Chapter 8: Electron Configuration & Chemical Periodicity
Chapter 8: Electron Configuration & Chemical Periodicity CH 121
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This 0 page Class Notes was uploaded by Amelia Notetaker on Tuesday December 1, 2015. The Class Notes belongs to CH 121 at University of Alabama - Huntsville taught by Pamela D Twigg (P) in Fall 2015. Since its upload, it has received 18 views. For similar materials see GENERAL CHEMISTRY I - 90514 - CH 121 - 02 in Chemistry at University of Alabama - Huntsville.
Reviews for Chapter 8: Electron Configuration & Chemical Periodicity
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Date Created: 12/01/15
Chapter 8 Lecture Notes Periodic table 0 We set up the periodic table according to electron orbital con gurations o The number of the period or row of the periodic table indicates the principal quantum number of the highest shell occupied 0 Period table is divided into regions by the type of subshell that is the highest lled orbital o No more than 2 electrons assigned to an orbital Arrangement of electrons in atoms 0 Electrons in atoms are arranged as Shells n major gt subshells l angular gt orbitals ml magnetic 0 Each orbital can be assigned no more than 2 electrons o This is tied to the existence of a 4th quantum number the electron spin quantum number ms The electron behaves as if it were spinning about a center axis thereby generating a magnetic eld whose direction depends on the direction of the spin 0 The 2 directions for the magnetic eld correspond to the 2 possible values for the spin quantum number ms where ms 12 and 12 Pauli exclusion principle 0 No two electrons in the same atom can have the same set of 4 quantum numbers 0 That is each electron has an unique address Electrons in atoms 0 When n1 then I 0 o This shell has a single orbital ls to which 2e can be assigned 0 When n2 then l01 25 orbital 2e Three 2p orbitals 6e Total 8e 0 When n4 then I 0 1 2 3 0 4s orbital 2e Three 4p orbitals 6e Five 4d orbitals 10e Seven 4f orbitals 14e Total 32e Assigning electrons to atoms 0 Electrons generally assigned to orbitals of successively higher energy 0 For H atoms E C1nquot2 E depends only on n o For manyelectron atoms energy depends on both n and l Assigning electrons to subshells o In H atoms all subshells of the same n have the same energy 0 In manyelectron atoms Subshells increase in energy as value of n increases 0 For subshells of same nl subshell with the lower n is lower in energy 0 Electron lling order 0 The n value is constant horizontally o The I value is constant vertically o N is constant diagonally Factors affecting assignment of electrons to subshells 0 Nuclear attractions Nuclear charge Z equals the number of protons in nucleus 0 Electron repulsions shielding Orbital shape 0 Nuclear attractions Higher nuclear charge increases nucleus electron attractions and lowers sublevel energy Z number of protons in the nucleus 0 H ls least stable O L2 ls most stable Effective nuclear charge 2 O O O O Shielding Reduces the attraction felt by the electron for the nucleus 2 is the effective nuclear charge experienced by the electrons Lower 2 makes it easier to remove an electron Shielding by other electrons in the same energy level Electron repulsion increases the energy of the sublevel Shielding by electrons in inner energy levels reduces Z for outer electrons Shielding explains why EZs lt E2p 2 increases across a period owing to incomplete shielding by inner electrons Estimate 2 by gt Z number of inner electrons Factors affecting assignment of electrons to subshells O Orbital shape penetration Increases nuclear attraction for 25 e over 2p e Decreases shielding of a 25 e by ls e Orbital shape causes electrons in some orbitals to penetrate close to the nucleus Penetration increases nuclear attraction and decreases shielding Splitting of levels into sublevels The lower the I value the more electrons penetrate greater nuclear attraction Lower I value more stable for a given shell n value Order of sublevel energies Sltpltdltf Writing atomic electron con gurations O 2 ways of writing them Spdf notation Orbital box notation Lithium 0 Group 1A 0 Atomic number 3 o lsquot2 Zsquot1 3 total electrons Sodium 0 Group 1A 0 Atomic number 11 o lsquot2 25quot2 2pquot6 3squot1 Or quotneon corequot 3squot1 o Ne3squot1 Rare gas notation Note that we have begun a new period c All group 1A elements have corensquot1 con gurations Transition metals 0 All transition metals have the con guration corensquot2 n1dquotx and so are dblock elements lanthanides and actinides o All these elements have the con guration corensquot2 n1dquotx n 2fquoty and so are f block members Electron con guration and group 0 Elements in the same group of the periodic table have the same outer electron con guration 0 Elements in the same group of the periodic table exhibit similar chemical behavior 0 Categories of electrons 0 Inner core electrons are those an atoms has in common with the previous noble gas and any complete transition series 0 Outer electrons are those in the highest energy level highest n value 0 Valence electrons those involved in forming compounds 0 General periodic trends 