Matter Waves Find the wavelength of any object given v and m Orbitals Square of Schrödinger wave-function gives the probability density or electron density or orbital Orbitals give the probability of finding an electron in a given region of space (boundary surface encloses 90% of electron density) Orbitals specify the energy of the electron (differ in E, some degenerate, increase in energy and size as distance from nucleus increases) Quantum Numbers a. n must be an integer, size of orbital, distance from nucleus b. l is an integer from 0 to n- 1, defines shape (l =0, 1, 2, 3 are called s, p, d, f) c. m l is an integer from -l to + l, defines orientation in space d. What is a shell, what is a subshell? Characterized by 3 quantum numbers Dr. Lori Stepan Van Der Sluys 1 Chapter 7 Dr. Lori Stepan Van Der Sluys 2 Chapter 7
Orbital Shapes s orbitals are radially symmetric, size increases with n, 1 orbital per shell p orbitals have dumbell shape, differ in orientations, 3 degenerate orbitals per shell d orbitals have four-leaf-clover shape plus one dumbell with a doughnut, differ in orientation, 5 degenerate orbitals per shell Energy Level Diagrams H atom vs. many-electron atoms; which orbitals are degenerate? Be able to place the orbitals in order of increasing energy. Explain why the two are different. Calculate the orbital energy for a given principal quantum number n For an H atom For a many-electron atom Effective Nuclear Charge Know the definition, explain the effect of core electrons Compare Z eff for orbitals; be able to tell which orbital experiences the greatest Z eff Compare Z eff for atoms; be able to tell which atom experiences the greatest Z eff Electron Spin What is the spin quantum number, and what are the choices of values? What is the Pauli Exclusion Principal, and how does it govern the arrangement of electrons in orbitals? Dr. Lori Stepan Van Der Sluys 3 Chapter 7 Dr. Lori Stepan Van Der Sluys 4 Chapter 7
Chapter 7; Periodic Trends Electron Configurations What is Hund s rule and why does it determine the most stable configuration? Write the electron configuration of any element Write electron configurations using condensed notation (noble gas in brackets) Know the difference between core and valence electrons Recognize the special stability of a half-filled or full shell of electrons. Know which 2 families of elements have unexpected s and d orbital configurations because they prefer to have half-filled or full shells. Know how to use the periodic table to determine electron configurations. Read: BLB 2.5; 7.1 7.6 HW: BLB 2:37; 7:11, 23, 25, 27, 31, 45a-c, 47a, e, f, 53, 61, 94 Supplemental 7:1 12 Know: screening effects periodic properties atomic and ion sizes isoelectronic series electron configurations of ions ionization energies electron affinities What Bonus deadlines are Coming Up? When and where is Exam 1? What do I Bring to Exam 1?: pencils, student ID and a calculator Absolutely NO text-programmable calculators or wireless devices (will be confiscated) Dr. Lori Stepan Van Der Sluys 5 Chapter 7 Dr. Lori Stepan Van Der Sluys 6 Chapter 7
The Periodic Table - Ch. 2.5 Developed in 1869 Properties have a repeating pattern Provides atomic number, ave. atomic wt. Groups - similar properties because same Alkalai Metals - Group? Alkaline Earth Metals - Group? Chalcogens - Group? Halogens - Group? Noble gases - Group? Electron Configuration and the Periodic Table Electron configurations determine: 1. organization of the Periodic Table 2. the properties of the elements atomic size ionization energy electron affinities reactivity Properties are determined by: 1. Size ( ) and shape ( ) of orbitals where valence electrons reside 2. Atomic number (nuclear charge) Dr. Lori Stepan Van Der Sluys 7 Chapter 7 Dr. Lori Stepan Van Der Sluys 8 Chapter 7
Understanding periodic trends As n, atomic orbitals become Energy & less stable As Z (# protons), any given orbital becomes & more stable; Z eff ATOMIC SIZE Size increases going down a group WHY? as n increases orbital increases Size decreases going from left to right across a period Trade-off: Z e attraction e e repulsion WHY? number of increases, i.e., nuclear charge increases added outer electrons ineffectively effective nuclear charge so electrons are recall Coulomb s Law: column = row = Dr. Lori Stepan Van Der Sluys 9 Chapter 7 Dr. Lori Stepan Van Der Sluys 10 Chapter 7
