Chapter 8 Periodic Properties & Electron Configurations Bushra Javed 1 Contents Electron Spin Electron Configurations of elements The development of Periodic Table Periodic Trends 2 Electron Configurations Quantum-mechanical theory describes the behavior of electrons in atoms The electrons in atoms exist in orbitals A description of the orbitals occupied by electrons is called an electron configuration 1s11s1 principal energy level of orbital occupied by the electron sublevel of orbital occupied by the electron number of electrons in the orbital 3 How Electrons Occupy Orbitals Calculations with Schrdingers equation show hydrogens one electron occupies the lowest energy orbital in the atom Schrdingers equation calculations for multi- electron atoms cannot be exactly solved due to electron-electron interactions in multi-electron atoms 4 How Electrons Occupy Orbitals The interactions that occur in multi-electron atoms are due to : electron spin & energy splitting of sublevels 5 The Property of Electron Spin Spin is a fundamental property of all electrons All electrons have the same amount of spin The orientation of the electron spin is quantized, it can only be in one direction or its opposite spin up or spin down The electrons spin adds a fourth quantum number to the description of electrons in an atom, called the Spin Quantum Number, m s not in the Schrdinger equation 6 The Property of Electron Spin Experiments by Stern and Gerlach showed a beam of silver atoms is split in two by a magnetic field. The experiment reveals that the electrons spin on their axis spinning charged particles generate a magnetic field Electron Spin If there is an even number of electrons, about half the atoms will have a net magnetic field pointing north and the other half will have a net magnetic field pointing south 8 Electron Spin 9 The two possible spin orientations of an electron and the conventions for m s are illustrated here. m s, and Orbital Diagrams m s can have values of + or Orbital Diagrams use a square to represent each orbital and a half-arrow to represent each electron in the orbital By convention, a half-arrow pointing up is used to represent an electron in an orbital with spin up Spins must cancel in an orbital paired 11 Orbital Diagrams We often represent an orbital as a square or a circle and the electrons in that orbital as arrows the direction of the arrow represents the spin of the electron orbital with one electron unoccupied orbital orbital with two electrons Pauli Exclusion Principle No two electrons in an atom may have the same set of four quantum numbers Therefore no orbital may have more than two electrons, and they must have the opposite spins Knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel 13 Pauli Exclusion Principle s sublevel has 1 orbital, therefore it can hold 2 electrons p sublevel has 3 orbitals, therefore it can hold 6 electrons d sublevel has 5 orbitals, therefore it can hold 10 electrons f sublevel has 7 orbitals, therefore it can hold 14 electrons 14 Sublevel Splitting in Multi-electron Atoms The sublevels in each principal energy shell of Hydrogen all have the same energy We call orbitals with the same energy degenerate For multi-electron atoms, the energies of the sublevels are split caused by charge interaction, shielding and penetration The lower the value of the l quantum number, the less energy the sublevel has s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3) 15 Shielding & Effective Nuclear Charge Each electron in a multi-electron atom experiences both the attraction to the nucleus and repulsion by other electrons in the atom These repulsions cause the electron to have a net reduced attraction to the nucleus it is shielded from the nucleus The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron Shielding & Penetration 17 penetration & Effective Nuclear Charge The closer an electron is to the nucleus, the more attraction it experiences The better an outer electron is at penetrating through the electron cloud of inner electrons, the more attraction it will have for the nucleus Penetration causes the energies of sublevels in the same principal level to not be degenerate Effect of Penetration and Shielding In the fourth and fifth principal levels, the effects of penetration become so important that the s orbital lies lower in energy than the d orbitals of the previous principal level The energy separations between one set of orbitals and the next become smaller beyond the 4s the ordering can therefore vary among elements causing variations in the electron configurations of the transition metals and their ions 19 Filling the Orbitals with Electrons Energy levels and sublevels fill from lowest energy to high s p d fs p d f Aufbau Principle Orbitals that are in the same sublevel have the same energy No more than two electrons per orbital Pauli Exclusion Principle When filling orbitals that have the same energy, place one electron in each before completing pairs Hunds Rule 20 Filling the Orbitals with Electrons The lowest-energy configuration of an atom is called its ground state. Any other allowed configuration represents an excited state. 