2008 Brooks/Cole 2. Frequency (Hz)

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1 Electromagnetic Radiation and Matter Oscillating electric and magnetic fields. Magnetic field Electric field Chapter 7: Electron Configurations and the Periodic Table Traveling wave moves through space like the ripples on a pond All types ( colors ) have the same velocity (through a vacuum). c = speed of light = x 10 8 ms -1 (exact) 2008 Brooks/Cole Brooks/Cole 2 Electromagnetic Radiation and Matter + λ Electromagnetic Radiation and Matter E increases from radio waves (low ν, long λ) to gamma rays (high ν, short λ) Amplitude 0 distance Frequency (Hz) γ-rays X-rays UV IR Microwave Radiowave FM AM Long radio waves Wavelength (m) Atom Virus Bacterial Animal Thickness Width Dog cell cell of a CD of a CD Wavelength (nm) Visible light is a very small portion of the entire spectrum 2008 Brooks/Cole Brooks/Cole 4 Planck s Quantum Theory Heated solid objects emit visible light Intensity and color distribution depend on T Planck s Quantum Theory As T, the wavelength of maximum intensity shifts toward the blue Increasing filament T 2008 Brooks/Cole Brooks/Cole 6 1

2 Planck s Quantum Theory Classical theory: no restriction on the E emitted by hot atoms. didn t fit experimental data. The Photoelectric Effect Light can cause ejection of e - from a metal surface. Metal cathode (-) An anode (+) attracts e -. Current is measured. vacuum window Anode (+) 2008 Brooks/Cole Brooks/Cole 8 The Photoelectric Effect If λ > threshold (E too low), no e - emission. Higher intensity does NOT cause e - emission if E < threshold! Thresholds: metal λ (nm) color Cs 579 yellow K 539 green Na 451 blue Li 428 violet current (# of ejected e - ) increasing E threshold High I low I The Photoelectric Effect Imagine photons (balls) hitting e - embedded in glue: If the E of the ball: is low, it can t eject an e -. exceeds the strength of the glue, an e - is released Higher intensity = more photons (balls). If E > threshold, more balls eject more e Brooks/Cole Brooks/Cole 10 The Photoelectric Effect Light waves waves match waves cancel + = + = The Bohr Model of the Hydrogen Atom Heated solid objects emit continuous spectra. Excited atomic gases emit line spectra. Each element has a unique pattern. Hydrogen, H wavelength (nm) Mercury, Hg wavelength (nm) 2008 Brooks/Cole Brooks/Cole 12 2

3 The Bohr Model of the Hydrogen Atom Neils Bohr (1913): Energy The Bohr Model of the Hydrogen Atom visible emission ir emission n 3 absorption 2 absorption: ΔE > 0, n emission: ΔE < 0, n Bohr s model exactly predicts the H-atom spectrum. E = x J 1 n = 1, 2, 3,... n 2 ultraviolet emission wavelength (nm) 2008 Brooks/Cole Brooks/Cole 14 The Bohr Model of the Hydrogen Atom H-atom transitions: ΔE = x J n f 2 n i 2 The Bohr Model of the Hydrogen Atom Calculate the E and wavelength (in nm) for an H-atom n = 4 n = 2 transition. ΔE = x [(½) 2 (¼) 2 ] J = x J = x J (negative sign omitted. Losing energy = emission) ν = ΔE/h = x J /6.626 x Js = x s -1 = x Hz λ = c/ν = x 10 8 ms -1 /6.166 x s -1 λ = x 10-7 m = nm 2008 Brooks/Cole Brooks/Cole 16 Bohr s model predicts the H-atom: λ = h mv λ = wavelength (m) h = Planck s constant (J s) m = mass (kg) v = velocity (m s -1 ) Davidson and Germer (1927) observed e - diffraction by metal foils. Wave-like behavior! 2008 Brooks/Cole Brooks/Cole 18 3

