Arrangement of Electrons in Atoms

Page III-6b- / Chapter Six Part II Lecture Notes The Structure of Atoms and Periodic Trends Chapter Six Part Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) CH Professor Michael Russell ORBITALS (ml) Arrangement of Electrons in Atoms Electron Spin Quantum Number, ms Each orbital can be assigned no more than electrons! This is tied to the existence of a 4th quantum number, the electron spin quantum number, ms. ms arises naturally when relativity (Einstein) combined with quantum mechanics (Paul Dirac) Paul Dirac Magnetism In diamagnetic systems, all electron spins are paired - no net magnetic moment. In paramagnetic systems, unpaired spins present. Magnetic fields randomly arranged unless placed in an external magnetic field. Electron spin can be proven experimentally. Two spin directions are given by ms where ms = +/ and -/. Leads to magnetism in atoms and ions In ferromagnetic substances the orientations of magnetic fields from unpaired electrons are affected by spins from electrons around them. When an external field is applied and then removed, the substance maintains the magnetic moment and becomes a permanent magnet. Page III-6b- / Chapter Six Part II Lecture Notes Magnetism H - a magnetic field

Page III-6b- / Chapter Six Part II Lecture Notes Electron Spin Quantum Number Measuring Paramagnetism Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons. Diamagnetic: NOT attracted to a magnetic field; spin paired. Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons. Sample of FeO with strong magnet Sample of FeO Diamagnetic: NOT attracted to a magnetic field Pauli Exclusion Principle No two electrons in the same atom can have the same set of 4 quantum numbers. QUANTUM NUMBERS,,, 4,... l ---> subshell,,,... (n - ) ml ---> orbital -l...... +l ms ---> electron spin +/ and -/ That is, each electron has a unique address which will consist of its own values of n, l, ml and ms. See: Quantum Numbers Handout Wolfgang Pauli Electrons in Atoms When n =, then l = and ml = this shell has a single orbital (s) to which e- can be assigned n =, l =, ml =, ms = + / - this is electron # n =, l =, ml =, ms = -/ - this is electron # When n =, then l = (s), (p) e- three p orbitals 6e- TOTAL = 8e- electron number s orbital n l ml ms / -/ s - / p 4 - -/ p 5 / p 6 -/ p 7 + / p 8 + -/ p s Electrons in Atoms When n =, then l = (s), (p), (d) s orbital ethree p orbitals 6efive d orbitals etotal = 8eEach electron has its own set of four quantum numbers! Wolfgang Pauli Page III-6b- / Chapter Six Part II Lecture Notes electron number n ---> shell 4 5 6 7 8 9 4 5 6 7 8 n l ml - - + + - - - - + + + + ms / -/ / -/ / -/ / -/ / -/ / -/ / -/ / -/ / -/ s s p p p p p p d d d d d d d d d d

electron number Electrons in Atoms n l ml ms 4 / 4s 4 -/ 4s 4 - / 4p 4 4 - -/ 4p 5 4 / 4p 6 4 -/ 4p 7 4 + / 4p 8 4 + -/ 4p 9 4 - / 4d 4 - -/ 4d 4 - / 4d 4 - -/ 4d 4 / 4d 4 4 -/ 4d 5 4 + / 4d 6 4 + -/ 4d 7 4 + / 4d 8 4 + -/ 4d When n = 4, l = (s), (p), (d), (f) 4s orbital e- three 4p orbitals 6e- five 4d orbitals e- seven 4f orbitals 4e- TOTAL = e- and so on and so on... Page III-6b- / Chapter Six Part II Lecture Notes electron number n l ml ms 9 4 - / 4f 4 - -/ 4f 4 - / 4f 4 - -/ 4f 4 - / 4f 4 4 - -/ 4f 5 4 / 4f 6 4 -/ 4f 7 4 + / 4f 8 4 + -/ 4f 9 4 + / 4f 4 + -/ 4f 4 + / 4f 4 + -/ 4f Distribution of Electrons in Shells Assigning Electrons to Atoms Electrons generally assigned to orbitals of successively higher energy. For H atoms, E = - Rhc(/n ). E depends only on n. For many-electron atoms, energy depends on both n and l. Assigning Electrons to Subshells In H atom all subshells of same n have same energy. In many-electron atom: a) subshells increase in energy as value of n + l increases. (The important n + l rule) b) for subshells of same n + l, subshell with lower n is lower in energy. See Electron Configurations Handout Using the n + l rule assumes zero point energy - the lowest energy state possible, or ground state An Aufbau Diagram Electron Filling Order Writing Atomic Electron Configurations Two ways of writing configs. One is called the spectroscopic or "spdf" notation. SPECTROSCOPIC NOTATION for H, atomic number = s no. of electrons Aufbau comes from a German word meaning building up, formulated by Bohr and Pauli in 9s Page III-6b- / Chapter Six Part II Lecture Notes value of n value of l

