Announcements EM radiation --Exam 3 Oct 3...Includes chapters 7/8/9/10 The excluded items include: 1. Classical distinction between energy and matter (p. 217) 2. Numerical problems involving the Rydberg equation (equations 7.3 and 7.4) 3. Spectral analysis in the laboratory (pp. 226-227) 4. Numerical problems involving the Heisenberg uncertainty principle (p. 231) 5. Trends among the transition elements (p. 261) 6. Trends in electron affinity (pp. 265-266) 7. Pseudo-noble gas (p. 269) 8. Lattice energy (pp. 283-285) 9. IR spectroscopy (p. 292) 10.Numerical problems involving electronegativity (p. 296) 11.Electronegativity and oxidation number (p. 297) 12.Section 11.3: MO theory and electron delocalization 13.All sections in chapter 12 except 12.3 is units called are absorbe d a wave and a particle having Quant a emitted amplitude energy involve energy changes in atoms atoms electron s frequency wavelength related by related by E = hv c =!v molecules described by Wave functions e- filling gives having quantum uunmbers e- comprising determined by Aufbau Rules described by Wave Function (Orbital) e- filling spdf electronic Core Electrons comprising Valence Electrons determined by described by which are Periodic Table which involve Oribital Energy Pauli Exclusion Principal n = 1,2,3,.. basis for Aufbau Rules Quantum Numbers Quantum Numbers Hund s Rule which summarizes Periodic Properties define s Orbital size & energy defines Angular momentum, l defines Orbital shape Magnetic ml defines Orbital orientation Spin, ms defines Electron spin I. The Periodic Law and the Periodic Table 1864 Newland Law of Octaves 1869 Dimitri Mendeleev and Lother Meyer Electronic Configuration and Periodicity Chapter 8 When the elements are arranged in order of increasing atomic mass, certain sets of chemical and physical properties recur periodically. 1913 Henry Mosely relates X-ray frequency to atomic number
When the Elements Were Discovered Quantum numbers (n,l,ml,ms) specify allowed states or orbitals which are regions of space where electrons are likely to be found around the nucleus. 1. Principal Quantum Number (n): Defines the size and energy level of the orbital. n = {1,2,3,4,...}. Also called a shell (K = 1, L = 2, M = 3, N = 4,...). 2. Angular Momemtum Quantum Number (l): Defines the shape of the orbital which is a volume in space where the electron is likely to be found. Also called a subshell. l = {0,1,2,3...up to n-1} where (0=s, 1=p, 2=d, 3=f) 3. Magnetic Quantum Number (ml): Defines the spatial orientation of an orbital of the same energy. ml = {-l, 0, +l} 4. Magnetic Spin Quantum Number (ms): Defines the orientation of electron spin. ms = {+1/2 or -1/2}. Electronic of the elements: four quantum numbers describe an electron in a ground state atom. Name Symbol Permitted Values Property principal n positive integers (1,2,3, orbital energy (size)!) angular momentum l integers from 0 to n-1 orbital shape magnetic m l integers from -l to 0 to +l orbital orientation in space spin m s +1/2 or -1/2 direction of e - spin The lowest energy (ground state) electronic of all elements are constructed by filling lowest energy orbitals sequentially in what is called the Aufbau Process. 1. Lower energy (n-quantum number) orbitals fill first. 2. Hund s Rule-orbitals fill one electron at a time before electrons are paired. 3. Pauli Exclusion Principle: No two electrons can have same 4-quantum numbers) Electrons fill the lowest energy orbitals first, 2 at a time! 4s 3p 3s 2s 1s 2p The order of filling of the orbitals can be remembered using a mnemonic device. Memorize this to help you! Chemists use spdf notation and orbital box diagrams to denote or show the ground state electronic of elements. For an Hydrogen atom orbital energy only depends on the n quantum number. For many electrons atoms the energy of an orbital or electron depends on both n and l (3s < 3p < 3d) Element H He spdf Notation 1s 1 1s 2 orbital box diagram Spin quantum number. An arrow denotes an electron with spin up (+1/2) or spindown (-1/2). n principal quantum # l quantum number # of electrons in orbital