0 Atomic and ionic size 0 O Ionization energy Electron af nity Higher effective nuclear charge electrons held more tightly Larger orbitals electrons held less tightly Effective nuclear charge 0 O O In a manyelectron atoms electrons are both attracted to the nucleus and repelled by other electrons The nuclear charge that an electron experiences depends on both factors 2 ZS Z is the atomic number and S is a screening constant usually close to the number of inner electrons The 25 electron penetrates the region occupied by the ls electron 25 electron experiences higher positive charge than expected 0 Atomic size 0 0 Size goes up on going down a group Electrons are added further from the nucleus there is less attraction Size goes down on going across a period 0 Because of 2 Each added electron feels a greater and greater positive charge 0 Size of transition elements 0 O 0 3d subshell is inside the 4s subshell 4s electrons feel a more or less constant 2 Sizes stay about the same and chemistries are similar lon con gurations 0 To form cations from elements 0 0 Remove 1 or more e from subshell of highest n or highest nl For transition metals removes ns electrons and then n1 electrons Beware electrons are taken out in a different order than they are lled in lon sizes 0 Cations are smaller than the atoms from which they come 0 The electronproton attraction Zeff has gone up and so size decreases o Anions are larger than the atoms from which they come from o The electronproton attraction has gone down so size increases Sizes of ions 0 In an isoelectronic series ions have the same number of electrons 0 But the nuclear charge increases across a period 0 In an isoelectronic series ionic size decreases with an increasing nuclear charge lon con gurations o Diamagnetic sample unaffected by magnetic eld No unpaired electrons o Paramagnetic sample attracted to the magnet Due to unpaired electrons Ionization energy 0 Amount of energy required to remove an electron from the ground state of a gaseous atom or ion 0 More energy to take an electron every time because 2 goes up and the other remaining electrons are pulled closer 0 Once all of the electrons are removed on the outer shell the ionization energy takes a quantum leap Trends in ionization energy 0 Ionization energy increases across period because 2 increases 0 Metals are good reducing agents 0 Nonmetals lose electrons with dif culty 0 Atoms with high ionization energy tend to form anions Except noble gases 0 Atoms with low ionization energy tend to form cations 0 As you go down the column less energy is required to remove the rst electron For atoms in the same group 2 is essentially the same 0 From left to right row it gets harder and harder to remove electrons because 2 increases 0 2 discontinuities The groups 2A and 3A Electrons removed from p orbital rather than 5 orbital Electron is farther from the nucleus Small amount of repulsion by s electron Occurs between groups 5A and 6A Electron removed comes from doubly occupied orbital Repulsion from other electrons in orbital helps in its removal 0 Electron af nity 0 Few elements gain electrons to form anions 0 Electron af nity is the energy involved when an atom gains an electron to form an anion o Atoms with a low electron af nity tend to form cations o Atoms with a high electron af nity tend to form anions Electron af nity of oxygen 0 Delta E is exothermic because O has an af nity for an electron Electron af nity of nitrogen 0 Delta E is zero for nitrogen due to electronelectron repulsion Trends in electron af nity o In general electron af nity becomes more exothermic as you go from left to right across a row 0 There are again however 2 discontinuities in this trend Occurs between groups 1A and 2A 0 Added electron must go in porbital not sorbital Electron is further from nucleus and feels repulsion from selectrons Occurs between groups 4A and 5A 0 Group SA has no empty orbitals Extra electron must go into occupied orbital creating repulsion o Reactive nonmetals have high ionization energy and high negative electron af nities These elements attract electrons strongly and tend to form negative ions in ionic compounds 0 Reactive metals have low ionization energy and slightly negative electron af nities These elements lose electrons easily to tend to form positive ions in ionic compounds 0 Noble gases have very high ionization energy and slightly positive electron af nities These elements tend to neither lose nor gain electrons Trends in 3 atomic properties 0 Atomic size Increases from right to left and up to down 0 Ionization energy Increases down to up and left to right 0 Electron af nity Increases with many exceptions left to right and down to up Metals vs nonmetals o Metals Have shiny luster various colors but most are silver 0 Good conductors of heat and electricity Most metal oxides are ionic solids that are basic Tend to form cations in aqueous solutions 0 Nonmetals No luster but various colors Solids are brittle some soft or hard 0 Poor conductors Most nonmetal oxides are molecular substances that form acidic solutions Tend to form anions or oxyanions in aqueous solutions
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