Shells of an atom contain a number of stacked orbitals

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1 1 Shells of an atom contain a number of stacked orbitals The orbitals have different energies and for the d and f orbitals, the energies overlap s-orbital energies in the next principal level. 8 s p d f 7 6 1

2 8 7 6 d f p 1 s s p d f Electrons fill the orbitals much like this upside down hotel del orbital. There is a specific sequence on the filling order of the occupancy. 1 1s H = 1s 1 He = 1s Li = 1s s 1 Be = 1s s B = 1s s p 1 C = 1s s p N = 1s s p O = 1s s p F = 1s s p Ne = 1s s p 6 s + + p 7

3 H 1s s p He 1s s p 8 Aufbau Degenerate Pauli Exclusion Principle Hund s Rule Magnetism Diamagnetism Paramagnetism Pseudo Noble gas Core electrons Isoelectronic 9 Li = + + 1s s p Be = + + 1s s p 10

4 + + B = 1s s p + + C = 1s s p + + N = 1s s p O = 1s s p + + F = 1s s p + + Ne = 1s s p 1 Pictorial view of the atomic orbitals arrangements for the first 10 elements of the periodic table. 1

5 1s 1 H = 1s 1 He = 1s Li = 1s s 1 Be = 1s s B = 1s s p 1 C = 1s s p N = 1s s p O = 1s s p F = 1s s p Ne = 1s s p 6 Na = 1s s p 6 s 1 Mg = 1s s p 6 s Al = 1s s p 6 s p 1 Si = 1s s p 6 s p P = 1s s p 6 s p 1 S = 1s s p 6 s p Cl = 1s s p 6 s p = 1s s p 6 s p 6 s + + p s + + p Core s K = [] s 1 Ca = [] s Sc = [] s d 1 Ti = [] s d V = [] s d Cr = [] s 1 d Mn = [] s d Fe = [] s d 6 Co = [] s d 7 Ni = [] s d 8 Cu = [] s 1 d 10 Zn = [] s d 10 Ga = [] d 10 s p 1 Ge = [] d 10 s p As = [] d 10 s p Se = [] d 10 s p Br = [] d 10 s p Kr = [] d 10 s p 6 d + + p p s d Why is the d orbital higher in energy than the s orbital? Understanding the fundamental reason as to why the s orbital is lower in energy than the d orbital explains why the s orbital is filled first in the electron configuration for K: K = [Ne] s 1 not K=[] d 1 This concept also explains the electron configuration for all transition metal elements. For example Sc: Sc = []s d 1 and not Sc = [] d. 1 16

