AP Chemistry Periodicity.

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Presentation transcript:

AP Chemistry Periodicity

Brief Review of the Periodic Table metals: left side of Table; form cations properties: ductile (can pull into wire) malleable (can hammer into shape) lustrous (shiny) good conductors (heat and electricity)

Brief Review of the Periodic Table (cont.) nonmetals: right side of Table; form anions properties: good insulators gases or brittle solids neon sulfur iodine bromine Ne S8 I2 Br2

nonmetals Brief Review of the Periodic Table (cont.) metalloids (semimetals): “stair” between metals and nonmetals (B, Si, Ge, As, Sb, Te, Po, At) metals nonmetals computer chips Si and Ge properties: in-between those of metals and nonmetals; “semiconductors” computer chips Ge and Si 

alkaline earth metals: alkali metals: group 1 (except H); 1+ charge; very reactive alkaline earth metals: group 2; 2+ charge; less reactive than alkalis transition elements: groups 3–12; variable charges chalcogens: group 16; 2– charge; reactive halogens: group 17; 1– charge; very reactive noble gases: group 18; no charge; unreactive lanthanides: elements 58–71 contain f orbitals actinides: elements 90–103 main block (representative) elements: groups 1, 2, 13–18

What family of elements has an ns2 valence electron configuration? alkaline earth metals

Anomalies in the Electron Configurations Your best guide to writing e– configs is “The Table,” but there are a few exceptions. e.g., Cr: [ Ar ] 4s1 3d5 Cu: [ Ar ] 4s1 3d10 These exceptions are due to the closeness in energy of the upper-level orbitals. “RuRh…!!” Other exceptions are… Mo, Ru, Rh, and Ag. All of these exceptions have a single valence-level s electron.

Sodium and potassium react w/water to produce hydrogen gas. valence orbitals: outer-shell orbitals -- elements in the same group have the same valence-shell electron configuration -- since valence e– are involved in bonding, elements within a group have many of the same properties Sodium and potassium react w/water to produce hydrogen gas.

Development of the Periodic Table -- few elements appear in elemental form in nature (Au, Ag, Hg, a few others) -- most are in combined forms with other elements -- In 19th century, advances in chemistry allowed more elements to be identified. Au Ag Hg

Dmitri Mendeleev Lothar Meyer 1834–1907 1830–1895 1869: Independently, Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) published classification schemes based on similarities in element properties. Dmitri Mendeleev 1834–1907 ** Mendeleev used his scheme to predict the existence of undiscovered elements, and so is given credit for inventing the first periodic table. ** D.M.’s guiding principle was… “atomic mass.” Lothar Meyer 1830–1895

Henry Moseley Graph of Moseley’s data 1887–1915 -- 1913: Henry Moseley bombarded atoms with high-energy electrons and measured the frequency of the X rays given off. X ray frequency generally increased as atomic mass increased, but VERY nicely increased as ____________ increased. atomic number Henry Moseley 1887–1915 Graph of Moseley’s data

Gilbert Lewis Ar Ne He 1875–1946 e– density distance from nucleus Electron Shells Even before Bohr, the American Gilbert Lewis had suggested that e– are arranged in shells. -- Experiments show that e– density is a maximum at certain distances from nucleus. Gilbert Lewis Ar e– density distance from nucleus 1875–1946 Ne -- no clearly defined boundaries between shells He (shells are diffuse, i.e., “fuzzy”)

r d r Approximate bonding atomic radii for the elements have been tabulated. The distance between bonded nuclei can be approximated by adding radii from both atoms. e.g., Bonding atomic radii are as follows: C = 0.77 A, Br = 1.14 A So the approximate distance between bonded C and Br nuclei = 0.77 + 1.14 = 1.91 A