The Pauli Exclusion principle states: No two electrons can have the same 4-quantum numbers. The spin numbers can not be the same (spin up and spin down allowed only). The order of filling of the orbitals can be remembered using a mnemonic device. Memorize how to write it out as it determines electronic structure. (n, l, m l and m s) Example: Atomic Number/Element Orbital Box Diagram Full-electronic Condensed-electronic Li 1s 2 2s 1 [He]2s 1 s-block main group d-block transition metals p-block main group Electronic using Aufbau Process Atomic Number/Element H Orbital Box Diagram Full-electronic 1s 1 Condensed-electronic 1s 1 f-block inner transition metals He Li 1s 2 1s 2 2s 1 1s 2 written with noble gas [He]2s 1 Be 1s 2 2s 2 [He]2s 2 Atomic Number/Element Orbital Box Diagram Full-electronic Condensed-electronic B 1s 2 2s 2 2p 1 [He]2s 2 2p 1 C 1s 2 2s 2 2p 2 [He]2s 2 2p 2 1s 2 2s 2 2p 3 [He]2s 2 2p 3 1s 2 2s 2 2p 4 [He]2s 2 2p 4 1s 2 2s 2 2p 5 [He]2s 2 2p 5 1s 2 2s 2 2p 6 [He]2s 2 2p 6
Odd-filling behavior here! 4th and 9th position. Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field. Diamagnetic species contain all paired electrons and is repelled by the magnetic field. Diamagnetic atoms or ions: All e - are paired. Weakly repelled in a magnetic field. Paramagnetic atoms or ions: Unpaired e - exist in an orbital Attracted to an external magnetic field. Diamagnetic all electrons paired 2p Paramagnetic unpaired electrons 2p Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field. Diamagnetic species contain all paired electrons and is repelled by the magnetic field. Magnetic field off Magnetic field on Magnetic field on When a cation is formed from an atom of a transition metal, electrons are removed first from the ns orbital, then from the (n-1)d orbital. Fe: [Ar]4s 2 3d 6 Fe 2+ : [Ar]4s 0 3d 6 or [Ar]3d 6 Fe: [Ar]4s 2 3d 6 Fe 3+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Paramagentic Diamagentic Mn: [Ar]4s 2 3d 5 Mn 2+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Paramagnetic Diamagnetic
Metals loose electrons (oxidized) to become cations. Non-metals gain electrons to become anions. The electronic of each reflects this change in the number of electrons. Na [Ne]3s 1 Ca [Ar]4s 2 Al [Ne]3s 2 3p 1 Non-metals gain electrons so that anion has a noble-gas outer electron. Na + [Ne] Ca 2+ [Ar] Al 3+ [Ne] H 1s 1 F 1s 2 2s 2 2p 5 O 1s 2 2s 2 2p 4 N 1s 2 2s 2 2p 3 Metals lose electrons so that cation has a noble-gas outer electron. H - 1s 2 or [He] F - 1s 2 2s 2 2p 6 or [Ne] O 2-1s 2 2s 2 2p 6 or [Ne] N 3-1s 2 2s 2 2p 6 or [Ne] Isoelectronic species are two different elements with the same electronic --but not the same nuclear. oxidation Na: [1s 2 2s 2 2p 6 3s 1 ] =====> Na + : [1s 2 2s 2 2p 6 ] = [Ne] oxidation Al: [1s 2 2s 2 2p 6 3s 2 3 p1 ] =====> Al 3+ : [1s 2 2s 2 2p 6 ] = [Ne] reduced N: [1s 2 2s 2 2p 3 ] =====> N 3- : [1s 2 2s 2 2p 6 ] = [Ne] reduced O: [1s 2 2s 2 2p 4 ] =====> O 2- : [1s 2 2s 2 2p 6 ] = [Ne] reduced F: [1s 2 2s 2 2p 5 ] =====> F - : [1s 2 2s 2 2p 6 ] = [Ne] Na +, Al 3+, F -, O 2-, and N 3- are all said to be isoelectronic with Ne as they have the same electronic...all subshells are filled. Metals and non-metals form ions with electronic s closest to their nearest noble gas. 1A 2A 3A 4A 5A 6A 7A 8A Metals and non-metal ions tend to form electronic states closest to their nearest noble gas. What is the spdf and condensed electron of Mg and Mg 2+? Mg 12 electrons Mg 1s 2 2s 2 2p 6 3s 2 [Ne]3s 2 Mg 2+ 1s 2 2s 2 2p 6 3s 0 [Ne]3s 0 = [Ne] What are the possible quantum numbers for the last (outermost) electron in Cl? Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s 2 2s 2 2p 6 3s 2 3p 5 2 + 2 + 6 + 2 + 5 = 17 electrons Last electron added to 3p orbital n = 3 l = 1 m l = -1, 0, or +1 m s =! or -! Using the periodic table on the inside cover of the text and give the full and condensed electrons s, partial orbital diagrams showing valence electrons, and number of inner electrons for the following elements: (a) potassium (K: Z = 19) (b) molybdenum (Mo: Z = 42) (c) lead (Pb: Z = 82) C) Is ground state F paramagenetic or diamagnetic?