6 8 7 6 d s 1 s 17 d 18 19

7 6 7 1 IA 1 H 1s 1 Li s 1 Na s 1 K s 1 Rb s 1 Cs 6s 1 Fr 7s 1 ne - number of electrons; the total number of electrons va - valence electrons; the number of valence electrons s - shell of valence electrons p - previous noble gas IIA Be s Mg s Ca s Sr s Ba 6s Ra 7s IIIB Sc d 1 Y d 1 La d 1 Ac 6d 1 IVB Ti d Zr d Hf d Db 6d VB V d Nb d Ta d 6 VIB 7 VIIB Cr Mn s 1 d d Mo Tc s 1 d d W Re 6s 1 d d Jl Rf Bh 6d 7s 1 6d 6d 8 9 VIIIB Fe Co d 6 d 7 Ru Rh d 6 d 7 Os d 6 Hn 6d 6 Ir d 7 Mt 6d IB Ni d 8 Ni d 8 Ni d 8 1 IIB Cu Zn s 1 d 10 d 10 Ag Cd s 1 d 10 d 10 Au Hg 6s 1 d 10 d 10 1 IIIA B p 1 Al p 1 Ga p 1 In p 1 Tl 6p 1 1 IVA C p Si p Ge p Sn p Pb 6p 1 VA N p P p As p Sb p Bi 6p 16 VIA O p S p Se p Te p Po 6p 17 VIIA F p Cl p Be p I p At 6p 18 VIIIA He 1s Ne p 6 p 6 Kr p 6 Xe p 6 Rn 6p IA H 1s 1 Li s 1 Na s 1 K s 1 Rb s 1 Cs 6s 1 Fr 7s 1 IIA Be s Mg s Ca s Sr s Ba 6s Ra 7s IIIB Sc d 1 Y d 1 La d 1 Ac 6d 1 IVB Ti d Zr d Hf d Db 6d VB V d Nb d Ta d Jl 6d 6 VIB 7 VIIB Cr s 1 d Mn d Mo Tc s 1 d d W 6s 1 d Rf 7s 1 6d Re d Bh 6d 8 9 VIIIB Fe d 6 Ru d 6 Os d 6 Hn 6d 6 Co d 7 Rh d 7 Ir d 7 Mt 6d IB Ni d 8 Ni d 8 Ni d 8 Cu s 1 d 10 Ag s 1 d 10 Au 6s 1 d 10 1 IIB Zn d 10 Cd d 10 Hg d 10 1 IIIA B p 1 Al p 1 Ga p 1 In p 1 Tl 6p 1 1 IVA C p Si p Ge p Sn p Pb 6p 1 VA N p P p As p Sb p Bi 6p 16 VIA O p S p Se p Te p Po 6p 17 VIIA F p Cl p Be p I p At 6p 18 VIIIA He 1s Ne p 6 p 6 Kr p 6 Xe p 6 Rn 6p 6 e - config for Sulfur S = [Ne]s p ne - number of electrons; the total number of electrons; this equals the number of protons or atomic number = 16 1 The 16 electrons for sulfur occupy the shells & orbitals of sulfur from the lowest energy to the highest. e - config for Sulfur S = [Ne]s p s p d f 16 total electrons S 1

8 8 7 6 s p d f e - config for Sulfur Tc = [Kr]s d total electrons Te Atomic Structure and Electronic 1 Configuration

9 What are ions? Charge atoms or group of atoms. cations are formed by losing electrons (+) lose electrons M+ + e - Cation M + e - M- Anion (-) gain electrons anions are formed by gaining electrons 6 Na + cation is form by loss of an electron from sodium Na Na + 11 p 1 n e- 8e- 1e- Na = [Ne]s 1 + 1e - 11 p 1 n e- 8e- 0e- 7 F 9 p 10 n e- 7e- F = [He] s p + 1e - 9 p 10 n e- 8e- F - 8

10 The 16 electrons for sulfur occupy the shells & orbitals of sulfur from the lowest energy to the highest. e - config for Sulfur S - = [Ne]s p 6 s p d f total electrons S What is the electron configuration for transition metal ions? Why is the electrons in the s-shell removed before the electrons in the d shell, even though the d-electrons are higher in energy? d s s d 1

11 6 7 1 IA 1 H 1s 1 Li s 1 Na s 1 K s 1 Rb s 1 Cs 6s 1 Fr 7s 1 IIA Be s Mg s Ca s Sr s Ba 6s Ra 7s ne - number of electrons; the total number of electrons va - valence electrons; the number of valence electrons s - shell of valence electrons IIIA IVA VA VIA p - previous noble gas IIIB Sc d 1 Y d 1 La d 1 Ac 6d 1 IVB Ti d Zr d Hf d Db 6d VB V d Nb d Ta d Jl 6d 6 VIB 7 VIIB Cr Mn s 1 d d Mo Tc s 1 d d W Re 6s 1 d d Rf 7s 1 6d Many anomalies Bh 6d 8 9 VIIIB Fe d 6 Ru d 6 Os d 6 Hn 6d 6 Co d 7 Rh d 7 Ir d 7 Mt 6d IB Ni d 8 Ni d 8 Ni d 8 1 IIB Cu Zn s 1 d 10 d 10 Ag Cd s 1 d 10 d 10 Au Hg 6s 1 d 10 d 10 B p 1 Al p 1 Ga p 1 In p 1 Tl 6p 1 C p Si p Ge p Sn p Pb 6p N p P p As p Sb p Bi 6p O p S p Se p Te p Po 6p 17 VIIA F p Cl p Be p I p At 6p 18 VIIIA He 1s Ne p 6 p 6 Kr p 6 Xe p 6 Rn 6p 6

12

13 1 Periodic Trends from the Periodic Table Periodic Trends from the Periodic Table Periodic Trends from the Periodic Table

