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CHAPTER 7 The Structure of Atoms and Periodic Trends Chapter Outline § § § § § § 7-1 The Pauli Exclusion Principle 7-2 Atomic Subshell Energies and Electron Assignments 7-3 Electron Configuration of Atoms 7-4 Electron Configuration of Ions 7-5 Atomic Properties and Periodic Trends 7-6 Periodic Trends and Chemical Properties Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (ml) Chapter Goal: Use our knowledge of atomic structure to rationalize chemical behavior. The Pauli Exclusion Principle • The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers n, l, ml and ms. – For a given orbital the values of n, l, and ml are fixed. – If we want to put more than one electron in an orbital they must have different ms values. – Thus, an orbital can hold a maximum of two electrons, and they must have opposite spins. Wolfgang Pauli 1900-‐1958, Austrian physicist, 1945 Nobel Prize in physics "for the discovery of the Exclusion Principle, also called the Pauli Principle" 4 Orbital energies In hydrogen electron energy is only dictated by n Because of electron interac5ons, larger atoms have energy dictated by n AND l 1.) The energy of an electron is ordered relaPve to n+l 2.) For electrons with same n+l, the one with a lower n is lower in energy 5 PenetraPon and Orbital Energy The 2s orbital distribuPon has a small probability contribuPon closer to nucleus (orbital penetra,on) This region lowers the screening of a 2s electron by the 1s electron 2p d6 oesn’t have this penetraPon, so it is higher in energy (less stable) Factors AffecPng Atomic Orbital Energies The Effect of Nuclear Charge (Zeffective or Z*) Higher nuclear charge lowers orbital energy (stabilizes the system) by increasing nucleus-electron attractions. The Effect of Electron Repulsions (Shielding or Screening) Additional electron in the same orbital An additional electron raises the orbital energy through electron-electron repulsions. Additional electrons in inner orbitals Inner electrons shield outer electrons more effectively than do electrons in the same sublevel. Problem Determine the relaPve energies of all the orbitals in n=1, 2 and 3 8 Orbital Filling: The Aufbau Principle & Hund’s Rule • Aufbau Principle: • Lower energy orbitals fill first to give the ground state. • Hund’s Rule: • Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before pairing up electrons. • Pauli exclusion principle: • Individual orbitals only hold two electrons, and each should have different spin. Orbital Filling: Notation Orbital diagram: represent the orbital with a box and the electron with an arrow spdf notation: write the orbital with a superscript indicating the number of electrons in that orbital Noble gas notation: represent the core electrons with a noble gas symbol [X] and write out the valence electrons Problems Write out (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas notation for the following: (a) Li (b) B (c) N (d) F (e) Ne (f) Sc and (g) Ti Transition Metals Transition elements have partially filled d-orbitals Electron filling follows the same principles as we have already discussed. Two notable exceptions involve having a ground state configuration with an unfilled s-orbital in order to: 1.) maximize the electron spin in the d-orbitals (Hund) 2.) Completely filling out the d-orbitals Problems Write out (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas notation for the following: (a) Cr (b) Cu Electron Configuration of Ions Start with electron configurations of the neutral element For: Anions: add electrons to obtain a noble gas configuration Cations: remove electrons from orbital with highest n until a particularly stable configuration is obtained (filled subshell, high spin) Problems Write out (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas notation for the following: (a) Ca2+ (b) Ti2+ (c) Cr+ (d) Ag+ Problems We have already encountered transition metal ions that can take on more than one charge state. Predict the (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas notation the following: (a) Fe2+ and Fe3+ (b) Cu+ and Cu2+ (c) Co2+ and Co3+ Magnetic Properties of Atoms and Ions Paired