6 Oct 1997Chemical Periodicity1 Electron Configurations  Chemical Periodicity (Ch 8) Electron spin & Pauli exclusion principle configurations spectroscopic, orbital box notation Hund’s rule - electron filling rules configurations of ATOMS: the basis for chemical valence configurations and properties of IONS periodic trends in : size ionization energies electron affinities Na + Cl  NaCl Mg +  O 2  MgO

6 Oct 1997Chemical Periodicity2 Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (  ) ORBITALS (m  ) Arrangement of Electrons in Atoms... Because there is a 4th quantum number, the electron spin quantum number, m s. Each orbital can be assigned up to 2 electrons! WHY ?

6 Oct 1997Chemical Periodicity3 Electron Spin Quantum Number, m s It can be proved experimentally that the electron has a spin. This is QUANTIZED. The two allowed spin directions are defined by the magnetic spin quantum number, m s m s = +1/2 and -1/2 ONLY.

6 Oct 1997Chemical Periodicity4 Electron Spin Quantum Number Diamagnetic : NOT attracted to a magnetic field All electrons are pairedN 2 MAGNETISM is a macroscopic result of quantized electron spin 5_magnet.mov Paramagnetic : attracted to a magnetic field. Substance has unpaired electronsO 2

6 Oct 1997Chemical Periodicity5 Pauli Exclusion Principle No orbital can have more than 2 electrons No two electrons in the same atom can have the same set of 4 quantum numbers (n, l, m l, m s )OR “Each electron has a unique address.” electrons with the same spin keep as far apart as possible electrons of opposite spin may occupy the same “region of space” (= orbital) Consequences:

6 Oct 1997Chemical Periodicity6 QUANTUM NUMBERS n (shell)1, 2, 3, 4,...  (subshell) 0, 1, 2,... n - 1 m  (orbital) -   m s (electron spin) +1/2, -1/2

6 Oct 1997Chemical Periodicity7 Shells, Subshells, Orbitals n  #orbitals #e - TotalPERIOD 10 s1221 (H, He) 20 s12 1 p3682 (Li…Ne) 30 s12 1 p36 3 (Na.. Ar) 2 d s12 1 p36 2 d510 3 f71432 n0..(n-1)(2  +1) 2*(2  +1) 2n 2 etc, for n = 5, 6  = 0 s  = 1 p  = 2 d  = 3 f

6 Oct 1997Chemical Periodicity8 Element Mnemonic Competition Hey! Here Lies Ben Brown. Could Not Order Fire. Near Nancy Margaret Alice Sits Peggy Sucking Clorets. Are Kids Capable ?

6 Oct 1997Chemical Periodicity9 Assigning Electrons to Atoms Electrons are assigned to orbitals successively in order of the energy. For H atoms, E = - R(1/n 2 ). E depends only on n. For many-electron atoms, orbital energy depends on both n and . E(ns) < E(np) < E(nd)...

6 Oct 1997Chemical Periodicity10 Assigning Electrons to Subshells In many-electron atom: a) subshells increase in energy as value of (n +  ) increases. 5_manyelE.mov In H atom all subshells of same n have same energy. (n +  )= 4 (n +  )= 5 b) for subshells of same (n +  ), subshell with lower n is lower in energy.

6 Oct 1997Chemical Periodicity11 2s e- spends more time close to Li 3+ nucleus than the 2p e- Therefore 2s is lower in E than 3s Effective Nuclear Charge The difference in SUBSHELL energy e.g. 2s and 2p subshells is due to effective nuclear charge, Z*. Charge felt by 2s e- of Li atom

6 Oct 1997Chemical Periodicity12 Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by an electron. Z* increases across a period owing to incomplete shielding by inner electrons. For VALENCE electrons we estimate Z* as: Charge felt by 2s e- in LiZ* = = 1 Be Z* = = 2 B Z* = = 3 and so on! Z* = [ Z - (no. of inner electrons) ]

6 Oct 1997Chemical Periodicity13 Inner shell or CORE ELECTRONS VALENCE ELECTRONS Photoelectron Spectroscopy - Measuring IE Photoelectric effect: h + A  A + + e - forms basis for DIRECT determination of IE Kinetic energy of electron = h - IE therefore: IE = h - KE(e - ) 2s 2p Ne 1s 2s2p3p 3s IE (MJ/mol) Signal Ar 1s 309

