Ch. 5 - The Periodic Table · Ch. 5 - The Periodic Table 0 50 100 150 200 250 0 5 10 15 20 Atomic...
Transcript of Ch. 5 - The Periodic Table · Ch. 5 - The Periodic Table 0 50 100 150 200 250 0 5 10 15 20 Atomic...
I II III
III. Periodic
Trends
(p. 140 - 154)
Ch. 5 - The Periodic Table
0
50
100
150
200
250
0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
A. Periodic Law
When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
0
50
100
150
200
250
0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
B. Reaction Patterns
Valence electrons: main group (s & p)
electrons in the outermost energy level
O: 2s2 2p4 = 6 valence electrons
S: 3s2 3p4 = 6 valence electrons
Se: 4s2 3d10 4p4 = 6 valence electrons!!
Similar valence e- within
a GROUP result in
similar chemical
properties in the group.
B. Reaction Patterns
Only main group electrons are in the
valence
This means that 8 is the largest number of
valence electrons possible. (2 + 6 = 8)
This is where the term “full octet” comes
from. A full octet is the most stable
type of electron configuration.
The noble gases are the only
elements that start with full
octets.
1 2 3 4 5 6 7 8
# valence electrons increases, left to right:
B. Reaction Patterns
Valence #s go
by column,
which is why
properties go
by column.
1+ 2+
Ions: charged atoms created when
electrons are gained or lost to create a full
octet (noble) configuration.
B. Reaction Patterns
3+ 4± 3- 2-1- 0
These
“common”
ion charges
are based on
valence
numbers.
Periodic Trends
Periodic trends are due to:
The organization of
electrons
(in energy levels, sublevels and
orbitals due to the laws of
quantum mechanics)
The fact that opposites attract:
The nucleus (positive charge) attracts the
electron cloud (negative charge)
+-
Periodic Trends
Attraction between charged particles
can be mathematically calculated with
Coulomb’s Law:
F𝑒𝑙𝑒𝑐 =kq
1q2
r2
F𝑒𝑙𝑒𝑐 = attraction (force)
k = a constant
r = radius (distance from nucleus to electrons)
q1 = positive charge
(protons)
q2 = negative charge
(electrons)
We’re not going to solve
this equation, but it is the
basis of most periodic
trends, so let’s take a look:
Periodic Trends
Attraction between charged particles
can be mathematically calculated with
Coulomb’s Law:
F𝑒𝑙𝑒𝑐 =kq
1q2
r2
The equation is saying two things:
More protons & electrons = more attraction
Smaller atom = more attraction
Periodic Trends
Vertical patterns:
Going down a column,
energy levels are being added.
For a higher energy level, the
valence electrons are at greater
distance (r) from the nucleus
(Higher energy levels = larger orbitals)
The added distance decreases
attractive force.
1
2
3
4
5
6
7
Periodic Trends
Vertical patterns:
When comparing elements in the
same GROUP (column/family)
Elements at the top of a group
(lower atomic number) hold onto
their valence electrons tightly.
Elements at the bottom of a group
(greater atomic number) hold onto
their valence electrons loosely.
1
2
3
4
5
6
7
Let me explain you a thing…
This next section,
your job is to listen
and UNDERSTAND,
not to write down.
Periodic Trends
It’s all about the Valence Electrons
Only outermost (valence) electrons are
involved in reactions.
Reactions happen because an atom
doesn’t have a stable (full octet)
configuration.
But actually the inner electrons are
always stable already.
Only the valence electrons need work.
Mg: [Ne-10] 3s2
Periodic Trends
Inner electrons are already stable, they don’t do
anything, let’s ignore them.
Valence electrons’ behavior (and element
properties) are based on their attraction to
the protons in the nucleus.
Valence electrons don’t feel the attraction
of all the positive charge in the nucleus.
The valence electrons are
“shielded” from the nucleus
by the inner electrons. Mg: [Ne-10] 3s2
Periodic Trends
The ten inner
electrons cancel
the charge of ten
protons.