Writing Atomic Electron Configurations Page III-6b-4 / Chapter Six Part II Lecture Notes Two ways of writing configs. Other is called the orbital box notation. ORBITAL BOX NOTATION for He, atomic number = s s Arrows depict electron spin One electron has n =, l =, m l =, m s = + / Other electron has n =, l =, m l =, m s = - / Atomic Electron Configurations Diagram Electron Configurations and the Periodic Table Lithium Group A Atomic number = s s ---> total electrons paramagnetic s s p p s Beryllium Boron s p Group A Atomic number = 4 s s ---> 4 total electrons diamagnetic s p Group A Atomic number = 5 s s p ---> 5 total electrons paramagnetic s p s p s s Page III-6b-4 / Chapter Six Part II Lecture Notes

Page III-6b-5 / Chapter Six Part II Lecture Notes s s s p p Carbon Group 4A Atomic number = 6 s s p ---> 6 total electrons paramagnetic Friedrich Hund Here we see for the first time HUND'S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible. s s s p p Nitrogen Group 5A Atomic number = 7 s s p ---> 7 total electrons paramagnetic Oxygen Fluorine s p Group 6A Atomic number = 8 s s p 4 ---> 8 total electrons paramagnetic s p Group 7A Atomic number = 9 s s p 5 ---> 9 total electrons paramagnetic s p s p s s s s s p p Neon Group 8A Atomic number = s s p 6 ---> total electrons diamagnetic Note that we have reached the end of the nd period, and the nd shell is full! Page III-6b-5 / Chapter Six Part II Lecture Notes Electron Configurations of p-block Elements

Page III-6b-6 / Chapter Six Part II Lecture Notes Sodium Group A Atomic number = s s p 6 s or "neon core" + s [Ne] s (uses noble gas notation) And: we have begun a new period! All Group A elements have [core]ns configurations. Aluminum Group A Atomic number = s s p 6 s p or [Ne] s p All * Group A elements have [core] ns np configurations where n is the period number. * some have (n-)d as well s s s p p Phosphorus Group 5A Atomic number = 5 s s p 6 s p or [Ne] s p Relationship of Electron Configuration and Region of the Periodic Table All * Group 5A elements have [core ] ns np configurations where n is the period number. * some have (n-)d also s s s p p Gray = s block Orange = p block Green = d block Violet = f block Transition Metals Fourth Period Electron Configurations All transition metals have the configuration [core]ns x (n - )d y and so are "d-block" elements. K s s p 6 s p 6 4s or [Ar]4s Ca s s p 6 s p 6 4s or [Ar]4s Sc s s p 6 s p 6 4s d or [Ar]4s d Ti s s p 6 s p 6 4s d or [Ar]4s d V s s p 6 s p 6 4s d or [Ar]4s d Cr s s p 6 s p 6 4s d 5 or [Ar]4s d 5 Mn s s p 6 s p 6 4s d 5 or [Ar]4s d 5 Note: exceptions! Chromium Iron Copper Cu s s p 6 s p 6 4s d or [Ar]4s d Page III-6b-6 / Chapter Six Part II Lecture Notes

Page III-6b-7 / Chapter Six Part II Lecture Notes Electron Configuration Anomalies Cr s s p 6 s p 6 4s d 5 or [Ar]4s d 5 Cu s s p 6 s p 6 4s d or [Ar]4s d Chromium, copper and other elements do not follow the n + l filling orders Anomalies arise from stability associated with half-filled and completely filled d-subshells. Know how n + l rule works, and know that anomalies exist on the periodic table Lanthanides and Actinides All these elements have the configuration [core]ns x (n - )d y (n - )f z and so are "f-block" elements Exceptions exist: Cerium [Xe] 6s 5d 4f [Xe] 6s 4f expected Uranium [Rn] 7s 6d 5f [Rn] 7s 5f 4 expected Transition Metals Electron Configurations Filling Order Iron: [Ar] 4s d 6 Zinc: [Ar] 4s d Technetium: [Kr] 5s 4d 5 Niobium: [Kr] 5s 4d Osmium: [Xe] 6s 4f 4 5d 6 Meitnerium: [Rn] 7s 5f 4 6d 7 notice f orbitals in 6th period & beyond Anion Configurations Cation Configurations To form anions from elements add or more e- using normal n + l rules P [Ne] s p + e- ---> P - [Ne] s p 6 To form cations from elements remove or more e- from subshell of highest n [or highest (n + l)]. P [Ne] s p - e- ---> P + [Ne] s s p s p s p s p s p s p s p s p s s s s Page III-6b-7 / Chapter Six Part II Lecture Notes