14 Periodic Trends from the Periodic Table A map of the building block of matter. Periodic Trends from the Periodic Table 6 Periodic Trends from the Periodic Table

15 1 IA 1 H 1s Li s 1 Na s 1 K s 1 Rb s 1 Cs 6s 1 Fr 7s 1 IIA Be s Mg s Ca s Sr s Ba 6s Ra 7s IIIB Sc d 1 Y d 1 La d 1 Ac 6d 1 IVB Ti d Zr d Hf d Db 6d VB V d Nb d Ta d 6 VIB 7 VIIB Cr Mn s 1 d d Mo s 1 d W 6s 1 d Jl Rf 6d 7s 1 6d Tc d Re d Bh 6d 8 9 VIIIB Fe d 6 Ru d 6 Os d 6 Hn 6d 6 Co d 7 Rh d 7 Ir d 7 Mt 6d IB Ni d 8 Ni d 8 Ni d 8 1 IIB Cu Zn s 1 d 10 d 10 Ag Cd s 1 d 10 d 10 Au Hg 6s 1 d 10 d 10 1 IIIA B p 1 Al p 1 Ga p 1 In p 1 Tl 6p 1 1 IVA B C p 1 Si p Ge p Sn p Pb 6p 1 VA N p P p As p Sb p Bi 6p 16 VIA O p S p Se p Te p Po 6p 17 VIIA F p Cl p Be p I p At 6p 18 VIIIA He 1s Ne p 6 p 6 Kr p 6 Xe p 6 Rn 6p 6 7 Periodic Trends from the Periodic Table 8 Periodic Trends from the Periodic Table 9 Periodic Trends from the Periodic Table

16 Z eff - The net charge an electron experience due to its environment. Two factors contribute to the degree of stability of an electron: 1. Proton - electron attraction: Coulombic attraction: E = k [Q 1 Q ] / r. electron - electron repulsion Shielding phenomena Together these two factors contribute to the effective nuclear charge that an electron experience. Z eff = Z - σ Z eff - Effective nuclear charge Z - Total number of protons in nucleus σ - # of shielding electrons 10 Periodic Trends from the Periodic Table What factors determine the distance between the valence electron from the nucleus? Consider the following: p and 1-e. #Ve/#p =1/1 = 1 1-p and -e. #Ve/#p = /1 = -p and 1-e. #Ve/#p = 1/ = 0. -p and -e. #Ve/#p = / = 1 11 Periodic Trends from the Periodic Table What factors determine the distance between the valence electron from the nucleus? Consider the following: p and 1-e. #Ve/#p =1/1 = 1 Closer e - is to nucleus greater the attraction. 1-p and -e. #Ve/#p = /1 = Two e - in the same orbitals will tend to repel each other. The larger the size -p and 1-e. #Ve/#p = 1/ = 0. Greater number of protons the stronger the attraction. The smaller the size -p and -e. #Ve/#p = / = 1 Electrons in the inner core shield electrons in the outer orbitals. New orbital, larger size. 1 Periodic Trends from the Periodic Table 1

17 What factors determine attraction and therefore the extent (distance) a valence electron is from the nucleus (protons)? 1 Periodic Trends from the Periodic Table As Se 1 Periodic Trends from the Periodic Table 1 Periodic Trends from the Periodic Table

18 Ions: Consider: anion atom cation Si - Si Si + # of protons : Shielding e - : Z eff : Valence e - : All else being equal, the number of electrons in the valence shell determines the size of the specie. e - are attracted to Zeff =, but e - do have significant e- e repulsion. e - are attracted to Zeff =, with e - having some e-e repulsion. e - are attracted to Zeff =, but with only e - having the lowest e-e repulsion. 16 Periodic Trends from the Periodic Table Ions Consider: Ca + K + Cl - S - # p: Core e - : Z eff : Valence e - : V e / Z eff 8/10 8/9 8/8 8/7 8/6 smallest largest 17 Periodic Trends from the Periodic Table 18 Periodic Trends from the Periodic Table

