The Periodic Table. Dmitri Mendeleev This photograph is in the public domain.

Transcription

1 History Electron configurations The organization of the modern table Some different types of tables Periodic trends in the table Some example problems

2 Dmitri Mendeleev This photograph is in the public domain.

3 ordered elements by atomic mass saw a repeating pattern of chemical and physical properties Periodic Law When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically put elements with similar properties in the same column

4 used patterns to predict properties of undiscovered elements where atomic mass order did not fit other properties, he re-ordered by other properties ie. Te & I

5 This diagram is in the public domain.

6 Predicted Properties of New Elements Property Ekaaluminium atomic mass 68 density (g/cm³) 6.0 melting point ( C) Low oxide EaO3 chloride EaCl6 Gallium GaO3 GaCl6

7 Predicted Properties of New Elements Property Ekasilicon atomic mass 7 density (g/cm³) 5.5 melting point ( C) high oxide EaO chloride EaCl4 germanium GeO GeCl4

8 Electron Configurations The four quantum numbers n, l, ml and ms enable us to label completely an electron in any orbital in any atom. It is as though the set of quantum numbers for an individual electron are an 'address' for the electron. For example, the set of quantum numbers for a s electron would be (n, l, ml, ms), 0, 0, +1/ or -1/.

9 Electron Configurations The hydrogen atom is a particularly simple system since it contains only one electron. The electron may reside in the 1s orbital (the ground state) or it may be found in some higher energy orbital (an excited state). For manyelectron atoms, however, we must know the electron configuration of the atom... in other words, how the electrons are distributed among the various atomic orbitals.

10 Electron Configurations The Aufbau Principle: electrons occupy orbitals in a way that minimizes the energy of the atom. In other words, electrons fill the atomic orbitals starting with 1s and proceeding to the next highest energy orbital

11 energy Electron Configurations 011 K. Brown

12 Electron Configurations 011 K. Brown

13 Electron Configurations The Pauli Principle: two electrons with the same spin cannot occupy the same orbital. In other words, no two electrons can have the same set of quantum numbers (the same 'address'). Hund's Rule: the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. In other words, when degenerate orbitals are available for filling, the electrons will occupy them singly with the same spin.

14 Electron Configurations Hydrogen 1 electron The Aufbau principle indicates that the electron will go into the lowest energy orbital, the 1s orbital: H[ ] or 1s 1

15 energy Electron Configurations 011 K. Brown

16 Electron Configurations Helium electrons He[ ] or 1s

17 energy Electron Configurations 011 K. Brown

18 Electron Configurations Lithium 3 electrons Li [ ] [ ] or 1s s 1

19 energy Electron Configurations 011 K. Brown

20 Electron Configurations Berylium 4 electrons Be[ ] [ ] ][ ][ or 1s s Boron 5 electrons B[ ] 1s [ ] [ s p ] or 1s s p 1

21 energy Electron Configurations 011 K. Brown

22 Electron Configurations Carbon 6 electrons C[ ] 1s [ ] [ s ][ ][ p ] or 1s s p

23 energy Electron Configurations 011 K. Brown

24 Electron Configurations Nitrogen 7 electrons N[ ] [ 1s ] [ s 3 4 ][ ] or 1s s p ][ ] or 1s s p ][ p Oxygen 8 electrons O[ ] 1s [ ] s [ ][ p

25 Electron Configurations Fluorine 9 electrons F[ ] [ 1s ] [ ][ s ][ ] or 1s s p 5 ] or 1s s p 6 p Neon 10 electrons Ne[ ] [ 1s ] s [ ][ ][ p

26 Electron Configurations Most often we are interested only in the outermost shell of electrons... these are the ones that are most involved in chemical reactions. This is reflected in the way the electron configuration is often written. Instead of explicitly showing all electrons we write the symbol for the previous noble gas followed by the outer shell or valence shell electron configuration.

27 Electron Configurations Carbon 1s s p [ He] s p long form compact form valence electrons Aluminum 6 1s s p 3s 3p long form 1 1 [ Ne] 3s 3p compact form

28 Electron Configurations In the periodic table the left-most columns include the alkali metals and the alkaline earth metals. In these elements the valence s orbitals are being filled On the right hand side, the right-most block of six elements are those in which the valence p orbitals are being filled

29 Electron Configurations In the middle is a block of ten columns that contain the transition metals. These are elements in which d orbitals are being filled Below this group are two rows with 14 columns. These are commonly referred to the f-block metals or the lanthanides and actinides. In these columns the f orbitals are being filled

30 Electron Configurations 013 K. Brown

31 Electron Configurations The electron configurations of some of the transition metal elements do not follow the regular rules. Chromium 4 [ Ar ] 4s 3d expected configuration 1 5 [ Ar ] 4s 3d actual configuration

32 Electron Configurations The electron configurations of some of the transition metal elements do not follow the regular rules. Copper 9 [ Ar ] 4s 3d expected configuration 1 10 [ Ar ] 4s 3d actual configuration

33 Electron Configurations The electron configurations of ions are generally straightforward with some exceptions. For anions we simply add the extra electrons using the regular rules for constructing electron configurations. For cations we remove electrons in reverse order for the main group elements. For transition metal elements however we remove electrons from the orbital with the highest value of n first.