21 Electron Configuration of Atoms The electron configuration is a listing of the sublevels in order of filling with the number of electrons in that sublevel written as a superscript. Kr = 36 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 22 Building up order of Filling in sublevels (Ground State Electron Configurations) 1s1s 2s2s2p2p 3s3s3p3p3d3d 4s4s4p4p4d4d4f4f 5s5s5p5p5d5d5f5f 6s6s6p6p6d6d 7s7s Start by drawing a diagram putting each energy shell on a row and listing the sublevels, (s, p, d, f), for that shell in order of energy (left-to-right) Next, draw arrows through the diagonals, looping back to the next diagonal each time 23 Electron Configuration & the Periodic Table 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 of the valence electrons 24 s1s1 s2s2 d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 s2s2 p 1 p 2 p 3 p 4 p 5 p6p6 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 f 14 d Blocks of Elements 26 Blocks of Elements For main-group (representative) elements, an s or a p subshell is being filled. For d-block transition elements, a d subshell is being filled. For f-block transition elements, an f subshell is being filled. 27 Electron Configurations Example 1 Write the ground state electron configuration of the chlorine atom, Cl, 1s21s2 2s22s2 2p62p6 3s23s2 3p53p5 For chlorine, Cl, Z = Electron Configurations Example 2 Write the ground state electron configuration of the manganese atom, Mn 29 Electron Configurations Example 3: What is the ground-state electron configuration of tantalum (Ta)? a.1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 10 4f 7 b.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 3 c.1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 3 d.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 30 Electron Configurations Example 4 Which of the following electron configurations corresponds to the ground state electron configuration of an atom of a transition element? a.1s 2 2s 2 2p 2 b.1s 2 2s 2 2p 6 3s 2 3p 5 c.1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 d.1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2 31 Exceptions to the Building up order of Filling About 21 of the predicted configurations are inconsistent with the actual configurations observed. One Possible Reason: Half filled and fully filled subshells are highly Stable Cr, Cu,Ag and U are some of the exceptions. 32 Exceptions to the Building up order of Filling Chromium (Z=24) and copper (Z=29) have been found by experiment to have the following ground- state electron configurations: Cr:1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 Cu:1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 In each case, the difference is in the 3d and 4s subshells. 8 | 33 Exceptions to the Building up order of Filling Expected Cr = [Ar]4s 2 3d 4 Cu = [Ar]4s 2 3d 9 Mo = [Kr]5s 2 4d 4 Ru = [Kr]5s 2 4d 6 Pd = [Kr]5s 2 4d 8 Found Experimentally Cr = [Ar]4s 1 3d 5 Cu = [Ar]4s 1 3d 10 Mo = [Kr]5s 1 4d 5 Ru = [Kr]5s 1 4d 7 Pd = [Kr]5s 0 4d 10 34 Electron Configurations Example 5 Which of the following electron configurations represents an allowed excited state of the indicated atom? a.He: 1s 2 b.Ne: 1s 2 2s 2 2p 6 c.Na: 1s 2 2s 2 2p 6 3s 2 3p 2 4s 1 d.P: 1s 2 2s 2 2p 6 3s 2 3p 2 4s 1 35 The Noble Gas core Electron Configuration The noble gases have eight valence electrons. except for He, which has only two electrons We know the noble gases are especially non-reactive He and Ne are practically inert The reason the noble gases are so non-reactive is that the electron configuration of the noble gases is especially stable 36 Noble Gas Core Electron Configuration A short-hand way of writing an electron configuration is to use the symbol of the previous noble gas in [] to represent all the inner electrons, then just write the last set. Rb = 37 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 = [Kr]5s 1 37 Noble Gas Core Electron Configuration Example 6 Which ground-state electron configuration is incorrect? a.Fe: [Ar] 3d 5 b.Ca: [Ar] 4s 2 c.Mg: [Ne] 3s 2 d.Zn: [Ar] 3d 10 4s 2 38 Writing Electron Configurations The pseudo-noble-gas core includes the noble- gas subshells and the filled inner, (n 1), d subshell. For bromine, the pseudo-noble-gas core is [Ar]3d 10 39 Hunds rule Hunds rule states: that the lowest-energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons 40 Electron Configurations 41 Electron Configurations Example 7 Draw an orbital diagram for nitrogen. 