19 For transition metal atoms, outer most electrons (valence) are located in ns-orbital. Across the d-block series, e- of metal enter the (n-1)d-orbital or subshell. As such, these electron shield the outer most ns electron from the nucleus. In other words, the last electrons are located in the d orbitals which is in the n-1 level but the valence electrons are in the ns-level. Size of radius does not change drastically 19 Periodic Trends from the Periodic Table Lanthanide Contraction Between La and Hf Z = 7 to 7, and the number of protons increases resulting in increase in Z eff The increase in Zeff for Hf, offset the increase in n when going from Zr to Hf, therefore Zr and Hf are about the same size. 0 Periodic Trends from the Periodic Table 1 Periodic Trends from the Periodic Table

20 Periodic Trends from the Periodic Table

21 1 Periodic Trend: Ionization Energy and Electron Affinity Si Si + Si + Si + Si + Si + [Ne]s p [Ne]s p 1 [Ne]s [Ne]s 1 [He]s p ] [He]s p I I 1 I I Periodic Trend: Ionization Energy and Electron Affinity Ionization Energy versus atomic number Periodic Trend: Ionization Energy and Electron Affinity

22 Periodic Trend: Ionization Energy and Electron Affinity Consider the first Ionization potential for a series of elements. - Trend is that 1st IP. Determine by how strong valence electron is held by the nucleus. Periodic Trend: Ionization Energy and Electron Affinity 6 Periodic Trend: Ionization Energy and Electron Affinity

23 7 Periodic Trend: Ionization Energy and Electron Affinity 8 Periodic Trend: Ionization Energy and Electron Affinity Electron Affinity: The energy release when an electron is added to a gaseous atom to form an anion: A(g) + e A- (g) A(g) + e Energy release downhill A - (g) 9 Periodic Trend: Ionization Energy and Electron Affinity

24 Electron Affinity: Across PT Most main group elements have negative EA (exothermic) Favorable process Example: O (g) + e - O- (g) EA= -11 kj N (g) + e - N- (g) EA= 0 kj C (g) + e - C- (g) EA= -1 kj Why the anomaly? C=[He]s p [He] s p Z N=[He]s p [He] s p eff = Z O=[He]s p [He] s p eff = 6 Not a favorable process since there is now considerable e - e repulsion. Still some e-e- repulsion, but the Zeff increase in going from N to O over-comes this factor. 10 Periodic Trend: Ionization Energy and Electron Affinity Electron Affinity Down the PT: The electron affinity should become less favorable since the e- are being added to an orbital that is further away from the nucleus. Exception with Halogen 7A family: The electron is going into a smaller orbital in F relative to the orbitals found in Cl, Br and I. The e - e repulsion makes E.A an unfavorable process. That is there is less space in the nd shell in the F element which leads to substantial e -e repulsion. Br I 11 Periodic Trend: Ionization Energy and Electron Affinity 1 Periodic Trend: Ionization Energy and Electron Affinity

25 1 Periodic Trend: Ionization Energy and Electron Affinity 1 Periodic Trend: Ionization Energy and Electron Affinity IV. Reactivity: Note: Chemistry deals with the valence electrons of the atoms and how these electrons interact with each other to form bonds. Therefore, Elements with Low Ionization energy Reactive Elements with favorable Electron affinity Reactive Elements with Noble gas configuration Un reactive Elements with pseudo noble gas e- config. Not as reactive 1 Periodic Trend: Ionization Energy and Electron Affinity

26

27 1 Chemistry of Nonmetals Metals are found in the South West of the Periodic Table Parallels: Atomic radius Nonmetals are found in the North East of the Periodic Table Parallels, Ionization Energy, Electron affinity and Electronegativity Metalloids borders the metals and nonmetals Chemistry of Nonmetals Families of elements in the periodic table have similar chemical behavior. First member of each family, however, have some dramatic differences compared to the rest of the elements in that family. B C N O F metalloid nm nm nm nm metal character 6 7 # val e- () () () (1) # bonds Chemistry of Nonmetals