4 Schrödinger equation (1926): Treats e - as standing waves (not particles). Developed by analogy to classical equations for the motion of a guitar string. Called wave mechanics or quantum mechanics Explains the structure of all atoms and molecules. Complicated math; important results Brooks/Cole Brooks/Cole 20 ψ 2 = probability of finding an e - at a point in space. Each ψ describes a different energy level. An electron density (probability) map plots ψ 2 for each point in space. Bigger value = darker shade. The H-atom ground-state orbital 2008 Brooks/Cole Brooks/Cole 22 Quantum Numbers Quantum Numbers Most important in determining the orbital energy. Defines the orbital size. Orbitals with equal n are in the same shell. l Code s p d f g h Brooks/Cole Brooks/Cole 24 4

5 Quantum Numbers l = 0, or 1, or 2 if n = 3 and l = 2 (3d), m l is -2, -1, 0, 1 or 2. if n = 3 and l = 1 (3p), m l is -1, 0, or 1. if n = 3 and l = 0 (3s), m l must be 0. Every (n, l, m l ) set has a different shape and/or orientation. Quantum Numbers Number of Number of Maximum Electron Subshell Orbitals Electrons Electrons Shell type Available Possible for n th Shell (n) (=2l + 1) in Subshell (=2n 2 ) 1 s s 1 2 p s 1 2 p 3 6 d s 1 2 p 3 6 d 5 10 f s 1 2 p 3 6 d 5 10 f 7 14 g* Brooks/Cole Brooks/Cole 26 Electron Spin Experiments showed a 4 th quantum no. was needed +½ or ½ only. s Orbitals l = 0 orbital: Every shell (n level) has one s orbital. Spherical. Larger n value = larger sphere View an e - as a spinning sphere. Spinning charges act as magnets. 1s 2s 3s Probability of finding e - at distance r from nucleus Distance from nucleus, r (pm) 2008 Brooks/Cole Brooks/Cole 28 p Orbitals d Orbitals Three p orbitals (l = 1): Related to m l = -1, 0, +1. p x, p y and p z Five d orbitals (l = 2): 3d xz 3d xy 3d yz 3d x 2 - y 2 3d z Brooks/Cole Brooks/Cole 30 5

6 Hydrogen Atom Energies E = x n 2 (in J/atom). Many-Electron Atoms All other atoms are more complex. Subshells do not have equal E. Energy 4s 3s 2s 4p 3p 2p 3d Note: E 3s E 3p E 3d. But, E 3px = E 3py = E 3pz 1s Also: 3d is above 4s (but E 4s E 3d ) 2008 Brooks/Cole Brooks/Cole 32 Atom Electron Configurations Add e - to orbitals in increasing E order Paired spins: +½ and -½ ; Unpaired spins: all +½ or -½ ; or H He Li Be B C N O F Ne Atom Electron Configurations 1s 2s 2p Electron configurations Expanded Condensed 1s 1 1s 1 1s 2 1s 2 1s 2 2s 1 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 1 1s 2 2s 2 2p 1 2p 1 1s 2 2s 2 2p 2 1s 2 2s 2 2p 1 2p 1 2p 1 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 2p 1 2p 1 1s 2 2s 2 2p 4 1s 2 2s 2 2p 2 2p 2 2p 1 1s 2 2s 2 2p 5 1s 2 2s 2 2p 2 2p 2 2p 2 1s 2 2s 2 2p 6 Energy 3s 2s 1s 3p 2p 2008 Brooks/Cole Brooks/Cole 34 Atom Electron Configurations Increasing (n + l), then increasing n n value l value s 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 5d 5f n + l = 8 4p 4d 4f n + l = 6 n + l = 7 3p 3d n + l = 4 n + l = 5 2p n + l = 2 n + l = 3 n + l = 1 Atom Electron Configurations 1s 2 s 3 s 4 s 3d 5 s 4d 6 s 5d 7 s 6d Main group s block 4 f 5 f Lanthanides and actinides f block 3d 4d 5d 6d Transition elements d block Block identities show where successive e - add. Note: d steps down, f steps down again. 2p 3p 4p 5p 6p 7p 1s Main group p block 2008 Brooks/Cole Brooks/Cole 36 6

