Chapter 6
“The Periodic Table”
The Elements by Tom Lehrer
Organizing the Elements
 used
properties of elements to sort
into groups.
 1829 J. W. Dobereiner arranged
elements into triads – groups of 3 w/
similar properties
 One
element in each triad
had properties intermediate
of the other two elements
 Cl, Br, and I look different,
but similar chemically
Mendeleev’s Periodic Table
 mid-1800s,
about 70 elements
known
 Dmitri Mendeleev – Russian chemist
& teacher
 Arranged elements by
increasing atomic mass
Mendeleev
 blanks for undiscovered
elements
When
discovered, his predictions
accurate
 Problems
Co to Ni
Ar to K
Te to I
w/ order
A better arrangement
 1913, Henry Moseley – British
physicist, arranged elements
according to increasing atomic
number
The Elements by Tom Lehrer
Periodic Law
 When
elements arranged in order of
increasing atomic #, periodic
repetition of phys & chem props
 Horizontal rows = periods
7
periods
 Vertical
column = group (or family)
Similar
phys & chem prop.
ID’ed by # & letter (IA, IIA)
Areas of periodic table

3 classes of elements:
1) Metals: electrical conductors, have
luster, ductile, malleable
2) Nonmetals: generally brittle and
non-lustrous, poor conductors of heat
and electricity





Some gases (O, N, Cl)
some brittle solids (B, S)
fuming red liquid (Br)

3) Metalloids: border the line-2 sides

Properties are intermediate between
metals and nonmetals
Section 6.2
Classifying the Elements
 OBJECTIVES:
Describe
the information
in a periodic table.
Section 6.2
Classifying the Elements
 OBJECTIVES:
Classify
elements based
on electron
configuration.
Section 6.2
Classifying the Elements
 OBJECTIVES:
Distinguish
representative elements
and transition metals.
Groups of elements - family names
 Group
IA – alkali metals
Forms
“base” (or alkali) when
reacting w/ H2O (not just dissolved!)
 Group
2A – alkaline earth metals
form bases with H2O; don’t
dissolve well, hence “earth metals”
Also
 Group
7A – halogens
“salt-forming”
Electron Configurations in Groups

Elements sorted based on econfigurations:
1) Noble gases
2) Representative elements
3) Transition metals
4) Inner transition metals
Let’s
now
take a
closer
look at
these.
Electron Configurations in Groups
1) Noble gases in Group 8A
(also
called Group 18)
 very stable = don’t react
 e- configuration w/ outer s & p
sublevels full
Electron Configurations in Groups
2) Representative Elements Groups
1A - 7A
 wide range of properties
 “Representative” of all elements
 s & p sublevels of highest energy level
NOT filled
 Group # equals # of e- in highest
energy level
Electron Configurations in Groups
3) Transition metals in “B” columns
 outer s sublevel full
 Start filling “d” sublevel
 “Transition” btwn metals &
nonmetals
Electron Configurations in Groups
4) Inner Transition Metals below
main body of PT, in 2 horizontal
rows
 outer s sublevel full
 Start filling “f” sublevel
 Once called “rare-earth” elements
 not true b/c some abundant

1A
2A
Elements 1A-7A groups called
8A
representative elements
3A 4A 5A 6A 7A
outer s or p filling
The group B called transition
elements
 These
are called the inner
transition elements, and they
belong here
Group 1A called alkali metals (but NOT H)
Group 2A called alkaline earth metals
H
Group 8A are noble gases
 Group 7A called halogens


Let’s take a quick break……
Periodic table rap
H
Li
1s1
1
1s22s1
Do you notice any similarity in these
configurations of the alkali metals?
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10
5p66s1
1s22s22p63s23p64s23d104p65s24d105p66
s24f145d106p67s1
Do you notice any similarity in the
configurations of the noble gases?
1s2 He 2
Ne
2
2
6
1s 2s 2p
10
1s22s22p63s23p6 Ar18
1s22s22p63s23p64s23d104p6 Kr
36
1s22s22p63s23p64s23d104p65s24d105p6
Xe
54
1s22s22p63s23p64s23d104p65s24d10 Rn
5p66s24f145d106p6 86
Elements
in
the
s
blocks
1
s
s2
He
Alkali metals end in s1
 Alkaline earth metals end in s2

 should
include He, but…
He has properties of noble gases
 has a full outer level of e-’s


group 8A.
Transition Metals - d block
Note the change in configuration.
1
d
2
d
3
d
s1
5
d
s1
5
6
7
8
10
10
d d d d d d
The P-block p1
p2
p3
p4
p5
p6
F - block

Called “inner transition elements”
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1
2
3
Period
Number
4
5
6
7

Each row (or period) is energy level for
s & p orbitals.