28 Nonmetals: Do not have luster: Various color Solids are brittle, few are hard and others soft Poor conductor of electricity Form molecular compounds with each other Form salts when combined with metals Form (acidic) oxides when combine with oxygen Form anions or oxyanions in solution. Group VIA elements. From left to right: Oxygen, sulfur, selenium and tellurium. Chemistry of Nonmetals Metalloid: Properties intermediate of nonmetals and metals Semiconductors Most common: Silicon Group IV elements. From left to right: Carbon (as graphite), silicon, germanium, tin and lead Chemistry of Nonmetals Hydrogen 1 val e- isotopes 1 H-protium H Deuterium H Tritium No particular family for H, although in the PT it is placed in the alkali (but it is a nonmetal) because of one valence e-. Sometimes it is also placed above halogens because it needs one more e- to get noble gas electron configuration. H H + + e- (common oxidation state) H + e- H- (when combined with electropositive element i.e., NaH) H-H bond is fairly strong but it can be broken (reacts) when activated by heat or catalyst H-O bonds very common and strong i.e., H O, H O H reacts well with carbon to form hydrocarbons (Organic) 6 Combines with alkali metal to form hydrides, i.e., NaH, CaH. Chemistry of Nonmetals

29 Nitrogen Atmospheric N 78% valence e- N N is very stable with a triple bond VA-group is the most paramagnetic family in representative. Chemistry Reaction with H to form ammonia (Haber Process) H + NH NH Ammonia raw ingredients for NH +, NO, NO, HNO, NO -, NO - oxidation state - to Important oxides N O (nitrous oxide) NO (nitric oxide) NO (nitrogen dioxide) Phosphorus: Most important among nitrogen family Allotropes White-P and Red-P 7 Chemistry of Nonmetals Oxygen 1% in Atmosphere 6 Valence allotropes; O (oxygen) O (ozone) Reaction O O ΔH = 8.6 (endothermic) Rxn w/ O M form metal oxides (O has Ox no. = - ) Oxide compound, Rust or protective layer. Peroxide are O - (O has Ox no. = -1) Super oxides O-O - (O has Ox no.= -1/) Sulfur: S 8 Less reactive than oxygen (lower electron affinity) Important contamination in petroleum and coal Found mostly in coal, when burned forms SO SO, SO - Air pollutant contributes to acid rain 8 Chemistry of Nonmetals Properties Diatomic (1- bonded specie) 7 val e- Highest EN among families of elements F has oxidation state of 0 or -1. All halogens (except F ) have variable charge (-1 to +7) 9 Chemistry of Nonmetals

30 Chemistry Properties : Diatomic (- bonded specie) 7 val e- Reacts with hydrogen to form strong acids, i.e., HCl. Form Oxyacids O m X(OH) n, i.e., HClO, HClO 10 Chemistry of Nonmetals 11 Chemistry of Nonmetals

31 1 Metals and Metallurgy Metals and Metallurgy Metals and Metallurgy

32 Properties and Metallic Bonding Physical Properties of Metals Important physical properties of pure metals: Malleable, ductile, good conductors, and feel cold. Shiny silvery Luster, various colors. Most metals are solids with the atoms in a close packed arrangement. There are not enough electrons for the metal atoms to be covalently bonded to each other. Tend to from cations in aq medium. Metal oxides are basic Metals and Metallurgy Alkali - ashes One valence electron Lowest IP 1 among families of elements. Most active among elements Chemistry Combines with H to form hydrides MH Combines with H O to form base, NaOH Combines with oxygen to from oxides, Li O Metals and Metallurgy Alkaline Earth Two valence electron Next lowest IP 1 among families of elements. Reactivity increase with atomic mass Chemistry Reacts w/ H O to form base, Ca(OH) or MgO Combines with oxygen to from oxides, MgO Combines with halogens to from salts, MgCl 6 Metals and Metallurgy

33 Theories describing properties & behavior. VBT (electron - sea model) Electron-sea model proposes: -Delocalized model for electrons in a metal. MO (Bond Theory) Delocalized model by Overlap of atomic orbitals 7 Metals and Metallurgy Valence orbitals overlap such that e - are free to migrate. The metal nuclei exist in a sea of electrons. No e - localized between any two metal atoms. Therefore, the electrons can flow freely. Without any definite bonds, metals easy to deform (and are malleable and ductile). High e- conductivity an heat capacity (heat conductor) ElectronSea model: metal bonds parallel number of valence electron, the greater the valence electrons, the greater the bond strength. 8 Metals and Metallurgy 9 Metals and Metallurgy