Only two protons
are “effective”
Zeff = +2
Mg: [Ne-10] 3s2
Result: Zeff = valence
Zeff = effective nuclear charge
(That’s important.)
Mg: [1s22s22p6] 3s2
10p+
1st energy level shield: 1s2
2p+
2nd E.L. shield: 2s22p6
Periodic Trends
Conclusion:
The shielding effect causes the values
of q1 and q2 to count from 1 to 8 (based
on valence number) for every period.
Which is why patterns restart every
period even though proton number
doesn’t “restart.”
1 2 3 4 5 6 7 8
Periodic Trends
Horizontal patterns:
Going across a period, valence
electrons and Zeff increase. (q1 and q2)
Greater charge (q1 & q2) increasesattractive force.
In a period, higher atomic number =
greater attraction between nucleus and
electron cloud.
1 2 3 4 5 6 7 8
F𝑒𝑙𝑒𝑐 =kq1q2r2
Periodic Trends
Horizontal patterns:
When comparing elements in the
same PERIOD (row)
Elements at the beginning of a
period (lower atomic number) hold
onto their valence electrons loosely.
Elements at the end of a period
(greater atomic number) hold onto
their valence electrons tightly.
1 2 3 4 5 6 7 8
Periodic Trends
Rule of thumb:
Up and to the right, electrons are held
more tightly.
Lower and to the left, electrons are held
more loosely.
C. Johannesson
Atomic Radius
0
50
100
150
200
250
0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
D. Atomic Radius
Li
ArNe
K
Na
Atomic Radius: SIZE
the distance from nucleus to edge of electron cloud
Or: half the distance between two adjacent nuclei
D. Atomic Radius
Increases to the LEFT in a period
Increases DOWN in a group
D. Atomic Radius
Why larger going
down?
More energy levels
Why smaller to the
right?
Increased nuclear
charge (without
additional shielding)
pulls e- in tighter
D. Atomic Radius
Ionization Energy: Energy required to remove one e- from a neutral atom.
The more stable an atom is, the more energy is required to ionize it.
i.e., the closer to a full octet, the higher the ionization energy.
D. Ionization Energy
First Ionization Energy
0
500
1000
1500
2000
2500
0 5 10 15 20Atomic Number
1s
t Io
niz
ati
on
En
erg
y (
kJ
)
E. Ionization Energy
KNaLi
Ar
NeHe
(First) Ionization Energy
Increases to the right in a period
Increases to the top in a group
E. Ionization Energy
Why is this opposite of atomic radius?
In small atoms, e- are close to the nucleus
where the attraction is stronger
Stronger attraction means more difficult to
remove electrons from.
Why small jumps within each group?
Stable e- configurations (full and half-full
sublevels) don’t want to lose electrons
E. Ionization Energy
First Ionization Energy
0
500
1000
1500
2000
2500
0 5 10 15 20Atomic Number
1s
t Io
niz
ati
on
En
erg
y (
kJ
)
E. Ionization Energy
Full
sublevel
Half-full
sublevel
Successive Ionization Energies
1st I.E. 736 kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ
Large jump in I.E. occurs when a CORE
(non-valence) e- is removed.
Mg = [Ne-10] 3s2
E. Ionization Energy
Al 1st I.E. 577 kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ
Successive Ionization Energies
Where is the “jump” for aluminum?
Al = [Ne-10] 3s2 3p1
3 v.e. means it should jump between 3rd & 4th ionizations.
E. Ionization Energy
Ionic RadiusCations (+)
Created by losing e-
smaller than parent
atom
Anions (–)
Created by
gaining e-
larger than parent
atom
G. Ionic Radius
Which atom has the larger radius?
Be or Ba
Ca or Br
Ba
Ca
Examples
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne
N
Ne
Examples
Which particle has the larger radius?
S or S2-
Al or Al3+
S2-
Al
Examples