7 Atom Electron Configurations Valence Electrons The first 20 elements have only s and p e - : 2008 Brooks/Cole Brooks/Cole 38 Valence Electrons atom configuration core valence N 1s 2 2s 2 2p 3 [He] 2s 2 2p 3 Si 1s 2 2s 2 2p 6 3s 2 3p 2 [Ne] 3s 2 3p 2 Se 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 4 [Ar] 3d 10 4s 2 4p 4 Mn 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2 [Ar] 3d 5 4s 2 Note: # of valence e - = A group # Valence Electrons Lewis dot symbols: Dots represent valence e -. Usually only used for s- and p-block elements. N 2008 Brooks/Cole Brooks/Cole 40 Valence Electrons Electron Configurations of Transition Metals 1A 2A 3A 4A 5A 6A 7A 8A ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 Li Be B C N O F Ne Na Mg Al Si P S Cl Ar Sn (5 th period, group 4A). Noble-gas core: [Kr] Complete 5s and 4d. 2 e - into 5p: [Kr] 4d 10 5s 2 5p Brooks/Cole Brooks/Cole 42 7

8 Electron Configurations of Transition Metals Electron Configurations of Transition Metals [Kr] 4d 5 5s 2 Sc 3d 1 4s 2 Ti 3d 2 4s 2 V 3d 3 4s 2 Fe 3d 6 4s 2 Co 3d 7 4s 2 Ni 3d 8 4s 2 Y 4d 1 5s 2 Zr 4d 2 5s 2 Nb 4d 4 5s 1 Ru Rh 4d 7 5s 1 4d 8 5s 1 Ni (4 th ; 8B; 8 e - into 3d) [Ar] 3d 8 4s 2 La 5d 1 6s 2 Hf 5d 2 6s 2 Ta 5d 3 6s 2 W 5d 4 6s 2 Os 5d 6 6s 2 Ir 5d 7 6s 2 Pt 5d 9 6s 1 Cu (4 th ; 1B; 9 e - into 3d) [Ar] 3d 9 4s 2 Cu has lower E with filled d-block and half-filled s-block [Ar] 3d 10 4s Brooks/Cole Brooks/Cole 44 Ion Electron Configurations Same approach. Positive ion: remove one e - for each + Negative ion: add one e - for each - Ion Electron Configurations A-group ions usually adopt the nearest noble-gas configuration many ions are isoelectronic Brooks/Cole Brooks/Cole 46 Transition Metal Ions Paramagnetism & Unpaired Electrons Spinning e - = tiny magnet. If all e - are paired: Fe [Ar] 3d 6 4s 2 Fe 2+ [Ar] 3d 6 Fe 3+ [Ar] 3d 5 Mn [Ar] 3d 5 4s 2 Mn 2+ [Ar] 3d 5 Mn 4+ [Ar] 3d 3 Mn 7+ [Ar] 2008 Brooks/Cole Brooks/Cole 48 8

9 Paramagnetism & Unpaired Electrons Paramagnet Ferromagnet Periodic Trends: Atomic Radii Estimate of atomic size ½(homonuclear bond length) Cl = 100 pm (Cl 2 bond = 200 pm) H = 37 pm (H 2 bond =74 pm) Cl 200 pm Cl Radii are additive. HCl has a ( ) = 137 pm bond 100 pm 2008 Brooks/Cole Brooks/Cole 50 Periodic Trends: Atomic Radii Periodic Trends: Atomic Radii Increasing Size p + add to the nucleus. Larger charge pulls all e - in, shrinking the atom. Increasing Size 2008 Brooks/Cole Brooks/Cole 52 Periodic Trends: Ionic Radii Group 1A Group 2A Group 3A Li Li + Be Be 2+ B B Na Na + Mg Mg 2+ Al Al K K + Ca Ca 2+ Ga Ga Rb Rb + Sr Sr 2+ In In Periodic Trends: Ionic Radii Group 6A Group 7A O O 2- F F S S 2- Cl Cl Se Se 2- Br Br Te Te 2- I I Main block: outer shell completely removed. e - /e - repulsion reduced (fewer e - ). More e - /e - repulsion (more e - ). The shell swells Brooks/Cole Brooks/Cole 54 9