“d” orbitals fill up in levels 1 less than
period #
 first
d is 3d found in period 4.
1
2
3
4
4d
5d
5
6
7
3d
1
2
3
4
5
6
7
4f
5f
f
orbitals start filling at 4f….2 less
than period #
Demo p. 165
Section 6.3
Periodic Trends
 OBJECTIVES:
Describe
trends among the
elements for atomic size.
Section 6.3
Periodic Trends
 OBJECTIVES:
Explain
how ions form.
Section 6.3
Periodic Trends
 OBJECTIVES:
Describe
periodic trends
for first ionization energy,
ionic size, and
electronegativity.
Trends in Atomic Size
}
Radius
Measure
Atomic Radius - half distance
btwn 2 nuclei of diatomic molecule (i.e. O2)
Units
of picometers (10-12 m… 1 trillionth)
ALL Periodic Table Trends
 Influenced
by 3 factors:
1. Energy Level
Higher energy levels further away
from nucleus.
2. Charge on nucleus (# protons)
More charge pulls electrons in
closer. (+ and – attract each other)
 3. Shielding effect
What do they influence?
Energy levels & Shielding
have effect on GROUP (  )
Nuclear
charge has effect on
PERIOD (  )
#1. Atomic Size - Group trends
Going down a
group, each atom
has another
energy level (floor)
 atoms get

r
H
Li
Na
K
bigge
Rb
#1. Atomic Size - Period Trends
 left
to right across period:
size
gets smaller
e-’s occupy same energy level
more nuclear charge
Outer e-’s pulled closer
Na
Mg
Al
Si
P
Here is an
animation to
explain the
trend.
S Cl Ar
Rb
K
Atomic Radius (pm)
Period 2
Na
Li
Kr
Ar
Ne
H
3
10
Atomic Number
Trends of Atomic Radius
Ions
 Some
compounds composed of
“ions”
ion is atom (or group of atoms) w/ + or charge