42 1s1s2s2s2p2p Electron Configurations Example 8 Which of the following electron configurations or orbital diagrams are allowed and which are not allowed by the Pauli exclusion principle? If they are not allowed, explain why? 43 a.1s 2 2s 1 2p 3 b.1s 2 2s 1 2p 8 c.1s 2 2s 2 2p 6 3s 2 3p 6 3d 8 d.1s 2 2s 2 2p 6 3s 2 3p 6 3d 11 e. 1s1s2s2s Electron Configurations Example 9 write the complete ground state orbital diagram and electron configuration of potassium 44 Valence Electrons The electrons in all the sublevels with the highest principal energy shell are called the valence electrons Electrons in lower energy shells are called core electrons Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons 45 46 47 Valence- shell configuration For main-group elements, the valence configuration is in the form ns A np B The sum of A and B is equal to the group number. So, for an element in Group VA of the third period, the valence configuration is 3s 2 3p 3 48 Valence-shell configuration Example 10 What are the valence-shell configuration for arsenic and zinc? 49 Arsenic is in period 4, Group VA. Its valence configuration is 4s24p3.4s24p3. Zinc, Z = 30, is a transition metal in the first transition series. Its noble-gas core is Ar, Z = 18. Its valence configuration is 4s 2 3d 10. Valence-shell configuration Example: 11 What is the valence -shell configuration of Tc, Z = 43 50 Mendeleevs Periodic Table Order elements by atomic mass Saw a repeating pattern of 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 51 Mendeleevs Periodic Table Used pattern to predict properties of undiscovered elements Mendeleevs Periodic Law allows us to predict what the properties of an element will be based on its position on the table 52 Mendeleev's Predictions 53 Everyone Wants to Be Like a Noble Gas! The Alkali Metals The alkali metals have one more electron than the previous noble gas In their reactions, the alkali metals tend to lose one electron, resulting in the same electron configuration as a noble gas forming a cation with a 1+ charge 54 Periodic Trends Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. We will look in more detail at three periodic properties: 1. atomic radius, 2. ionization energy, and 3. electron affinity. 55 Atomic Radius While an atom does not have a definite size, we can define it in terms of covalent radii (the radius in covalent compounds). 56 Trend in Atomic Radius Main Group 57 There are several methods for measuring the radius of an atom, and they give slightly different numbers van der Waals radius = nonbonding covalent radius = bonding radius Trend in Atomic Radius Main Group Atomic Radius Increases down group valence shell farther from nucleus effective nuclear charge fairly close Atomic Radius Decreases across period (left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer 58 Trend in Atomic Radius Main Group Trends Within each group (vertical column), the atomic radius increases with the period number. This trend is explained by the fact that each successive shell is larger than the previous shell. 59 Trend in Atomic Radius Main Group Within each period (horizontal row), the atomic radius tends to decrease with increasing atomic number (nuclear charge). 60 A representation of atomic radii is shown below. 61 Trend in Atomic Radius Main Group Example 12: Refer to a periodic table and arrange the following elements in order of increasing atomic radius: Br, Se, Te Br 34 Se 52 Te Te is larger than Se. Se is larger than Br. Br < Se < Te Trend in Atomic Radius Main Group Example 13 An atom of which of the following elements has the largest atomic radius? a) Rb b)Cl c)Mg d) P 63 Trend in Atomic Radius Main Group Example 14 An atom of which of the following elements has the smallest atomic radius? a)Mg b)Cl c)As d)Rb 64 Trend in Ionization Energy 65 Minimum energy needed to remove an electron from an atom or ion in the gaseous state endothermic process valence electron easiest to remove, lowest IE M(g) + IE 1 M 1+ (g) + 1 e - M +1 (g) + IE 2 M 2+ (g) + 1 e - first ionization energy = energy to remove electron from neutral atom; 2nd IE = energy to remove from 1+ ion; etc. Trend in Ionization Energy Going down a group, first ionization energy decreases. This trend is explained by understanding that the smaller an atom, the harder it is to remove an electron, so the larger the ionization energy. In general : E 1 < IE 1 < IE 3 66 Trend in Ionization Energy Generally, ionization energy increases with atomic number. Ionization energy is proportional to the effective nuclear charge divided by the average distance between the electron and the nucleus. Because the distance between the electron and the nucleus is inversely proportional to the effective nuclear charge, ionization energy is inversely proportional to the square of the effective nuclear charge. 67 Trend in Ionization Energy Example 15 Refer to a periodic table and arrange the following elements in order of increasing ionization energy: As, Br, Sb. 8 | 68 Sb is larger than As. As is larger than Br. Ionization energies: Sb < As < Br 35 Br 33 As 51 Sb Trend in Ionization Energy Small deviations occur between Groups IIA and IIIA and between Groups VA and VIA. Examining the valence configurations for these groups helps us to understand these deviations: IIAns 2 IIIAns 2 np 1 VA ns 2 np 3 VIA ns 2 np 4 69 It takes less energy to remove the np 1 electron than the ns 2 electron. It takes less energy to remove the np 4 electron than the np 3 electron. Trend in Ionization Energy Example16 An atom of which of the following elements has the largest second ionization energy? a)Li b)C c)F d)Be 70 Trend in Ionization Energy Example 17 Which of the following elements has larger first IE and WHY ? a)Al b)N c)O d)K 71 Electron affinity (E.A.) The energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion. A negative energy change (exothermic) indicates a stable anion is formed. The larger the negative number, the more stable the anion. Small negative energies indicate a less stable anion. A positive energy change (endothermic) indicates the anion is unstable. 