Page III-6b-8 / Chapter Six Part II Lecture Notes Ion Configurations For transition metals, remove ns electrons and then (n - ) electrons. Fe [Ar] 4s d 6 ---> Fe + [Ar] d 6 ---> Fe + [Ar] d 5 Fe Fe + Magnetic Properties Magnetic properties of ions assist us with charges DIAMAGNETIC ions have no unpaired electrons. Ions with unpaired electrons are PARAMAGNETIC. As number of unpaired electrons increases, the degree of paramagnetism also increases 4s d To form cations, always remove electrons of highest n value first! 4s 4s d Fe + d 4s 4s d d Zn + diamagnetic Mn + most paramagnetic, 5 unpaired e - s 4s d Fe + paramagnetic, 4 unpaired e - s Periodic Trends CH Q&D Guide to Periodic Trends: Atomic and ionic size: increase left and down Ionization energy and Electron affinity: increase right and up See Periodic Trends Handout CH Periodic Trends "Cheat Sheet" Larger orbitals. Electrons held less tightly. Electrons held more tightly Effective Nuclear Charge, Z* Z* is the effective nuclear charge experienced by the outermost electrons Inner electrons screen the effect of the nuclear charge from the outermost electrons Explains why E(s) < E(p) Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* by --> [ Z - (no. inner electrons) ] Charge felt by s e- in Li: Z* = - = Be Z* = 4 - = B Z* = 5 - = and so on! Effective Nuclear Charge, Z* The s electron PENETRATES the region occupied by the s electron. s electron experiences a higher positive charge than expected. Page III-6b-8 / Chapter Six Part II Lecture Notes

Area screened by s electrons Page III-6b-9 / Chapter Six Part II Lecture Notes Effective Nuclear Charge, Z* Electron cloud for s electrons Effective Nuclear Charge Z* is the nuclear charge experienced by the outermost electrons due to screening by inner electrons Atom Z* Experienced by Electrons in Valence Orbitals Li +.8 Be??? (~) B +.58 C +. N +.85 O +4.49 F +5. Increase in Z* across a period Values calculated using Slater s Rules for determining effective nuclear charge: http://www.wellesley.edu/chemistry/chem/slater.html Orbital Energies Orbital energies drop as Z* increases Atomic Size SIZE Size increases as you go down a group. Because electrons are added further from the nucleus, there is less attraction. Size increases on going right to left across a period. n = n = n = n= 4 n = 5 n = 6 r (pm) Moving down group IA, the atomic radii increase with the principle quantum number 5 5 75 Group() radii 5 6 n Moving across the rd period we see the atomic radii of the elements decrease. Atomic radii generally increase going right to left on the periodic table. Page III-6b-9 / Chapter Six Part II Lecture Notes r (pm) 9 6 5 8 8 rd period radii 4 5 7 group number

Atomic Radii Page III-6b- / Chapter Six Part II Lecture Notes Trends in Atomic Size Radius (pm) 5 K 5 nd period Li rd period Na st transition series Ne Ar Kr 5 He 5 5 5 5 4 Atomic Number Li,5 pm e and p Ion Sizes Does + the size go up or down when Li +, 6 pm e and p losing an electron to form a cation? Li,5 pm e and p Ion Sizes Li +, 6 pm e and p CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES. + Forming a cation. F,64 pm 9e and 9p Ion Sizes Does the size go up or down - when gaining an F - electron, 6 pmto form an anion? e and 9 p F,64 pm 9e and 9p Ion Sizes Forming an anion. ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. Trends in ion sizes are the same as atom sizes. - F -, 6 pm e and 9 p Page III-6b- / Chapter Six Part II Lecture Notes

Page III-6b- / Chapter Six Part II Lecture Notes Trends in Ion Sizes Redox Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg + ions and not Mg +? Why do nonmetals take on electrons? Ionization Energy Trends in Ionization Energy Mg (g) + 78 kj ---> Mg + (g) + e- Mg + (g) + 45 kj ---> Mg + (g) + e- Mg + (g) + 77 kj ---> Mg + (g) + e- Mg + has protons and only electrons. Therefore, IE for Mg + > Mg. Energy IE = energy cost required is very high to remove to dip into an electron a shell of lower from an n. atom This is in why the gas ox. no. phase. = Group no. Trends in Ionization Energy IE increases R to L across a period because Z* increases. Metals lose electrons more easily than nonmetals. Metals are good reducing agents. Nonmetals lose electrons with difficulty. Trends in Ionization Energy IE increases up a group because size decreases. Reducing ability generally increases down the periodic table (such as in the reactions of Li, Na, and K) Page III-6b- / Chapter Six Part II Lecture Notes

Lithium Page III-6b- / Chapter Six Part II Lecture Notes Periodic Trend in the Reactivity of Alkali Metals with Water Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy involved when an atom gains an electron. A(g) + e- ---> A - (g) E.A. = E Sodium Potassium Electron Affinity of Oxygen Electron Affinity of Nitrogen E is EXOthermic because O has an affinity for an e-. O atom [He] O - ion [He] + electron E is zero due to electron-electron repulsion. N atom N- ion [He] [He] + electron EA = 4 kj EA = kj Trends in Electron Affinity Trends in Electron Affinity Affinity for electron increases left to right across a period (EA becomes more negative). Affinity increases as you move up a group (EA becomes more negative). Page III-6b- / Chapter Six Part II Lecture Notes Notice: EA (F) < EA (Cl) unknown mechanism, electron repulsion? atom size?

Implications of Periodic Trends Useful in predicting reactivities, chemical formulas, etc. Page III-6b- / Chapter Six Part II Lecture Notes End of Chapter 6 Part Metals: low ionization energy, give up electrons easily Nonmetals: high electron affinity, love electrons from metals See also: Chapter Six Part Study Guide Chapter Six Part Concept Guide Page III-6b- / Chapter Six Part II Lecture Notes