Transcript Chapter 6 The Periodic Table - (Home) Collinsville Public

Chapter 6
The Periodic Table
p. 154
The Elements by Tom Lehrer
Organizing the Elements
 Chemists
used elements properties
to sort into groups.
 1829 - J. W. Dobereiner
 triads

– groups of 3 w/ similar properties
One element in triad
had properties intermediate
of other 2 elements
 Cl, Br, and I look different….
similar chemically
Mendeleev’s Periodic Table
 1800s,
about 70 elements known
 1869 - Dmitri Mendeleev – Russian
chemist & teacher
 Arranged elements by atomic mass
Mendeleev’s Periodic Table
 Blank
spaces
undiscovered
elements
 Predicted
properties
predictions very accurate
 Problems
Te
w/ order
to I atomic mass decreases
I belongs w/ Br & Cl
Mendeleev broke rule – put Te before I
A better arrangement
 1913, Henry Moseley
British
physicist
Determined atomic #’s
Modern PT arranged
by atomic #
The Elements by Tom Lehrer
Periodic Law
 Elements
arranged by increasing
atomic #, periodic repetition of
properties present
 Horizontal rows = periods
7
periods
 Vertical
column = group (family)
Similar
properties
IUPAC labels (1-18)
 U.S.
system (# & letter…i.e. IA, IIA)
Areas of periodic table

3 classes of elements:

1) Metals: electrical conductors,
lustrous, ductile, malleable

2) Nonmetals: generally brittle/nonlustrous, 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
p. 161
Groups of elements - family names
 Group
IA – alkali metals
Forms
“base” (or alkali) when reacting
w/ H2O
 Group
(not just dissolved!)
2A – alkaline earth metals
form bases with H2O; don’t
dissolve well, hence “earth metals”
Also
 Group
7A – halogens
Greek
hals (salt) & genesis (to be born)
Electron Configurations in Groups

sorted based on e- configs:
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
called Group 18)
 very stable = don’t react
 e- configuration
 full outer s & p sublevels
(also
Electron Configurations in Groups
2) Representative Elements
Groups 1A - 7A
 Properties vary
 “Represent” all elements
 s & p sublevels of highest PEL NOT
filled
 Group # = valence e-’s
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
PT, 2 horizontal rows
 outer s sublevel full
 Start filling “f” sublevel

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
metals end in s1
 Alkaline earth metals end in s2
 Alkali
should
include He, but…
 properties
full
of noble gases
outer EL
 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 period # = 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
p. 170
Trends in
Atomic Size
Atomic
Radius - half
distance btwn 2 nuclei
of identical atoms
Increases top-bottom
Decreases L-R
picometers
 10-12
m… 1 trillionth
Radius
ALL PT Trends
 Influenced
by 3 factors:
1. Energy Level
Higher energy levels further from
nucleus
2. Charge on nucleus (# protons)
More + charge pulls e-’s in closer
3. Shielding effect
#1. Atomic Size - Group trends

Going down a
group, atoms gain
another PEL
(floor)
 atoms
get…..
b
i
H
Li
g
Na
g
K
e
r
Rb
#1. Atomic Size - Period Trends
L
to R across period:
Here is an animation
to explain the trend.
More
p+ in nucleus
More e-’s occupy same energy level
stronger nuclear charge
 Pulls e- cloud closer to nucleus

atoms
Na
S
get….
Al
Si
P
m a
l
l
Mg
S Cl Ar
e
r
Trends of Atomic Radius
increases
decreases
Ions
p. 172
 Some compounds composed of
“ions”
Ion
- atom (or group of atoms) w/ + or -
charge
formed
when e- transferred btwn
atoms
(loses e-’s…+ ion)
Anion (gains e-’s… - ion)
Cation
Cation Formation
Na atom
1 valence e-
Effective nuclear
charge on
remaining e-’s
increases.
11p+
Remaining epulled closer to
nucleus. Ionic
size decreases.
Valence elost in ion
formation
Result: a smaller
sodium ion, Na+
Anion Formation
Chlorine
atom with 7
valence e17p+
One e- is added
to the outer
shell (from Na
for example).
Effective nuclear charge is
reduced and the e- cloud
expands.
A chloride ion is
produced. It is
larger than the
original atom.
#2. Trends in Ionization Energy
p.173
 Ionization
energy - energy required to
completely remove e- (from gaseous
atom)
 energy required to remove only 1st ecalled first ionization energy.
Ionization Energy
 second
IE is E required to remove
2nd eAlways greater than first IE
 third IE greater than 1st or 2nd IE