34 MO is a more qualitative approach. This theory treats bonding in terms of bonding and antibonding. Classic MO Theory. Consider Chromium: [Cr] = s 1 d Bond Order = 6 Electron fills bonding orbital, highest bond order is when all bonding orbital are occupied and there are no antibonding contributions. 10 Metals and Metallurgy 11 Metals and Metallurgy Next available orbital not available for e-, therefore e- are immobilized. Large gap for e - therefore e - can t be promoted 1 Metals and Metallurgy

35 Transition metal characterized by incomplete d-orbitals d-orbitals lead to multiple oxidation states. Z eff slowly increase resulting in lower oxidation states for later transition elements. (Since Zeff larger for latter transition metals, it is more difficult to remove e- therefore latter elements can t attain high Ox. number.) Sc (+)...Cr (+...+6)...Mn (+...+7)...Zn (+) Note Zn has the smallest ionic radii, and largest Z eff 1 Metals and Metallurgy Similar trend going down Period (nd and rd row transition metal) Note: Going down pt, expect Atomic radii to decrease: Sc<Y<La For Ti, Zr, Hf, this pattern doesn t follow trend Ti < Zr = Hf. Why? The size systematically decreases going across the periodic table until the 8B family, and then the atomic radii increases again. 1 Metals and Metallurgy Between La and Hf, Z = 7 to 7, and the number of protons increases resulting in increase in Z eff The increase in Z eff for Hf, offset the increase in n when going from Zr to Hf, therefore Zr and Hf are about the same size. 1 Metals and Metallurgy

36 Some miscellaneous properties for the fourth period transition elements. 16 Metals and Metallurgy Presence of unpaired of e- lead to interesting magnetic properties diamagnetic, no unpaired e-, actually is repelled by magnetic field paramagnetic- presence of unpaired e-. will be influence by strong magnetic. field. Attraction ferromagnetic - also presence of unpaired e- except that in there are domains in solid that permanently align selves in presence of magnetic field. Ferromagnetic more magnetic than paramagnetic material 17 Metals and Metallurgy The properties of elements in terms of their metallic character. 18 Metals and Metallurgy

37 Lecture Guide Please try and keep answers to single page Periodic Table and Periodic Trend Name I. Periodic Table 1. Who is given credited for conceiving the Periodic Table?. Define electron affinity. Why do these energy have negative potential.. The periodic table is also referred to as a map of.... The f-block are also referred to the and metals. 6. Why is the electron affinity of Cl more favorable than F even though F has a higher ionization energy?. How are size of atoms measured? 7. Where are the Rare Earth Elements found?. What two phenomena explains the size of an atom? 6. What is meant by Z eff? 8. How does a chemical reactivity relate to ionization energy and electron affinity? 7. Explain why oxygen is a smaller atom than nitrogen. 8. Calculate the Z eff for Se and Sb. III Chemistry of nonmetals. 1. Nonmetals are best characterize by their ability to do what? 9. In which portion of the periodic table are the largest atom? 10. Place the following in order of decreasing size: O, O -, O +, O - K+, K-, K 11. What are isoelectronic specie?. What are some of the properties of nonmetals?. What nonmetal is place in the same family as other metals? 1. What is the difference between a period and a family? 1. Explain the trend of atomic size across the transition metals. Why does the size of the metal remain relatively constant?. What is the calcogen family?. Which halogen is the most reactive? Why? 1. What is the Lanthanide contraction, why are transition metals in the third series nearly identical in size to elements in the second series? 6. Name all the oxy anions for iodine 7. What element(s) have reacted with noble gases? II. Ionization Energy and Electron Affinity 1. Define ionization energy.. For elements in the Group IIIA family, which ionization energy represents the disruption of a noble gas electron configuration?. Which elements in the periodic table have the highest ionization energy?. What is the trend of the ionization energy across the transition metals? IV. Metals 1. Nonmetals are best characterize by their ability to do what?. What are some physical properties of metal?. The sea of electron model is also known as what theory?. What bond theory is used to explain how insulator works?. What is meant by ferromagnetism?

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