Slide 1

Transcript Slide 1

AP Chemistry Periodicity

Brief Review of the Periodic Table

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

Brief Review of the Periodic Table (cont.)

nonmetals: right side of Table; form anions properties: good insulators gases or brittle solids neon Ne sulfur S 8 iodine I 2 bromine Br 2

Brief Review of the Periodic Table (cont.)

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

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

What family of elements has an

ns 2

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 ] 4s 1 3d 5 Cu: [ Ar ] 4s 1 3d 10 These exceptions are due to the closeness in energy of the upper-level orbitals. Other exceptions are… Mo, Ru, Rh, and Ag.

All of these exceptions have a single valence-level s electron. “RuRh…!!”

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

/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

1869: Independently, Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) published classification schemes based on similarities in element properties.

** 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.” Dmitri Mendeleev

1834–1907

Lothar Meyer

1830–1895

-- 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

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. -- no clearly defined boundaries between shells (shells are diffuse, i.e., “fuzzy”)

Gilbert Lewis

1875–1946

Ne He

distance from nucleus

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, Z eff : the net (+) charge attracting an e – (a measure of how tightly particular e – s are held) Z Equation: eff = Z – S Z = atomic number S = # of e – and e – BETWEEN nucleus in question (NOT e – in same subshell) -- Within a given electron shell, s e – s have the greatest Z eff , f e – s the least. For Fe, the 3p e – s have Z eff = 26 the 3d e – s have Z eff – 12 = 14; = 26 – 18 = 8.

Fe 3d 6 4s 2 3p 6 3s 2 2p 6 2s 2 1s 2 nuc.

- 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).

v.e

– tougher to remove

v.e

– 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

Coulombic attraction depends on… amount of charge 2+ 2– distance between charges 2+ 2– 1+ 1– 2+

+ –

2– As we go , more coulombic attraction, no new energy level, more pull, smaller size H

+ + – –

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) + 1 st I.E.  M + (g) + 2 nd I.E.  M 2+ (g) + 3 rd I.E.  M M M + 2+ 3+ (g) + e (g) + e (g) + 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, 1 st 2 nd IE = 738 kJ/mol IE = 1,450 kJ/mol 3 rd IE = 7,730 kJ/mol (strong evidence that only valence e – are involved in bonding) As we go down a group, 1 st IE… - more e – /e – repulsion and more shielding decreases.

Generally, as we go from left to right, 1 st IE… Be: B: Exceptions: e.g., B < Be 1s 2 2s 2 B doesn’t like 1s 2 2s 2 2p 1

Subshells prefer to be either completely filled OR half-filled.

2p (easier to remove B’s single 2p e – than one of Be’s two 2s e – s)

…than any of these.

N: O: 1s 2 2s 2 2p 3 1s 2 2s 2 2p 4 More stable to have than to have

This e

– 2p

is easier to remove…

down a group…

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… eq. for e – affinity: A + e – A – released.

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. more (–) e – affinity = more willing to accept an e – Electron affinities don’t vary much going down a group.

–141

–200

–195

–190

–328

–349

–325

–295

Regions of the Table

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

-- 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.

Regions of the Table (cont.)

nonmetals: right side of Table; form anions properties: good insulators; gases or brittle solids neon Ne sulfur S 8 -- memorize the HOBrFINCl iodine I 2 bromine Br 2

Regions of the Table (cont.)

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

(i.e., a “basic” oxide)

Reactivity Trends

metal oxide + water MgO(s) + H 2 O(l) metal hydroxide Mg(OH) 2 (aq) metal oxide + acid CaO(s) + 2 HNO 3 (aq) metal + nonmetal 2 Al(s) + 3 Br 2 (l) salt + water Ca(NO 3 ) 2 (aq) + H 2 O(l) salt 2 AlBr 3 (s)

(i.e., an “acidic” oxide)

Reactivity Trends (cont.)

nonmetal oxide + water CO 2 (g) + H 2 O(l) acid H 2 CO 3 (aq) nonmetal oxide + base CO 2 (g) + 2 KOH(aq) salt + water K 2 CO 3 (aq) + H 2 O(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) + Cl 2 (g)

-- react with hydrogen to form metal hydrides: 2 M(s) + H 2 (g) 2 MH(s) -- react with water to form metal hydroxides: 2 M(s) + 2 H 2 O(l) 2 MOH(aq) + H 2 (g) -- react w /O 2 : Li yields Li 2 O, others yield (mostly) peroxides (M 2 O 2 ) 2 M(s) + O 2 (g) M 2 O 2 (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 /H 2 O to form metal hydroxides Ca(s) + 2 H 2 O(l) Ca(OH) 2 (aq) + H 2 (g) -- MgO is a protective oxide coating around substrate Mg Mg ribbon MgO

Hydrogen

-- a nonmetal, but belongs to no family -- reacts w /other nonmetals to form molecular (i.e., covalent) compounds The Hindenburg

Halogens

-- At isn’t considered to be a halogen; little is known about it -- at 25 o C, F 2 Br 2 and Cl is a liquid, I 2 2 are gases, is a solid -- their exo. reactivity is dominated by their tendency to gain e – -- Cl 2 is added to water; the HOCl produced acts as a disinfectant -- HF(aq) = weak acid; HCl(aq) HBr(aq) HI(aq) = strong acids

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.

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 (KrF 2 ); Xe has a few (XeF 2 , XeF 4 , XeF 6 )

professional-grade Rn detector

Fan for Rn mitigation

Ca atom 20 p + 20 e –

Ionic Radius

smaller Cations are _______ than parent atoms; larger Ca 2+ ion 20 p + 18 e – Cl atom 17 p + 17 e –

Ca Ca 2+ Cl

Cl – ion 17 p + 18 e –

Cl –

EX. Compare the sizes of Fe, Fe 2+ , and Fe 3+ .

Then compare Br with Br – . Fe > Fe 2+ > Fe 3+ Br – > Br

Electronegativity

electronegativity: the tendency for a bonded atom to attract e – to itself Electronegativity increases going... up and to-the-right.

Linus Pauling quantified the electronegativity scale.

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