Atoms are neutral because the number of
protons = electrons
+
& - ions formed when e- transferred (lost or
gained) btwn atoms
Ions
 Metals LOSE electrons, from outer
energy level
Sodium loses 1 e more
p+ (11) than e- (10)
 + charge particle formed…“cation”
Na+
called “sodium ion”
Ions
 Nonmetals
GAIN one or more
electrons
Cl gains 1 ep+ (17) & e- (18), so charge of -1
called “chloride ion”
anions
 Cl1
#2. Trends in Ionization Energy
 Ionization
energy - energy required to
completely remove e- (from gaseous
atom)
 energy required to remove only 1st ecalled first ionization energy.
Ionization Energy
 second
ionization energy is E
required to remove 2nd eAlways greater than first IE.
 third greater than 1st or 2nd IE.
Table 6.1, p. 173
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
5247
7297
1757
2430
2352
2857
3391
3375
3963
Third
Why did these values
increase so much?
11810
14840
3569
4619
4577
5301
6045
6276
What factors determine IE
 greater
nuclear charge = greater IE
 Greater distance from nucleus
decreases IE
 Filled & half-filled orbitals have lower
energy
 Easier
to achieve (lower IE)
 Shielding
effect
Shielding
in outer
energy level
“looks through” all
other energy
levels to see
nucleus
 e-’s
Ionization Energy - Group trends
 going
first
down group
IE decreases b/c...
e- further away from
nucleus attraction
more shielding
Ionization Energy - Period trends
 Atoms
in same period:
same
energy level
Same shielding
 Increasing nuclear charge
 So IE generally increases left - right
 Exceptions…full
& 1/2 full orbitals
First Ionization energy
He
H
 He
has greater IE
than H.
 Both have same
shielding (e- in 1st
level)
 He
= greater nuclear
charge
Atomic number
First Ionization energy
He
 Li
H
Li
lower IE than H
 more shielding
 further away
 These outweigh
greater nuclear
charge
Atomic number
First Ionization energy
He
 Be
H
Be
higher IE than
Li
 same shielding
 greater nuclear
charge
Li
Atomic number
First Ionization energy
He
B
H
Be
B
Li
has lower IE
than Be
 same shielding
 greater nuclear
charge
 By removing an
electron we make
s orbital half-filled
Atomic number
First Ionization energy
He
H
Be
C
B
Li
Atomic number
First Ionization energy
He
N
H
C
Be
B
Li
Atomic number
First Ionization energy
He
 Oxygen
N
H
C O
Be
B
Li
breaks
the pattern,
because
removing an
electron leaves it
with a 1/2 filled p
orbital
Atomic number
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
First Ionization energy
He
Ne
N F
H
C O
Be
B
Li
 Ne
has a lower
IE than He
 Both are full,
 Ne has more
shielding
 Greater
distance
Atomic number
Ne
First Ionization energy
He
N F
H
C O
Be
B
Li
 Na
has a lower
IE than Li
 Both are s1
 Na has more
shielding
 Greater
distance
Na
Atomic number
Atomic number
First Ionization energy
Trends in Ionization Energy (IE)
Driving Forces
 Full
Energy Levels require high E
to remove eNoble Gases = full orbitals
 Atoms want noble gas
configuration
2nd Ionization Energy
 For
elements w/ filled or ½ filled
orbital by removing 2 e-, 2nd IE
lower than expected.
 True for s2
 Alkaline earth metals form 2+
ions.
3rd IE
the same logic s2p1
atoms have an low 3rd IE.
 Atoms in the aluminum family
form 3+ ions.
 2nd IE and 3rd IE are always
higher than 1st IE!!!
 Using
Trends in Ionic Size: Cations
 Cations
form by losing electrons.
 metals
 Cations
are smaller than the atom they
came from –
 they
lose electrons
 they lose an entire energy level.
 Cations
of representative elements
have noble gas configuration before
them.
Trends in Ionic size: Anions
 Anions
gain electrons
 Anions
bigger than the atom they
came from –
same
energy level
greater area the nuclear charge needs to
cover
 Nonmetals
Configuration of Ions
Ions always have noble gas
configurations (full outer level)
 Na atom is: 1s22s22p63s1
 Forms a 1+ sodium ion: 1s22s22p6
 Same as Ne

Configuration of Ions
 Non-metals
form ions by
gaining electrons to achieve
noble gas configuration.
 They end up with the
configuration of the noble gas
after them.
Ion Group trends
 Each
step down a
group is adding an
energy level
 Ions get bigger
going down, b/c of
extra energy level
Li1+
Na1+
K1+
Rb1+
Cs1+
Ion Period Trends
 Across
period
 nuclear
charge increases
 Ions get smaller.
 energy
level changes between
anions and cations.
Li1+
B3+
Be2+
C4+
N3-
O2-
F1-
Size of Isoelectronic ions
means “the same”
 Isoelectronic ions have the same # of
electrons
 Al3+ Mg2+ Na1+ Ne F1- O2- and N3all have 10 electrons
 all have the same configuration:
1s22s22p6 (which is the noble gas: neon)
 Iso-
Size of Isoelectronic ions?

Positive ions that have more protons
would be smaller (more protons would
pull the same # of electrons in closer)
Al3+
13
12
Na1+
11
Mg2+
Ne
F1-
10
9
2O
8
N37
#3. Trends in Electronegativity
Electronegativity is tendency for
atom to attract e-’s when atom in a
compound
 Sharing e-, but how equally do they
share it?
 Element with big electronegativity
means it pulls e- towards itself
strongly!

Electronegativity Group Trend
 Further down a group,
farther e- is away from
nucleus, plus the more e-’s
an atom has
 more willing to share
 Low electronegativity
Electronegativity Period Trend
 Metals
low
let e-’s go easily
electronegativity
 Nonmetals
take
want more electrons
them away from others
High electronegativity.
Trends in Electronegativity
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Chemistry Song "Elemental Funkiness" Mark Rosengarten