IE helps predict what ions elements form



Li 1+
Mg 2+
Al 3+
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
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
Cation Formation
11p+
Anion Formation
17p+
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 Effect
in outer PEL
“look thru” other
PEL’s to “see”
nucleus
 Stays same thru
blocks
 Greater influence on
IE than nuclear
charge
 e-’s
Shielding Trends
increases
remains constant
Ionization Energy - Group trends
p. 174
 going
down group
first
IE decreases b/c...
e- further from p+ attraction
more
“shielding”
Ionization Energy - Period trends
p. 174
 Same
period atoms have same:
#
energy levels
 “shielding” (within a block – slight decrease
btwn “s” and “p”)
Increasing
nuclear charge
IE generally increases left - right
 Exceptions…full
& 1/2 full orbitals
First Ionization energy
He
H
 He
greater IE than H.
 Both w/ 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
higher IE than
Li
=
H
Be
Li
shielding
(period)
 greater nuclear
charge
Atomic number
B
has lower IE than
Be
First Ionization energy
He
 greater
H
Be
B
Li
nuclear
charge
 shielding has greater
influence on IE
Slight decrease (“p”
e-)
 “p” e- removed
 “s” orbital ½ 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
breaks
the pattern, b/c
removing eleaves it w/ 1/2
filled p orbital
B
Li
Atomic number
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
First Ionization energy
He
Ne
N F
 Ne
has a lower
IE than He
 Both
H
C O
Be
B
Li
full but…
 Ne
more
shielding
 b/c
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 - more
shielding
 Greater
distance
Na
Atomic number
Trends in Ionization Energy (IE)
decreases
increases
Trends in Ionic Size: Cations
 Cations
 metals
lose e-’s
 Cations smaller
than atom they came
from
 lose
e-’s
 lose entire energy level.
 Cations
of representative elements
have noble gas config before them
Trends in Ionic size: Anions
 Anions
gain e-‘s
nonmetals
 Anions
bigger than atom they came
from
same
energy level
greater area nuclear charge needs to cover
Configuration of Ions
Ions always have noble gas
configurations (full outer level)
 Na atom is: 1s22s22p63s1

1+ Na ion: 1s22s22p6
 Same as Ne
 Forms
Configuration of Ions
 Non-metals
form ions by
gaining e-’s to achieve noble
gas configuration
configuration
them
of noble gas after
Ion Group trends
 Each
step down a
group adds energy
level
 Ions - bigger going
down
 more
energy levels
Li1+
Na1+
K1+
Rb1+
Cs1+
Ion Period Trends
 Across
period
 nuclear
charge increases
 Ions get smaller
 energy
level changes btwn anions
& cations
Li1+
B3+
Be2+
C4+
N3-
O2-
F1-
#3. Trends in Electronegativity
Electronegativity (EN)- tendency for
atom to attract e-’s when atom in
cmpd
 Sharing e-, but how equally?
 Element w/ big EN pulls e- towards
itself strongly!

Electronegativity Group Trend
 Further down group, farther
e- away from nucleus
plus
 more
more e-’s atom has
willing to share
 Low EN
Electronegativity Period Trend
 Metals
low
let e-’s go easily
EN
 Nonmetals
want more e-’s
e-’s from others
High EN
take
Trends in Electronegativity
decreases
0
increases
Chemistry Song "Elemental Funkiness" Mark Rosengarten
The Elements – Tom Lehrer