1. AP
Chemistry
Periodicity

2. 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)

3. 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

4. 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 

5. 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

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

7. 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.

8. 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.

9. 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

10. 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

11. Henry Moseley Graph of Moseley’s data 1887–1915
: 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

12. 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”)

13. 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 =
= A

14. 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.

15. 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

16. 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

17. 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 + –

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

19. 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

20. 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

21. (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…

22. down a group…
across a period…
across a period…
First

23. 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

24. 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.

25. 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)

26. -- 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

27. 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

28. 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

29. (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)

30. (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)

31. 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)

32. Potassium in water, forming flammable hydrogen
-- react with hydrogen to form metal hydrides:
2 M(s) + H2(g) MH(s)
-- react with water to form metal hydroxides:
2 M(s) H2O(l) 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.

33. 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) H2O(l) Ca(OH)2(aq) + H2(g)
-- MgO is a
protective
oxide coating
around
substrate Mg
Mg ribbon
MgO

34. 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.)

35. 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

36. 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

37. 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

38. 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).

39. “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