51 First Ionization energy Atomic number H He Li Be B C N O F Ne Na < IE than Li Both are s 1 Na has more shielding Greater distance (>n) Na

52 First Ionization energy Atomic number

53 Driving Force for Ionization u Full Energy Levels are very low energy. u Noble Gases have full orbitals. u Atoms behave in ways to achieve noble gas configuration.

54 2nd Ionization Energy u For elements that reach a filled or half filled orbital by removing 2 electrons 2nd IE is lower than expected. u True for s 2 u Alkali earth metals form +2 ions.

55 3rd IE u Using the same logic s 2 p 1 atoms have a low 3rd IE. u Atoms in the aluminum family form + 3 ions. u 2nd IE and 3rd IE are always higher than 1st IE!!!

56 Electron Affinity (EA) The energy change associated with adding an electron to a gaseous atom. Easiest to add to group 7A. Filled energy level. EA increases from left to right atoms become smaller, with greater nuclear charge. EA decreases as go down a group.

57 Ionic Size Cations form by losing electrons. Cations are smaller that the atom they come from. Metals form cations. Cations of representative elements have noble gas configuration.

58 Ionic size Anions form by gaining electrons. Anions are bigger than the atom they come from. Nonmetals form anions. Anions of representative elements have noble gas configuration.

59 Ionic Radii

60 Configuration of Ions Ions always have noble gas configuration. Na is 1s 1 2s 2 2p 6 3s 1 Forms a +1 ion - 1s 1 2s 2 2p 6 Same configuration as neon. Metals form ions with the configuration of the noble gas before them - they lose electrons.

61 Configuration of Ions Non-metals form ions by gaining electrons to achieve noble gas configuration. They form the configuration of the noble gas after them.

62 Group trends Adding energy level Ions get bigger as you go down a column Li +1 Na +1 K +1 Rb +1 Cs +1

63 Periodic Trends Across the period nuclear charge increases so ions get smaller. Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

64 Size of Isoelectronic ions Iso - same Iso-electronic ions have the same # of electrons Al +3 Mg +2 Na +1 Ne F -1 O -2 and N electrons configuration: 1s 1 2s 2 2p 6

65 Size of Atoms and Ions

66 Size of Isoelectronic ions Positive ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

67 Atomic Radii and Ionic Radii Compared

68 Comparison af Atomic and Ionic Radii

Electronegativity

70 Electronegativity u The tendency for an atom to attract electrons to itself when it is chemically combined with another element. u How fair it shares. u Big electronegativity means it pulls the electron toward it. u Atoms with large negative electron affinity have larger electronegativity. u Scale designed by Linus Pauling

71 Group Trend u The further down a group the farther the electron is away and the more electrons an atom has. u More willing to share. u Low electronegativity.

72 Periodic Trend u Metals are at the left end. u They lose electrons easily u Low electronegativity u At the right end are the nonmetals. u They gain electrons. u High electronegativity.

73 Ionization energy, electronegativity Electron affinity INCREASE

74 Atomic size increases, shielding constant Ionic size increases

75 Summary ot trends in the Periodic Table

76 Oxidation States Check Blackboard

77 Metal, nonmetals, metalloids, & noble gases

78 Properties of Metals u Elements with 3 or less electrons in outer level u Electrons loosely bound u Become positive ions u Good electric and heat conductors u Malleable, ductile u Most are solids at room temperature (exception : mercury) u Grayish, silvery in color (exceptions: Cu, Au)

79 Properties of Nonmetals u Elements with 5 or more electrons in outer level u Gain electrons easily to become negative ions u Brittle u Can be solid, liquid, or gas at room temperature u Good insulators u Many have colors in naturals state (sulfur – yellow, iodine: purple)

80 Transition Elements u Located in B group elements (or groups 3-12) u Many have multiple oxidation states (because of the d-orbitals) u Many form colorful compounds: Ni +2 green; Cu +2 – blue; Mn +7, purple u All are metals, some with the highest known melting points (Tungsten)

81 Properties of Metalloids u Located on the zig-zag line u Behave both as metals and nonmetals with some exhibiting stronger metallic character (aluminum)

82 Noble Gases u Located in group 8A(or 18) u Most are inert (exceptions: Xe – some compounds with oxygen and fluorine are known) u High ionization energies. u Have octet of electrons (exception: He with only 2 electrons)