6 Oct 1997Chemical Periodicity14 Electron Filling Order (Figure 8.7)

6 Oct 1997Chemical Periodicity15 Writing Atomic Electron Configurations Two ways of writing configurations. One is called the One is called the spectroscopic notation:

6 Oct 1997Chemical Periodicity16 A second way is called the orbital box notation. One electron has n = 1,  = 0, m l = 0, m s = + 1/2 Other electron hasn = 1,  = 0, m l = 0, m s = - 1/2 Writing Atomic Electron Configurations (2)

6 Oct 1997Chemical Periodicity17 Electron Configuration tool - see “toolbox”.

6 Oct 1997Chemical Periodicity18 Lithium Group 1A Z = 3 1s 2 2s 1 Beryllium Group 2A Z = 4 1s 2 2s 2

6 Oct 1997Chemical Periodicity19 Boron Z = 5 1s 2 2s 2 2p 1 Carbon Z = 6 1s 2 2s 2 2p 2 Why not ? 

6 Oct 1997Chemical Periodicity20 Carbon Z = 6 1s 2 2s 2 2p 2 The configuration of C is an example of HUND’S RULE: the lowest energy arrangement of electrons in a subshell is that with the MAXIMUM no. of unpaired electrons Electrons in a set of orbitals having the same energy, are placed singly as long as possible.

6 Oct 1997Chemical Periodicity21 Nitrogen Z = 7 1s 2 2s 2 2p 3 Oxygen Z = 8 1s 2 2s 2 2p 4

6 Oct 1997Chemical Periodicity22 Fluorine Z = 9 1s 2 2s 2 2p 5 Neon Z = 10 1s 2 2s 2 2p 6 Note that we have reached the end of the 2nd period,... and the 2nd shell is full!

6 Oct 1997Chemical Periodicity23 GROUPS and PERIODS or “neon core” + 3s 1 [Ne] 3s 1 (uses rare gas notation) Na begins a new period. All Group 1A elements: Li Na K Rb Cs have [core] ns 1 configurations. (n = period #) Sodium Z = 11 1s 2 2s 2 2p 6 3s 1

6 Oct 1997Chemical Periodicity24 Periodic Chemical Properties 5_Li.mov 5_Na.mov 5_K.mov Li Na K Rb Cs Alkalis REACTIVITYSIZEIE (Ionization Energy) Be Mg Ca Sr Ba Alkaline Earths

6 Oct 1997Chemical Periodicity25 Alkaline Earths Metals (ns 2 ) - easily oxidized to M 2+ - less reactive than alkalis of same period reactivity:Be < Mg < Ca < Sr < Ba WHY? - Size INCREASES as  group VALENCE e - are farther from nucleus same Z * - Valence e - less tightly held Therefore valence e - are easier to remove Typical reactions / compounds Oxides: M +1/2O 2 (g)  MO (s) CaO (lime) - #5 Ind. Chem Halides: M + X 2 (g)  MX Carbonates: CaCO 3 ( limestone)  CaO + CO 2 RECALL: Solubility rules and PRECIPITATION REACTIONS Sulfates: CaSO 4.2H 2 O ( gypsum)  CaSO H 2 O (plaster-of-paris) + 3/2H 2 O

6 Oct 1997Chemical Periodicity26 Relationship of Electron Configuration and Regions of the Periodic Table f block s block p block d block

6 Oct 1997Chemical Periodicity27 Transition Metals Table 8.4 Transition metals (e.g. Sc.. Zn in the 4th period) have the configuration [argon] ns x (n - 1)d y also called “d-block” elements. CopperIronChromium 3d orbitals used for Sc - Zn

6 Oct 1997Chemical Periodicity28 To form cations from elements : remove 1 e- (or more) from subshell of highest n [or highest (n +  )]. Ion Configurations P [Ne] 3s 2 3p 3 - 3e-  P 3+ [Ne] 3s 2 3p 0

6 Oct 1997Chemical Periodicity29 Ion Configurations (2) Transition metals ions: remove ns electrons and then (n - 1)d electrons. E 4s ~ E 3d - exact energy of orbitals depend on whole configuration Fe [Ar] 4s 2 3d 6 loses 2 electrons  Fe 2+ [Ar] 4s 0 3d 6

6 Oct 1997Chemical Periodicity30 Ion Configurations (3) From the magnetic properties of ions. Ions (or atoms) with UNPAIRED ELECTRONS are: PARAMAGNETIC. Ions (or atoms) without unpaired electrons are: DIAMAGNETIC. How do we know the configurations of ions?