(a) for K (Z = 19) full condensed orbital diagram There are 18 inner electrons. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ar] 4s 1 Use condensed electron s to write the reaction for the formation of each transition metal ion, and predict whether the ion is paramagnetic. (b) for Mo (Z = 42) 36 inner electrons and 6 valence electrons full 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 4d 5 condensed [Kr] 5s 1 4d 5 partial orbital diagram (c) for Pb (Z = 82) 78 inner electrons and 4 valence electrons. 1s full 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 2 condensed 4s 1 3d 5s 1 4d 5 5p [Xe] 6s 2 4f 14 5d 10 6p 2 4p (a) Mn 2+ (Z = 25) (b) Cr 3+ (Z = 24) (c) Hg 2+ (Z = 80) Write the electron and remove electrons starting with ns to match the charge on the ion. If the remaining has unpaired electrons, it is paramagnetic. partial orbital diagram 6s 2 6p 2 Use condensed electron s to write the reaction for the formation of each transition metal ion, and predict whether the ion is paramagnetic. Identify n and l quantum numbers for each of the following. (a) Mn 2+ (Z = 25) (b) Cr 3+ (Z = 24) (c) Hg 2+ (Z = 80) Write the electron and remove electrons starting with ns to match the charge on the ion. If the remaining has unpaired electrons, it is paramagnetic. SOLUTION: (a) Mn 2+ (Z = 25) Mn([Ar]4s 2 3d 5 ) Mn 2+ ([Ar] 3d 5 ) + 2e - paramagnetic (b) Cr 3+ (Z = 24) Cr([Ar]4s 1 3d 5 ) Cr 3+ ([Ar] 3d 3 ) + 3e - paramagnetic third shell fourth shell What neutral element has the following orbital-filling diagram? (c) Hg 2+ (Z = 80) Hg([Xe]6s 2 4f 14 5d 10 ) Hg 2+ ([Xe] 4f 14 5d 10 ) + 2e - not paramagnetic (is diamagnetic) Identify n and l quantum numbers for each of the following. Many atomic properties show periodicity and trends. 3p 4dz 2 Amount of energy to remove 1 mole e - from 1 mole of gaseous atoms or element third shell fourth shell What neutral element has the following orbital-filling diagram? Gallium = Ga Amount of energy to add 1 mole e - to 1 mole of gaseous atoms or element
Many atomic properties show periodicity and trends. Electrons in elements are categorized either as inner core electrons or valence electrons. 1) Inner core electrons : electrons filling the lower n shells of an element. They are located closer to the nucleus. 2. Outer core or VALENCE e - : those e - in the highest energy level (highest n-value). The number of valence e - is given by the Group Number in the periodic table for Group A. --Responsible for chemistry and bonding of elements forming compounds or ions (true for representative but not transition metals--more complex). Periodicity in the chemical reactivity of elements occurs because of periodicity in the electronic structure of valence electrons! Inner core electrons shield outer electrons from the positive charge of the nucleus. 1-electron outer s-orbital 2-electrons outer d-orbital 5-electrons outer p-orbital Effective nuclear charge (Z eff ) is the electrostatic force felt by the outer valence electrons taking into shielding by core electrons. Z eff = Z core e- Core Valence Configuration Element Z (p + ) Electrons Electrons **** Z effective Radius (pm) To a good approximation: effective nuclear charge, Zeff is given by: Effective Nuclear charge Z eff = Z core e- # protons # of inner nonvalence electrons Bigger Zeff means more pull or electrostatic force between nucleus and electrons. [Ne]3s 1 Na 11 10 1 1 186 [Ne]3s 2 Mg 12 10 2 2 160 [Ne]3s 2 3p 1 Al 13 10 3 3 143 [Ne]3s 2 3p 2 Si 14 10 4 4 132 [Ne]3s 2 3p 3 P 15 10 5 5 128 [Ne]3s 2 3p 4 S 16 10 6 6 127 [Ne]3s 2 3p 5 Cl 17 10 7 7 99 [Ne]3s 2 3p 6 Ar 18 10 8 8 98 [Ar]4s 1 K 19 18 1 1 227 [Ar]4s 2 Ca 20 18 2 2 197 [Ar]4s 2 3d 1 Sc 21 18 3 3 135