34 Electron Configurations O : 1s s p Cl : 1s s p 3s 3p O : 1s s p K :1s s p 3s 3p 4s Cl : 1s s p 3s 3p Cr : 1s s p 3s 3p 4s 3d K : 1s s p 3s 3p Cr :1s s p 3s 3p 3d 3

35 Electron Configurations Species with the same electron configuration are isoelectronic. - O : 1s s p - F : 1s s p 6 Ne : 1s s p Na :1s s p + 6 Mg : 1s s p 6

36 Shielding Why do we see periodic physical and chemical properties of the elements? We need to first understand the concept of shielding before this can be satisfactorily explained.

37 Shielding in a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other outer shell electrons are shielded from full strength of nucleus by the inner shell electrons effective nuclear charge is net positive charge that is attracting a particular electron in an outer shell

38 Shielding This diagram is licensed under Creative Commons.

39 Shielding Z is nuclear charge, S is electrons in lower energy levels Z effective = Z S

40 Valence Electrons and Periodic Properties the electrons in all the subshells with the highest principal energy shell are called the valence electrons electrons in lower energy shells are called core electrons

41 Valence Electrons and Periodic Properties the Group number corresponds to the number of valence electrons the length of each block is the maximum number of electrons the sublevel can hold the Period number corresponds to the principal energy level, n, of the valence electrons

42 Valence Electrons and Periodic Properties the number of valence electrons largely determines the behavior of an element since the number of valence electrons follows a periodic pattern, the properties of the elements should also be periodic quantum mechanical calculations show that 8 valence electrons should result in a very unreactive atom, an atom that is very stable and the noble gases, that have 8 valence electrons are all very stable and unreactive

43 Valence Electrons and Periodic Properties conversely, elements that have either one more or one less electron should be very reactive and the halogens are the most reactive nonmetals and alkali metals the most reactive metals many metals and nonmetals form one ion, and the charge on that ion is predictable based on its position on the Periodic Table Group 1A = +1, Group A = +, Group 7A = -1, Group 6A = -, etc.

44 Valence Electrons and Periodic Properties these atoms form ions that will result in an electron configuration that is the same as the nearest noble gas

45 Periodic Properties of the Elements Atomic radius increases going down a group Atomic radius decreases across period (left to right)

46 Periodic Properties of the Elements licensed under Gnu Free Documentation.

47 Periodic Properties of the Elements

48 Periodic Properties of the Elements Ion size increases down the group Cations smaller than neutral atom; Anions bigger than neutral atom Cations smaller than anions Larger positive charge = smaller cation Larger negative charge = larger anion

49 Periodic Properties of the Elements radius of positive ions is smaller than that of a neutral atom, which in turn is smaller than that of a negative ion. Thus the sizes of the isoelectronic atoms and ions have the following order: S- > Cl- > Ar > K+ > Ca+

50 Periodic Properties of the Elements licensed under Creative Commons.

51 Periodic Properties of the Elements

52 Periodic Properties of the Elements

53 Periodic Properties of the Elements electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field. This is called paramagnetism electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field this is called diamagnetism

54 Periodic Properties of the Elements both Zn atoms and Zn+ ions are diamagnetic, showing that the two 4s electrons are lost before the 3d 10 Zn:[ Ar ] 4s 3d Zn :[ Ar ] 4s 3d

55 Periodic Properties of the Elements Since the number paired and unpaired electrons is determined by the electron configuration and the electron configuration is integrally linked to the modern periodic table, diamagnetism and paramagnetism are calculable properties of the elements in the table. 7 Co :[ Ar ] 4s 3d Mg :[ Ne] 3s paramagnetic diamagnetic

56 Periodic Properties of the Elements Ionization Energy The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. 013 K. Brown

57 Periodic Properties of the Elements Ionization Energy The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy.

58 Periodic Properties of the Elements Ionization Energy Ionization energies increase moving from left to right across a period (decreasing atomic radius and increasing effective nuclear charge). Ionization energy decreases moving down a group (increasing atomic radius, same effective nuclear charge). Group I elements have low ionization energies because the loss of an electron forms a stable octet of electrons.

59 Periodic Properties of the Elements Ionization Energy The effective nuclear charge can be used to explain the overall observations concerning the periodicity of ionization energies. As Zeff increases across a period the outer shell electrons are held more tightly and require more energy to remove from the atom.