Which is the largest element in the periodic table? Effective Nuclear Charge 1. In which subshell does an electron experience the greatest effective nuclear charge in a manyelectron atom? A) 3f B) 3p C) 3d D) 3s E) 4s Rank these elements in order of increasing size: Br B Be Dr. Lori Stepan Van Der Sluys 11 Chapter 7 Dr. Lori Stepan Van Der Sluys 12 Chapter 7
ELECTRON CONFIGURATIONS OF IONS Elements or electrons to form separate ions with filled. Na Cl 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 5 e core electrons valence electrons Na + Cl 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 [Ne] [Ar] "complete octets" Which electrons determine the chemistry of an element? Do core electrons participate in chemistry of reactions? Why or why not? Formation of Transition Metal Ions For a transition metal (TM): s electrons are close in energy to d electrons; included as part of the valence e Ag [Kr] 5s 1 4d 10 Ag + [Kr] 4d 10 Forming TM ions: remove e first (largest n) then remove e, if necessary Ions with different charges may be formed Fe [Ar] 4s 2 3d 6 Fe 2+ [Ar] [Ar] Fe 3+ Dr. Lori Stepan Van Der Sluys 13 Chapter 7 Dr. Lori Stepan Van Der Sluys 14 Chapter 7
ION SIZES cations are than parent atoms Charges of Common Ions Dr. Lori Stepan Van Der Sluys 15 Chapter 7 Na Na + 1.54Å 0.97Å anions are than parent atoms Cl Cl 0.99Å 1.81Å Dr. Lori Stepan Van Der Sluys 16 Chapter 7
Atom size increases going down family: ion size also increases Li + F Na + Cl size K + Br increases Rb + I ISOELECTRONIC SERIES Isoelectronic: Isoelectronic series: Example: O 2- F Na + Mg 2+ Al 3+ 10 electrons each: Configuration? Rank the following in order of increasing size: Ca 2+ Be 2+ Ba 2+ Z eff = Z 2 from [He] nuclear charge increases size decreases Put these ions in order of increasing size. Ca +2 S 2 K + Cl Dr. Lori Stepan Van Der Sluys 17 Chapter 7 Dr. Lori Stepan Van Der Sluys 18 Chapter 7
IONIZATION ENERGY Definition: Energy needed to an electron from an atom Periodic Trends in First Ionization Energy M(g) M + (g) + e - M + (g) M +2 (g) +e - I 1 First ionization energy I 2 Second ionization energy Outer (valence) electrons are more easily removed Example: Mg 1s 2 2s 2 2p 6 3s 2 core electrons [Ne] valence electrons I 1 = 735 kj/mol Mg + (g) [Ne] 3s 1 I 2 = 1445 kj/mol Mg +2 (g) [Ne] I 3 = 7730 kj/mol Mg +3 (g) 1s 2 2s 2 2p 5 Going down a family: I 1 (kj/mol) Li 520 Na I.E. 496 size K 419 Rb 403 Cs 376 Dr. Lori Stepan Van Der Sluys 19 Chapter 7 Dr. Lori Stepan Van Der Sluys 20 Chapter 7
Periodic Trends in First Ionization Energy Exceptions to the Trend: Requires to remove electrons from filled subshells (Mg, Ar, Zn) or half filled subshells (P) Going across the periodic table: I 1 (kj/mol) Na Mg Al Si P S Cl Ar 490 735 580 780 1060 1005 1225 1550 size I.E I 1 increases from to (with some exceptions) I n > 0; therefore IE is always endothermic Dr. Lori Stepan Van Der Sluys 21 Chapter 7 Dr. Lori Stepan Van Der Sluys 22 Chapter 7
ELECTRON AFFINITIES Periodic Trends in EA Definition: Energy needed to an electron to an atom or ion in the gas phase. Cl(g) + e Cl (g) EA = 349kJ/mol + EA: endothermic (ΔE +, energetically unfavorable) -EA: exothermic (ΔE, energetically favorable) Halogens: Want to fill subshell (p 6 ) large EA Group II metals (Be, Mg): Do not want to fill new subshell have EA Group I metals : Negative ions aren't stable but have ns 2 configuration EA Noble gases: REALLY do not want to fill new shell have values for EA Dr. Lori Stepan Van Der Sluys 23 Chapter 7 Dr. Lori Stepan Van Der Sluys 24 Chapter 7
METALLIC CHARACTER AND REACTIVITY 2M(s) +2H 2 O(l) 2M + (aq) +2OH (aq)+h 2 (g) Li Be Na Mg K Ca Rb Sr Metallic Character and Reactivity increases as ionization energy decreases NON-METALLIC CHARACTER AND REACTIVITY F 2 electron Cl 2 reactivity affinity Br 2 I 2 Non-Metallic Character and Reactivity increases as electron affinity increases (IE = energy need to form a positive ion) Noble Gases are relatively inert Dr. Lori Stepan Van Der Sluys 25 Chapter 7 Dr. Lori Stepan Van Der Sluys 26 Chapter 7
Compound Formation Metals prefer to form group 1A (alkali): 1+ cations valence e : group 2A (alkaline earth): 2+ cations valence e : transition metals: 1+, 2+, 3+ cations Why does reactivity as IE? Easier to as n Nonmetals prefer to form group 7A (halogens): 1 anions valence e : reactivity of halogens as EA group 6A (oxygen family): 2 anions valence e : Compounds consisting of only nonmetals are Compounds consisting of metals plus nonmetals are Dr. Lori Stepan Van Der Sluys 27 Chapter 7