Increasing Atomic Radius Because of increasing effective nuclear charge across a period, atomic radii decrease across a Period. As n increases down a group so does the radius. Decreasing Atomic Radius Many atomic properties show periodicity and trends. Amount of energy to remove 1 mole e - from 1 mole of gaseous atoms or element n increases Amount of energy to add 1 mole e - to 1 mole of gaseous atoms or element Periodicity of Atomic Radius The radii of cations are smaller than their parent neutral atoms, while anions are larger than its parent. Group I Cations get smaller (greater Z eff) Anions get larger (lower Z eff) Group VIII Using only the periodic table rank each set of main group elements in order of decreasing atomic size: (a) Ca, Mg, Sr (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb Using only the periodic table rank each set of main group elements in order of decreasing atomic size: (a) Ca, Mg, Sr SOLUTION: (a) Sr > Ca > Mg (b) K > Ca > Ga (c) Rb > Br > Kr (d) Rb > Sr > Ca (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb These elements are in Group 2A(2). These elements are in Period 4. Rb has a higher n engery level and is far to the left. Br is to the left of Kr. Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.
First ionization energies of the main-group elements. First Ionization Energy Ionization energy is the minimum energy (kj/mol) required to remove an 1 mole of electrons from one mole of a gaseous atom in its ground state (!H > 0). Ranking Elements by First Ionization Energy PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE 1 : (a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs I 1 + X (g) X + (g) + e - I 1 first ionization energy PLAN: IE decreases as you proceed down in a group; IE increases as you go across a period. I 2 + X (g) X 2+ (g) + e - I 2 second ionization energy I 3 + X (g) X 3+ (g) + e - I 3 third ionization energy I 1 < I 2 < I 3 Ranking Elements by First Ionization Energy PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE 1 : (a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs The ionization energy increases dramatically when an core electron is removed from a nonvalence shell. PLAN: IE decreases as you proceed down in a group; IE increases as you go across a period. SOLUTION: (a) He > Ar > Kr (b) Te > Sb > Sn (c) Ca > K > Rb (d) Xe > I > Cs Group 8A(18) - IE decreases down a group. Period 5 elements - IE increases across a period. Ca is to the right of K; Rb is below K. I is to the left of Xe; Cs is furtther to the left and down one period. 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
Identifying an Element from Successive Ionization Energies Name the Period 3 element with the following ionization energies (in kj/mol) and write its electron : Identifying an Element from Successive Ionization Energies Name the Period 3 element with the following ionization energies (in kj/mol) and write its electron : IE 1 IE 2 IE 3 IE 4 IE 5 IE 6 1012 1903 2910 4956 6278 22,230 IE 1 IE 2 IE 3 IE 4 IE 5 1012 1903 2910 4956 6278 IE 6 22,230 PLAN: Look for a large increase in energy which indicates that all of the valence electrons have been removed. The number valence electrons is reflected in the periodic table for Group A elements...find the group with that number of valence electrons. SOLUTION: The largest increase occurs after IE 5, that is, after the 5th valence electron has been removed. Five electrons would mean that the valence is 3s 2 3p 3 and the element must be phosphorous, P (Z = 15). The complete electron is 1s 2 2s 2 2p 6 3s 2 3p 3. Main Group (or representative) metals form ionic basic oxides when reacted with oxygen while nonmetals form covalent acidic oxides with oxygen. Increasing Acidity Properties of Oxides Across a Period Covalent Oxides basic acidic Ionic Oxides Increasing Basicity 2 4 1A Li2O BeO B2O3 CO2 OF2 Na2O K2O 2A CaO 3A 4A (14) Ga2O3 GeO2 6A (16) 7A (17) SeO3 Br2O7 5 Rb2O SrO In2O3 SnO2 TeO3 I2O7 6 Cs2O BaO Tl2O3 PbO2