chm115_lecture11

chm115_lecture11 - Chemistry 115 Lecture 11 Outline Chapter...

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Unformatted text preview: Chemistry 115 Lecture 11 Outline Chapter 7 Trends in Periodic Properties Atomic size Ionization energy Metallic behavior Recitation: Many electron atoms Quantum numbers, orbitals, Electron spin Aufbau principle, Trends in atomic properties HW 5: Due Friday, Feb 18, 11pm Energy levels for the hydrogen atom The values are exactly those predicted by the Bohr model, but now we include orbitals 19-2 Ordering of Orbital Energy Levels Across Periodic Chart In many-electron atom: a) subshells increase in energy as a) value of n + l increases. value b) for subshells of same n + l, b) subshell with lower n is lower in energy. energy. 20-3 Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3 See • Explains why E(2s) < E(2p) • Z* increases across a period owing to Z* incomplete shielding by inner electrons. incomplete • Estimate Z* = [ Z - (no. inner electrons) ] Estimate • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Charge • Be Z* = 4 - 2 = 2 Be Z* •B Z* = 5 - 2 = 3 and so on! Z* Electron Filling Order Order See Figure 7.2 1s 2s 3s 4s 5s 6s H Li Be Na Mg K Ca Rb Sr Sc Ti 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 4f 5f We have a simple system for building up a periodic table of the elements 7s He B CN P O S F Ne Al Si Cl Ar V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Sn Sb Te Pb Bi I Xe Y Zr Nd Mo Tc Ru Rh Pd Ag Cd In Ir Pt Au Hg Tl Cs Ba La Hf Ta W Re Os Po At Rn Fr Ra Ac Rf Du Sg Bo Ha Me Anomalous Electron Filling Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr This pattern divides elements into blocks according to which type of orbital (which l) is filling up The location of an element by the block it occupies predicts well some of its important physical and chemical properties. 7.5 Periodic Properties Key concept about period chart: Elements in a group have similar chemical properties because they have similar outer electron configurations. Note the S and P blocks (main group elements), the D block (transition elements) and the F block (inner transition elements). Electrons can be categorized as Inner core electrons – near nucleus Outer electrons – those found in highest energy level Valence electrons – participate in bonding (usually same as outer electrons). General Periodic Trends • • • Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. Trends in Periodic Properties 1. Atomic Size: A. bigger n bigger diameter due to larger orbitals size increases going down group B. larger Zeff smaller diameter Larger + charge pulls e- in Size decreases across period Holds for S, P blocks, but not so well in D block. size small big Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3 See • • • • • • Explains why E(2s) < E(2p) Z* increases across a period owing to Z* incomplete shielding by inner electrons. incomplete Estimate Z* = [ Z - (no. inner electrons) ] Estimate Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Charge Be Z* = 4 - 2 = 2 Be Z* B Z* = 5 - 2 = 3 and so on! Z* Effective Nuclear Charge, Z* • Atom in • Li • Be •B •C •N •O •F Z* Experienced by Electrons Valence Orbitals +1.28 Increase in Z* Increase ------across a period across +2.58 +3.22 +3.85 +4.49 (Actual values are non-integer) +5.13 Orbital Energies Orbital energies “drop” as Z* increases Electrons get pulled closer to larger nuclei across chart ChemNow Screens 8.9 - 8.13, Simulations ChemNow Simulations Atomic Radii Atomic Size Size decreases across a period owing to increase Size decreases in Z*. Each added electron feels a greater and greater + charge. greater Large Small Increase in Z* Trends in Atomic Size Radius (pm) 250 K 3rd period 2nd period Na Li 1st transition series 200 150 100 Kr Ne Ar 50 He 0 0 5 10 15 20 25 30 35 40 Atomic Number Sizes of Transition Elements Sizes See Figure 7.9 Ion Sizes Ion Li,152 pm 3e and 3p Does+ size go the up+or down when up , 60 pm Li l2e and 3 p electron osing an to form a cation? to Ion Sizes Ion Li,152 pm 3e and 3p Li + , 78 pm 2e and 3 p + Forming a Forming cation. cation. • CATIONS are SMALLER than the atoms SMALLER from which they come. from • The electron/proton attraction has gone UP The and so size DECREASES. DECREASES Trends in Ion Sizes Trends See Active Figure 7.12 Redox Reactions Why do metals lose Why electrons in their reactions? Why does Mg form Mg2+ Why ions and not Mg3+? ions Why do nonmetals take on Why electrons? electrons? 2. Ionization Energy: IE = energy required to remove 1 mol of electrons from 1 mol of atoms or ions. Atom(g) ion+(g) + e- ∆ E = IE1 > 0 First IE is key to reactivity. Low IE atoms tend to form cations, while high IE atoms tend to form anions. A. IE decreases down a group because electrons are further from nucleus and easier to ionize. IE: He > Ne > Kr A. IE increases across period as atoms get smaller. Ar > Cl > Al A. IE2 > IE1, and big gaps tell you when all valence electrons have been ionized. Be: 0.90 eV, 1.76 eV, 14.85 eV, 21.01 eV. IE high low Ionization Energy IE = energy required to remove an electron from an IE atom in the gas phase. atom PLAY MOVIE Mg (g) + 738 kJ Mg+ (g) + eMg Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. atom Mg (g) + 738 kJ Mg+ (g) + e- PLAY MOVIE Mg+ (g) + 1451 kJ Mg2+ (g) + eMg has 12 protons and only 11 electrons. Ionization Energy Mg (g) + 735 kJ Mg+ (g) + eMg+ (g) + 1451 kJ Mg2+ (g) + e- PLAY MOVIE Mg2+ (g) + 7733 kJ Mg3+ (g) + eEnergy cost is very high to dip into a shell of Energy lower n. This is why ox. no. = Group no. This Trends in Ionization Energy See Active Figure 7.10 See Periodic Trend in the Reactivity of Alkali Metals with Water Lithium Sodium Potassium 3. Electron Affinity 3. A few elements GAIN electrons to few GAIN form anions. anions Electron affinity is the energy involved Electron when an atom gains an electron to form an anion. form A(g) + e- A-(g) E.A. = ∆U 1. Electron Affinity (EA) EA is the energy needed to add 1 mol e- to 1 mol atoms. Follows trends of IE. Halogens like e-, so have high EA Atoms can gain electrons to form anions: Cl + e- Cl- ∆ H = -349 kJ/mol Energy is released Energy release (binding energy) = Electron Affinity EA high low Trends in Electron Affinity See Active Figure 7.11 Trends in Electron Affinity • See Figure 7.11 and See Appendix F Appendix • Affinity for electron Affinity increases across a period (EA becomes more positive). positive). • Affinity decreases down a Affinity group (EA becomes less positive). positive). Atom EA F +328 kJ Cl +349 kJ Br +325 kJ I +295 kJ Note effect of atom Note size on F vs. Cl size 1. Metallic Behavior Metals are located at the left and lower parts of the periodic chart Metals are typically hard, shiny metals with high melting points and that conduct electricity. There are exceptions: Hg, Cs (liquids at, near RT) graphite (conducts), I (shiny). Metals tend to lose electrons and form cations: Rb Rb + + e low Metallic high Conclusions Contraction within shell with increasing nuclear charge Metallic behavior associated with ease of ionization Stronger binding within shell with increasing nuclear charge NM M Metallic behavior Add electron to form noble gas configuration Tendency to within shell to better accept e- with increasing nuclear charge Lose electron to form noble gas configuration ...
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