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
sulfur
iodine
bromine
Ne
S8
I2
Br2
Brief Review of the Periodic Table (cont.)
metalloids (semimetals): “stair” between metals
(B, Si, Ge, As, Sb, Te, Po, At)
and nonmetals
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:
halogens:
noble gases:
lanthanides:
actinides:
group 2; 2+ charge;
less reactive than alkalis
groups 3–12; variable charges
group 16; 2– charge; reactive
group 17; 1– charge; very reactive
group 18; no charge; unreactive
elements 58–71
contain f orbitals
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.
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 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
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
-- 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.
(shells are diffuse, i.e., “fuzzy”)
e– density
-- no clearly defined
boundaries between shells
Ar
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, Zeff: the net (+) charge
attracting an e–
(a measure of how tightly particular e–s are held)
Equation:
Zeff = Z – S
Z = atomic number
3d6
S = # of e– BETWEEN nucleus
and e– in question (NOT e–
in same subshell)
-- Within a given electron shell,
s e–s have the greatest Zeff,
f e–s the least.
4s2
3p6
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.
-- 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–
Li
v.e–
K
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
Coulombic attraction depends on…
amount of charge
distance between charges
2+
2–
1+
1–
2+
2+
+
–
H
+ +
He
+
–
+
–
–
–
2–
2–
As we go , more coulombic
attraction, no new energy level,
more pull, smaller size
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) + 2nd I.E. 
M2+(g) + 3rd I.E. 
M+(g) + e–
M2+(g) + e–
M3+(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, 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…
--
more e–/e– repulsion and more shielding
decreases.
Generally, as we go from left to right, 1st IE…
Exceptions: e.g., B < Be
Be:
2p
1s2 2s2
B doesn’t like
B:
1s2 2s2 2p1
Subshells prefer to be
either completely filled
OR half-filled.
N:
1s2 2s2 2p3
O:
1s2 2s2 2p4
(easier to remove B’s
single 2p e– than one
of Be’s two 2s e–s)
…than any
of these.
More stable to have
than to have
2p
This e–
is easier to remove…
First
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–
released.
A–
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.
He
+
O
more (–)
e– affinity
=
more willing to
accept an e–
–141
S
–200
Electron affinities don’t vary
much going down a group.
Se
–195
Te
–190
F
Ne
–328
+
Cl
Ar
–349
+
Br
Kr
–325
+
I
Xe
–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
sulfur
Ne
S8
-- memorize the HOBrFINCl
iodine
I2
bromine
Br2
Regions of the Table (cont.)
metalloids (semimetals): “stair” between metals
(B, Si, Ge, As, Sb, Te, Po, At)
and nonmetals
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
metal hydroxide
MgO(s) + H2O(l)
metal oxide + acid
CaO(s) + 2 HNO3(aq)
metal + nonmetal
2 Al(s) + 3 Br2(l)
Mg(OH)2(aq)
salt + water
Ca(NO3)2(aq) + H2O(l)
salt
2 AlBr3(s)
(i.e., an
“acidic”
oxide)
Reactivity Trends (cont.)
nonmetal oxide + water
acid
CO2(g) + H2O(l)
nonmetal oxide + base
CO2(g) + 2 KOH(aq)
H2CO3(aq)
salt + water
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)
-- 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
Hydrogen
-- a nonmetal, but belongs to no family
-- reacts w/other nonmetals to form
molecular (i.e., covalent) compounds
The Hindenburg
(She
burned
up in
(She
was
scuttled
in May
June
1919,1937,
along
killing
with
71 other
36 passengers.)
German
ships.)
Halogens
-- 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–
-- Cl2 is added to water; the HOCl
produced acts as a disinfectant
-- HF(aq) = weak acid;
HCl(aq)
HBr(aq)
= strong acids
HI(aq)
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 (KrF2);
Xe has a few (XeF2, XeF4, XeF6)
professional-grade
Rn detector
Fan for
Rn mitigation
Ionic Radius
smaller
Cations are _______ than parent atoms;
anions are ______ thanlarger
parent atoms.
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
Cl–
Ca
Ca2+
EX. Compare the sizes of Fe,
Fe2+, and Fe3+.
Then compare Br with Br–.
Fe > Fe2+ > Fe3+
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