10 Periodic Trends: Ionic Radii Periodic Trends: Ionization Energies Isoelectronic Ions O 2- F - Na + Mg 2+ Ionic radius (pm) Number of protons Number of electrons Increasing nuclear charge decreasing size Mg(g) Mg + (g) + e - ΔE = Ionization Energy Also called IE Brooks/Cole Brooks/Cole 56 Periodic Trends: Ionization Energies Periodic Trends: Ionization Energies Mg(g) Mg + (g) + e - ΔE = IE 1 Mg + (g) Mg 2+ (g) + e - ΔE = IE 2 Down a group: IE. Larger atom = less tightly held e - Across a period: IE. Smaller atom = more tightly held e - IE 2 > IE 1 Mg + holds the 2 nd e - more tightly. Huge increase if e - removal breaks a complete shell (the core) Brooks/Cole Brooks/Cole 58 Periodic Trends: Ionization Energies Table 7.8 Ionization Energies Required to Remove Successive Electrons Ionization Energy Li Be B C N O F Ne (MJ/mol) 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2 1s 2 2s 2 2p 3 1s 2 2s 2 2p 4 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 IE IE IE IE IE IE IE IE 8 Core electrons IE IE Periodic Trends: Electron Affinities F(g) + e - F - (g) ΔE = Electron Affinity Usually exothermic (EA is negative) EA increases from left to right. Metals have low EA; nonmetals have high EA. Some tables list positive numbers. a sign-convention choice Brooks/Cole Brooks/Cole 60 10

11 Periodic Trends: Electron Affinities Table 7.9 Electron Affinities (kj/mol) 1A 2A 3A 4A 5A 6A 7A 8A (1) (2) (13) (14) (15) (16) (17) (18) H -73 Li -60 Na -53 K -48 Rb -47 Be Mg Ca -2 Sr -5 B -27 Al -43 Ga -30 In -30 C -122 Si -134 Ge -119 Sn -107 N P -72 As -78 Sb -103 O -141 S -200 Se -195 Te -190 F -328 Cl -349 Br -325 I -295 He Ne Ar Kr Xe Ion Formation and Ionic Compounds Ionic compound formation is very exothermic: K(s) + ½ F 2 (g) KF(s) Two, of several, steps are: ΔH f = kj K K + + e - IE = 419 kj [Ne] 4s 1 [Ne] F + e - F - EA = -328 kj [He] 2s 2 2p 5 [He] 2s 2 2p 6 = [Ne] 2008 Brooks/Cole Brooks/Cole 62 Energy in Ionic Compound Formation K(g) F(g) ΔH 1 = ΔH sub ΔH 3 = IE ΔH 4 = EA K + (g) + F - (g) Energy in Ionic Compound Formation Lattice energies are hard to measure K(g) F(g) ΔH 1 = +89 kj ΔH 3 = +419 kj ΔH 4 = -328 kj K + (g) + F - (g) ΔH 2 = ½ Bond E ΔH 5 = Lattice E ΔH 2 = +79 kj Lattice Energy K(s) + ½ F 2 (g) ΔH f KF(s) K(s) + ½ F 2 (g) ΔH f = kj KF(s) ΔH f = ΔH 1 + ΔH 2 + ΔH 3 + ΔH 4 + ΔH 5 (H is a state function) 2008 Brooks/Cole 63 Lattice Energy = ΔH f - ΔH 1 - ΔH 2 - ΔH 3 - ΔH 4 = (-328) = -826 kj 2008 Brooks/Cole 64 Energy in Ionic Compound Formation Table 7.10 Effect of Ion Size and Charge on Lattice E and Melting Point Compound Charges of Ions r + + r - Lattice Energy Melting Point (kj/mol) (K) NaCl +1, = 283 pm BaO +2, = 275 pm MgO +2, = 212 pm Brooks/Cole 65 11

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