In a many-electron atom, each e– is attracted to the nucleus and repelled by the other e–. -- effective nuclear charge, Zeff: the net (+) charge attracting an e– (a measure of how tightly particular e–s are held) Z = atomic number Equation: 3d6 S = # of e– BETWEEN nucleus and e– in question (NOT e– in same subshell) 4s2 Zeff = Z – S 3p6 -- Within a given electron shell, s e–s have the greatest Zeff, f e–s the least. Fe 3s2 2p6 2s2 For Fe, the 3p e–s have Zeff = 26 – 12 = 14; the 3d e–s have Zeff = 26 – 18 = 8. 1s2 nuc.

v.e– v.e– Li K -- The (+) charge “felt” by the outer e– is always less than the nuclear charge. This effect, due to the core (or kernel) electrons, is called the... screening effect (or shielding effect). K v.e– Li v.e– tougher to remove easier to remove

Atomic Radius As we go down a group, atomic radius… increases. -- principal quantum number increases (i.e., a new energy level is added) As we go from left to right across the Table, atomic radius… decreases. -- effective nuclear charge increases, but principal quantum number is constant more p+, but no new (i.e., farther away) energy levels

2+ 2– 1+ 1– Coulombic attraction depends on… amount of charge distance between charges As we go , more coulombic attraction, no new energy level, more pull, smaller size + – + – + – H He + –

Arrange the following atoms in order of increasing atomic radius: Sr, Ba, Cs Sr < Ba < Cs

Ionization Energy: the minimum energy needed to remove an e– from an atom or ion M(g) + 1st I.E.  M+(g) + e– M+(g) + 2nd I.E.  M2+(g) + e– M2+(g) + 3rd I.E.  M3+(g) + e– e– I.E. Successive ionization energies are larger than previous ones. -- (+) attractive force remains the same, but there is less e–/e– repulsion

The ionization energy increases sharply when we try to remove an inner-shell electron. e.g., For Mg, 1st IE = 738 kJ/mol 2nd IE = 1,450 kJ/mol 3rd IE = 7,730 kJ/mol (strong evidence that only valence e– are involved in bonding) As we go down a group, 1st IE… decreases. -- more e–/e– repulsion and more shielding

(easier to remove B’s single 2p e– than one of Be’s two 2s e–s) Generally, as we go from left to right, 1st IE… Exceptions: e.g., B < Be B doesn’t like 2p Be: 1s2 2s2 B: 1s2 2s2 2p1 (easier to remove B’s single 2p e– than one of Be’s two 2s e–s) Subshells prefer to be either completely filled OR half-filled. …than any of these. N: 1s2 2s2 2p3 More stable to have O: 1s2 2s2 2p4 than to have 2p This e– is easier to remove…

down a group… across a period… across a period… First

energy energy Electron affinity: the energy change that occurs when an e– is added to a gaseous atom For most atoms, adding an e– causes energy to be… released. eq. for e– affinity: A + e– A– e– energy e– I.E. energy Exceptions: noble gases: the added e– must go into a new, higher energy level group 2 metals: the added e– must go into a higher-energy p orbital group 15 elements: the added e– is the first one to double-up a p orbital

The halogens have the most (–) electron affinities, meaning that they become very stable when they accept electrons. –328 F Cl Br I O S Se Te Ne Ar Kr Xe He –349 –325 –295 –141 –200 –195 –190 + more (–) e– affinity more willing to accept an e– = Electron affinities don’t vary much going down a group.

Regions of the Table metals: left side of Table; form cations properties: ductile (can pull into wire) malleable (can hammer into shape) lustrous (shiny) good conductors (heat and electricity)

-- Because of their low ionization energies, they are often oxidized in reactions. (i.e., they lose e–) -- Metallic character of the elements increases as we go down-and-to-the-left. increasing metallic character

Regions of the Table (cont.) nonmetals: right side of Table; form anions properties: good insulators; gases or brittle solids neon sulfur iodine bromine Ne S8 I2 Br2 -- memorize the HOBrFINCl

nonmetals Regions of the Table (cont.) metalloids (semimetals): “stair” between metals and nonmetals (B, Si, Ge, As, Sb, Te, Po, At) metals nonmetals computer chips Si and Ge properties: in-between those of metals and nonmetals; “semiconductors” Si and Ge computer chips