electrons – diamagnetic Ferromagnetic- metals with magnetic properties Unpaired electrons - paramagnetic attracted by a magnetic field attraction proportional to number of unpaired e– Problem Suppose an element has the electronic configuration: 1s2 2s2 2p6 3s2 3p6 4s13d3 Is the element in the ground state? Problem (a) Depict the electron configuration of manganese, Mn, and its 4+ cation, Mn4+, using noble gas configuration and orbital diagrams. (b) Determine the magnetic properties of MnO2. Will this substance be more or less magnetic than solid manganese Mn (s) General Periodic Trends Elements were arranged in the periodic table historically based on their chemical properties. The modern understanding of atomic structure allows us to rationalize this arrangement Will look at three properties: Atomic and ionic size - radius Ionization energy - complete removal of an electron Electron affinity - addition of an electron General Periodic Trends All three properties can be rationalized by considering effective nuclear charge. Higher Z* yields more tightly bound electrons and: Smaller atomic and ionic size - electrons orbit closer to nucleus Larger Ionization energy - more tightly bound = more energy required to remove More negative electron affinity - addition of an electron more favorable Atomic Radii of Elements IonizaPon Energy • IonizaPon energy – energy (in kJ) required for the removal of 1 mol of electrons from 1 mol of atoms or ions in the gas phase. • First ionizaPon energy, IE1, is the energy needed to remove the first electron from a neutral atom. Na(g) → Na+(g) + e-‐ • Second ionizaPon energy, IE2, is the energy needed to remove the second electron. Na+(g) → Na2+(g) + e-‐ – IonizaPon energies are always posiPve, i.e., removing an electron is endothermic. 24 Problem Predict the periodic trend for first ionizaPon energy: (a) Across a period (b) Down a group Removing More Than 1 Electron The general trend: IE1 < IE2 < IE3 < IE4… Notice that the jump from IE1 to IE2 for Be is not as great as that for Li. Use orbital diagrams to rationalize Problem: Determine the group that element ‘X’ belongs to: 50,000 45,000 E, kJ 40,000 35,000 X 30,000 25,000 20,000 15,000 10,000 5,000 0 1st 2nd 3rd 4th 5th 6th Ionization Energies 7th 8th Electron Affinity • Electron Affinity (EA) – energy change accompanying the addiPon of 1 mol of electrons to 1 mol of gaseous atoms or ions. – In most cases, energy is released when the first electron is added. X(g) + e-‐ → X-‐(g) ΔE = EA1 28 Electron Affinity EA1 show numerous excepPons to expected periodic trends. Group 6A & 7A – Highly negaPve EA1 values, tend to form anions. Group 1A & 2A – Only slightly negaPve EA1 values, tend to form caPons. Noble Gases – Slightly posiPve EA1 and very high IE1, tend not to form ions. In general, groups 2 and 5 have less electron affinity than expected. Can this be ra5onalized with electron configura5ons? 29 Electron Affinity - Problem Would you expect the following reaction to occur: Br2 + I− → I2 + Br − Problem: ion sizes Rank the following ions in order of decreasing size. Na+, N3-, Mg2+, F– O2– Notice that they all have 10 electrons: They are isoelectronic (same electron configuration) as Ne. Ion size Group Trends for the Active Metals Group 1A: The Alkali Metals Form +1 ions because removing core electrons energetically unfavorable Alkali Metals + H2O 2 M (s) + 2 H2O (l) → 2 MOH (aq) + H2 (g) This reaction typically more favorable for metals with lower ionization energies Group Trends for the Active Metals Group 7A: The Halogens Form -1 ions because this give a noble gas configuration which fills out p subshell. Halogens Group 7A: The Halogens The chemistry of the halogens is dominated by gaining an electron to form an anion: X2 + 2e− → 2X− With water: 2F2(g) + 2H2O(l) → 4HF(aq) + O2(g) ΔH = −758.7 kJ This reaction typically more favorable for halogens with higher electron affinities Suggested problems: from the end of chapter 7 of the text book: Electron configura5ons: 1, 3, 5, 7, 9 Quantum numbers: 11, 13, 15 Ions: 17, 19, 21 Periodic proper5es: 23, 25, 27, 29, 31 General: 37, 39, 41, 43, 45, 47, 49, 51, 53, 55, 57, 59, 63, 65, 67, 69, 71, 75, 77