60 Periodic Properties of the Elements Ionization Energy licensed under Creative Commons.

61 Periodic Properties of the Elements Ionization Energy Going down a group, the average distance that the outer shell electrons are from the nucleus increases as Zeff remains constant. Consequently, the outer shell electrons are held less tightly and can be removed with less energy than those above them in the group.

62 Periodic Properties of the Elements Ionization Energy 013 K. Brown

63 Periodic Properties of the Elements Successive Ionization Energies removal of each successive electron costs more energy regular increase in energy for each successive valence electron large increase in energy when start removing core electrons

64 Periodic Properties of the Elements Successive Ionization Energies licensed under Creative Commons.

65 Periodic Properties of the Elements Electron Affinity Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. 013 K. Brown

66 Periodic Properties of the Elements Electron Affinity Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affinities of certain groups in the periodic table.

67 Periodic Properties of the Elements Electron Affinity The Group IIA elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high electron affinities because the addition of an electron to an atom results in a completely filled shell.

68 Periodic Properties of the Elements Electron Affinity 013 K. Brown

69 Periodic Properties of the Elements Metals and Non-metals Metals malleable & ductile shiny, lusterous, reflect light conduct heat and electricity most oxides basic and ionic form cations in solution lose electrons in reactions - oxidized

70 Periodic Properties of the Elements Metals and Non-metals Nonmetals brittle in solid state dull electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions - reduced

71 Periodic Properties of the Elements Metals and Non-metals 013 K. Brown

72 Periodic Properties of the Elements Alkali Metals atomic radius increases down the column ionization energy decreases down the column very low ionization energies Reactivity increases down the column electron affinity decreases down the column melting point decreases down the column

73 Periodic Properties of the Elements The Halogens atomic radius increases down the column ionization energy decreases down the column very high electron affinities reactivity decreases down the column react with hydrogen to form HX (X halogen) and chlorides with metals

74 Periodic Properties of the Elements The Noble Gases atomic radius increases down the column ionization energy decreases down the column very unreactive

75 The Stories of the Elements The prehistorical elements: Copper Iron Lead Platinum Sulfur Gold Carbon Mercury Silver Tin

76 The Stories of the Elements Hydrogen discovered by Cavendish, named by Lavoisier Oxygen discovered by Scheele and Priestly, named by Lavoisier Nitrogen discovered by Rutherford, named by Lavoisier Phosphorus discovered by Brand in 1669.

77 The Stories of the Elements

78 The Stories of the Elements Yttrium - discovered 1794 by Johann Gadolin Terbium - discovered 1843 by Mosander Erbium - discovered 1843 by Mosander Ytterbium - discovered 1878 by Marignac Gadolinium - discovered 1880 by Marignac Holmium - discovered 1878 by Cleve & Soret Thulium - discovered 1879 by Cleve

79 The Stories of the Elements

80 The Stories of the Elements

81 The Stories of the Elements

82 013 K. Brown

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91 Example Problems 1. Arrange the following in order of increasing first ionization energy: Na, Cl, Al, S, Cs. Cs < Na < Al < S < Cl

92 Example Problems. Arrange the following species in isoelectronic pairs: O+, Ar, S-, Ne, Zn, Cs+, N3-, As3+, N, Xe. O+, N Ar, SNe, N3Zn, As3+ Cs+, Xe 5s4d105p6 1ssp3 [Ne] 3s3p6 1ss3p6 [Ar] 4s3d10 [Kr]

93 Example Problems 3. Based on valences and typical anion and cation charges, predict the likely formulas for the reaction between francium and oxygen, strontium and bromine, radium and selenium. FrO ie. Fr+ and O- SrBr Sr+ and Br- RaSe Ra+ and Se-

94 Example Problems 4. Arrange the following atoms in order of decreasing atomic radius: Na, Al, P, Cl, Mg. Na > Mg > Al > P > Cl

95 Example Problems 5. Arrange the following isoelectronic species in order of increasing a) radius and b) ionization energy: O-, Mg+, F-, Na+ a) Mg+ < Na+ < F- < Ob) O- < F- < Na+ < Mg+

96 Example Problems 6. Element 10 (El) has just been discovered. What will the most likely formula of its oxide be? Will it be larger or smaller than the previous element in the table? Element 10 will have electron configuration [Rn] 7s6d107p6 8s... it has two valence shell electrons in the 8s orbital which puts it in group two. The 'standard' ion charge for a group two element is + so the formula for the oxide of element 10 should be ElO.

97 Example Problems 6. Element 10 (El) has just been discovered. What will the most likely formula of its oxide be? Will it be larger or smaller than the previous element in the table? The previous element would be element 119 which would have to be in group 1. The trend in atomic radii is for them to decrease from left to right across the table so that element 10 should be smaller than element 119.