6 Oct 1997Chemical Periodicity31 General Periodic Trends Atomic and ionic radii : SIZE Ionization energy : E(A + ) - E(A) Electron affinity : E(A - ) - E(A)

6 Oct 1997Chemical Periodicity32 Atomic Size INCREASES down a Group Size goes UP on going down a GROUP Because electrons are added further from the nucleus, there is less attraction.

6 Oct 1997Chemical Periodicity33 Atomic Size DECREASES across a period Size goes DOWN on going across a PERIOD. Size decreases due to increase in Z*. Each added electron feels a greater and greater +ve charge.

6 Oct 1997Chemical Periodicity34 Atomic Radii

6 Oct 1997Chemical Periodicity35 Trends in Atomic Size (Figure 8.10)

6 Oct 1997Chemical Periodicity36 Sizes of Transition Elements (Figure 8.11) 3d subshell is inside the 4s subshell. 4s electrons feel a more or less constant Z*. Sizes stay about the same and chemistries are similar!

6 Oct 1997Chemical Periodicity37 Ion Sizes - CATIONS Does the size go up or down when an atom loses an electron to form a cation? CATIONS are SMALLER than the parent atoms. The electron/proton attraction goes UP so size DECREASES. Forming a cation Li, 152 pm 3 e -, 3 p + Li +, 60 pm 2 e -, 3 p

6 Oct 1997Chemical Periodicity38 F -, 136 pm 10 e -, 9 p - Does the size go up or down when gaining an electron to form an anion? Ion Sizes - ANIONS F, 64 pm 9 e -, 9 p Forming an anion ANIONS are LARGER than the parent atoms. electron/proton attraction goes DOWN so size INCREASES.

6 Oct 1997Chemical Periodicity39 Trends in Ion Sizes ANIONSCATIONS Trends in relative ion sizes are the same as atom sizes. (59 pm) (207 pm)

6 Oct 1997Chemical Periodicity40 Oxidation-Reduction Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons? - related to IE and EA

6 Oct 1997Chemical Periodicity41 Mg (g) kJ  Mg + (g) + e-[Ne]2s 1 Ionization Energy (IE) Mg (g) atom[Ne]2s Mg Energy ‘cost’ is very high to remove an INNER SHELL e - (shell of n < n VALENCE ). This is why oxidation. no. = Group no. Mg 2+ (g) kJ  Mg 3+ (g) + e-[He]2s 2 2p 5 Mg 3+ Mg + (g) kJ  Mg 2+ (g) + e-[Ne]2s 0 Mg 2+ Mg +

6 Oct 1997Chemical Periodicity42 Trends in First Ionization Energy

6 Oct 1997Chemical Periodicity43 Trends in Ionization Energy (2) IE increases across a period because Z* increases. Metals lose electrons more easily than nonmetals. Metals are good reducing agents. Nonmetals lose electrons with difficulty. IE decreases down a group Because size increases, reducing ability generally increases down the periodic table. E.g. reactions of Li, Na, K

6 Oct 1997Chemical Periodicity44 2nd IE / 1st IE Li Na K 2nd IE: A +  A ++ + e -

6 Oct 1997Chemical Periodicity45 Electron Affinity (EA) A few elements GAIN electrons to form anions. Electron affinity is the energy released when an atom gains an electron. A(g) + e -  A - (g) E.A. =  E = E(A - ) - E(A) If E(A - ) < E(A) then the anion is more stable and there is an exothermic reaction

6 Oct 1997Chemical Periodicity46 Affinity for electron increases across a period (EA becomes more negative). Atom EA (kJ) B -27 C -122 N 0 O -141 F-328 Trends in Electron Affinity (Table 8.5, Figure 8.14) Affinity decreases down a group (EA becomes less negative). F-328 Cl-349 Br-325 I-295

6 Oct 1997Chemical Periodicity47 SUMMARY Electron spin: diamagnetism vs. paramagnetism Pauli exclusion principle - allowable quantum numbers configurations spectroscopic notation orbital box notation Hund’s rule - electron filling rules configurations of ATOMS: the basis for chemical valence period 2 ; groups transition metals configurations and properties of IONS periodic trends in : size ionization energies electron affinities