Chapter 6 The Periodic Table
Download
Report
Transcript Chapter 6 The Periodic Table
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 ep+ (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 eAlways 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 eNoble 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 N3all 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
0
Chemistry Song "Elemental Funkiness" Mark Rosengarten