(i.e., a “basic” oxide) Reactivity Trends metal oxide + water metal hydroxide MgO(s) + H2O(l) Mg(OH)2(aq) metal oxide + acid salt + water CaO(s) + 2 HNO3(aq) Ca(NO3)2(aq) + H2O(l) metal + nonmetal salt 2 Al(s) + 3 Br2(l) 2 AlBr3(s)

(i.e., an “acidic” oxide) Reactivity Trends (cont.) nonmetal oxide + water acid CO2(g) + H2O(l) H2CO3(aq) nonmetal oxide + base salt + water CO2(g) + 2 KOH(aq) K2CO3(aq) + H2O(l)

Group Trends Alkali Metals -- the most reactive metals (one e– to lose) -- obtained by electrolysis of a molten salt e.g., chloride ion is oxidized and sodium ion is reduced 2 NaCl(l) 2 Na(l) + Cl2(g)

Potassium in water, forming flammable hydrogen -- react with hydrogen to form metal hydrides: 2 M(s) + H2(g) 2 MH(s) -- react with water to form metal hydroxides: 2 M(s) + 2 H2O(l) 2 MOH(aq) + H2(g) -- react w/O2: Li yields Li2O, others yield (mostly) peroxides (M2O2) 2 M(s) + O2(g) M2O2(s) Potassium in water, forming flammable hydrogen and soluble potassium hydroxide.

Alkaline-Earths -- not as reactive as alkalis (two e– to lose) compared to alkalis: harder, denser, higher MPs -- Ca and heavier ones react w/H2O to form metal hydroxides Ca(s) + 2 H2O(l) Ca(OH)2(aq) + H2(g) -- MgO is a protective oxide coating around substrate Mg Mg ribbon MgO

The Hindenburg (She was scuttled in June 1919, along with 71 other Hydrogen -- a nonmetal, but belongs to no family -- reacts w/other nonmetals to form molecular (i.e., covalent) compounds The Hindenburg (She was scuttled in June 1919, along with 71 other German ships.) (She burned up in May 1937, killing 36 passengers.)

noble gas to fill halogen lamps. The halogen sets -- At isn’t considered to be a halogen; little is known about it -- at 25oC, F2 and Cl2 are gases, Br2 is a liquid, I2 is a solid -- their exo. reactivity is dominated by their tendency to gain e– A small amount of a halogen is mixed with a noble gas to fill halogen lamps. The halogen sets up an equilibrium with the tungsten filament to prevent the heated tungsten from being deposited on the inside of the bulb. -- Cl2 is added to water; the HOCl produced acts as a disinfectant -- HF(aq) = weak acid; HCl(aq) HBr(aq) HI(aq) = strong acids

professional-grade Rn detector Fan for Rn mitigation Noble Gases -- all are monatomic; have completely-filled s and p orbitals -- He, Ne, and Ar have no known compounds; Rn is radioactive -- Kr has one known compoud (KrF2); Xe has a few (XeF2, XeF4, XeF6) professional-grade Rn detector Fan for Rn mitigation

Ionic Radius Cations are _______ than parent atoms; anions are ______ than parent atoms. smaller larger Ca atom Ca2+ ion Cl atom Cl– ion 20 p+ 20 p+ 17 p+ 17 p+ 20 e– 18 e– 17 e– 18 e– Cl– Ca Cl Ca2+ EX. Compare the sizes of Fe, Fe2+, and Fe3+. Then compare Br with Br–. Fe > Fe2+ > Fe3+ Br– > Br

Linus Pauling quantified the electronegativity scale. the tendency for a bonded atom to attract e– to itself Electronegativity increases going... up and to-the-right. electronegativity increases Most electronegative element is... fluorine (F).

“Oh, man… I forgot which ones the most electronegative elements are.” F = 4.0 O = 3.5 N = Cl = 3.0 Others: C = 2.5 H = 2.1