72 Trends in Electron affinity (E.A.) Electron affinities in the main-group elements show a periodic variation when plotted against atomic number, although this variation is somewhat more complicated than that displayed by ionization energies. In a given period, the electron affinity rises from the Group IA element to the Group VIIA element but with sharp drops in the Group IIA and Group VA elements. 73 8 | 74 Trends in Electron affinity (E.A.) Broadly speaking, the trend is toward more negative electron affinities going from left to right in a period. Lets explore the periodic table by group. 75 76 Trend in Electron Affinity (EA) Example 18 An atom of which of the following elements has the most negative electron affinity? a)K b)Cl c)Br d)Se e)N 77 Groups IIA and VIIIA do not form stable anions; their electron affinities are positive. Group ValenceAnion Valence IAns 1 ns 2 stable IIIAns 2 np 1 ns 2 np 2 stable IVAns 2 np 2 ns 2 np 3 stable VAns 2 np 3 ns 2 np 4 not so stable VIAns 2 np 4 ns 2 np 5 very stable VIIAns 2 np 5 ns 2 np 6 very stable Except for the members of Group VA, these values become increasingly negative with group number. 78 Metallic Character Elements with low ionization energies tend to be metals. Those with high ionization energies tend to be nonmetals. This can vary within a group as well as within a period. 79 Metallic Character Metallic character is how closely an elements properties match the ideal properties of a metal more malleable and ductile, better conductors, and easier to ionize Metallic character decreases left-to-right across a period metals are found at the left of the period and nonmetals are to the right Metallic character increases down the column nonmetals are found at the top of the middle Main Group elements and metals are found at the bottom 80 81 Oxides of Main Group Elements Oxides A basic oxide reacts with acids. Most metal oxides are basic. If soluble, their water solutions are basic. An acidic oxide reacts with bases. Most nonmetal oxides are acidic. If soluble, their water solutions are acidic. An amphoteric oxide reacts with both acids and bases. 82 Trends in Metallic Character Group IA, Alkali Metals (ns 1 ) These elements are metals; their reactivity increases down the group. The oxides have the formula M 2 O. Hydrogen is a special case. It usually behaves as a nonmetal, but at very high pressures it can exhibit metallic properties. 83 Reactions of Alkali Metals with Water Alkali metals are oxidized to the 1+ ion H 2 O is split into H 2 (g) and OH ion The Li, Na, and K are less dense than the water so float on top The ions then attach together by ionic bonds The reaction is exothermic, and often the heat released ignites the H 2 (g) 84 Trends in Metallic Character Group IIA, Alkaline Earth Metals (ns 2 ) These elements are metals; their reactivity increases down the group. The oxides have the formula MO. 85 Trends in Metallic Character Group IIIA (ns 2 np 1 ) Boron is a metalloid; all other members of Group IIIA are metals. The oxide formula is R 2 O 3. B 2 O 3 is acidic. Al 2 O 3 and Ga 2 O 3 are amphoteric The Group IIIA oxides are basic. 86 Trends in Metallic Character Group IVA (ns 2 np 2 ) Carbon is a nonmetal; silicon and germanium are metalloids; tin and lead are metals. The oxide formula is RO 2 and, for carbon and lead, RO. CO 2, SiO 2, and GeO 2 are acidic (decreasingly so). SnO 2 and PbO 2 are amphoteric. 87 PbO (yellow) PbO 2 (dark brown) SiO 2 (crystalline solid quartz) SnO 2 (white) Some oxides of Group IVA 88 Group VA (ns 2 np 3 ) Nitrogen and phosphorus are nonmetals; arsenic and antimony are metalloids; bismuth is a metal. The oxide formulas are R 2 O 3 and R 2 O 5, with some molecular formulas being double these. Nitrogen, phosphorus, and arsenic oxides are acidic. Antimony oxides are amphoteric. Bismuth oxide is basic. 89 Group VIA, Chalcogens (ns 2 np 4 ) Oxygen, sulfur, and selenium are nonmetals; tellurium is a metalloid; polonium is a metal. The oxide formulas are RO 2 and RO 3. Sulfur, selenium, and tellurium oxides are acidic except for TeO 2, which is amphoteric. PoO 2 is also amphoteric. 90 Group VIIA, Halogens (ns 2 np 5 ) These elements are reactive nonmetals, with the general molecular formula being X 2. All isotopes of astatine are radioactive with short half-lives. This element might be expected to be a metalloid. Each halogen forms several acidic oxides that are generally unstable. 91 Group VIIIA, Noble Gases (ns 2 np 6 ) These elements are generally unreactive, with only the heavier elements forming unstable compounds. They exist as gaseous atoms. 92