Transcript Chapter 14 - Chemical Periodicity

Chapter 12
Chemical Periodicity
Killarney High School
Section 12.1
Classification of the Elements
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OBJECTIVES:
• Explain why you can infer the
properties of an element based on
those of other elements in the
periodic table.
Section 12.1
Classification of the Elements
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OBJECTIVES:
• Use electron configurations to
classify elements as noble gases,
representative elements, transition
metals, or inner transition metals.
Periodic Table Revisited
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Russian scientist Dmitri Mendeleev
taught chemistry in terms of
properties.
Mid 1800’s - molar masses of
elements were known.
Wrote down the elements in order
of increasing mass.
Found a pattern of repeating
properties.
Mendeleev’s Table
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Grouped elements in columns by
similar properties in order of
increasing atomic mass.
Found some inconsistencies - felt that
the properties were more important
than the mass, so switched order.
Also found some gaps.
Must be undiscovered elements.
Predicted their properties before they
were found.
The modern table
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Elements are still grouped by
properties.
Similar properties are in the same
column.
Order is by increasing atomic
number.
Added a column of elements
Mendeleev didn’t know about.
The noble gases weren’t found
because they didn’t react with
anything.
Periodic Table Video
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Horizontal rows are called
periods
There are 7 periods
Vertical columns called groups
Elements are placed in
columns by similar properties
Also called families
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1A
2A
The elements in the A
8A
0
groups are called the
3A 4A 5A 6A 7A
representative elements
outer s or p filling
The group B are called the
transition elements
 These
are called the inner
transition elements, and they
belong here
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Group 1A are the alkali metals
Group 2A are the alkaline earth
metals
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Group 7A is called the Halogens
Group 8A are the noble gases
Why Do Some Elements Share Properties?
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The part of the atom another atom
sees is the electron cloud.
More importantly the outside
orbitals.
The orbitals fill up in a regular
pattern.(Aufbau Principle)
If the outside orbital electron
configuration repeats.
The properties of atoms repeat.
H
Li
1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10
5p66s1
1s22s22p63s23p64s23d104p65s24d105
p66s24f145d106p67s1
1s2 He 2
Ne
2
2
6
1s 2s 2p
10
1s22s22p63s23p6 Ar18
1s22s22p63s23p64s23d104p6 Kr
36
1s22s22p63s23p64s23d104p65s24d105p6 Xe
54
1s22s22p63s23p64s23d104p65s24d10 Rn
5p66s24f145d106p6 86
S- block
s1
s2
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Alkali metals all end in s1
Alkaline earth metals all end
in s2
really should include He, but it
fits better later.
He has the properties of the
noble gases.
Transition Metals -d block
s1
1
d
2
d
3
d
d4
Exception!!
s1
5
6
7
8
10
9
d d d d d d
The P-block
p1 p2
p3
p4
p5
p6
F - block
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inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1
2
3
4
5
6
7
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Each row (or period) is the energy
level for s and p orbitals.
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d orbitals fill up after previous energy level,
so first d is 3d even though it’s in row 4.
1
2
3
4
5
6
7
3d
1
2
3
4
5
6
7
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6
7
f orbitals start filling at 4f
4f
5f
Summary: Fig. 12.5, p. 277
Sample Problem 12-1, p.278
The Periodic Law states that
when elements are arranged in
order of increasing atomic
number there is a periodic
pattern in their physical and
chemical properties.
Writing electron
configurations the
easy way
Yes there is a shorthand
Electron Configurations repeat
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The shape of the periodic table is
a representation of this
repetition.
When we get to the end of the
column the outermost energy
level is full.
This is the basis for our
shorthand.
The Shorthand
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Write symbol of the noble gas
before the element, in [ ].
Then, the rest of the electrons.
Aluminum’s full configuration:
1s22s22p63s23p1
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previous noble gas Ne is:
1s22s22p6
so, Al is: [Ne] 3s23p1
More examples
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Ge = 1s22s22p63s23p64s23d104p2
• Thus, Ge = [Ar] 4s23d104p2
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Hf =
1s22s22p63s23p64s23d104p65s2
4d105p66s24f145d2
• Thus, Hf = [Xe]6s24f145d2
The Shorthand Again
Sn- 50 electrons
The noble gas
before it is Kr
Takes care of 36
Next 5s2
Then 4d10
Finally 5p2
[ Kr ] 5s2 4d10 5p2
Assignment
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Objective Ws 1- 6
Questions 10-15 pg 295
Section 12.2
Periodic Trends
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OBJECTIVES:
• Interpret group trends in atomic
radii, ionic radii, ionization
energies, and electronegativities.
Section 12.2
Periodic Trends
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OBJECTIVES:
• Interpret period trends in atomic
radii, ionic radii, ionization
energies, and electronegativities.
Periodic Trends Online Lab
Atomic Size Video
Trends in Atomic Size
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First problem: Where do you
start measuring from?
The electron cloud doesn’t have
a definite edge.
They get around this by
measuring more than 1 atom at
a time.
Atomic Size
}
Radius
Atomic Radius = half the distance between
two nuclei of a diatomic molecule.
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Trends in Atomic Size
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Influenced by three factors:
1. Energy Level
• Higher energy level is further
away.
2. Charge on nucleus
• More charge pulls electrons in
closer.
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3. Shielding effect
(blocking effect?)
Group trends
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As we go down
a group...
each atom has
another energy
level,
so the atoms
get bigger.
H
Li
Na
K
Rb
Periodic Trends
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As you go across a period, the radius
gets smaller.
Electrons are in same energy level.
More nuclear charge.
Outermost electrons are closer.
Na
Mg
Al
Si
P
S Cl Ar
Rb
K
Atomic Radius (nm)
Overall
Na
Li
Kr
Ar
Ne
H
10
Atomic Number
Trends in Ionization Energy
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The amount of energy required
to completely remove an
electron from a gaseous atom.
Removing one electron makes a
1+ ion.
The energy required to remove
the first electron is called the
first ionization energy.
Ionization Energy
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The second ionization energy is
the energy required to remove
the second electron.
Always greater than first IE.
The third IE is the energy
required to remove a third
electron.
Greater than 1st or 2nd IE.
Ionization Energies Table 12.1, p. 281
kJ/mol
Symbol First
Atom
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
11810 Ion Formed
14840 Be 2+
3569
4619
4577
5301
6045
6276
What determines IE
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The greater the nuclear charge,
the greater IE.
Greater distance from nucleus
decreases IE
Filled and half-filled orbitals have
lower energy, so achieving them
is easier, lower IE.
Shielding effect
Shielding
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The electron on the
outermost energy
level has to look
through all the other
energy levels to see
the nucleus.
Second electron has
same shielding, if it
is in the same period
+
Group trends
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As you go down a group, first
IE decreases because...
The electron is further away.
More shielding.
Periodic trends
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All the atoms in the same period
have the same energy level.
Same shielding.
But, increasing nuclear charge
So IE generally increases from
left to right.
Exceptions at full and 1/2 full
orbitals.
First Ionization energy
He
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H
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He has a greater
IE than H.
same shielding
greater nuclear
charge
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
First Ionization energy
H to Br
kJ/mol
Atomic number
Driving Force
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Full Energy Levels require lots of
energy to remove their
electrons.
Noble Gases have full orbitals.
Atoms behave in ways to achieve
noble gas configuration.
2nd Ionization Energy
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For elements that reach a filled
or half-filled orbital by removing
2 electrons, 2nd IE is lower than
expected.
True for s2
Alkaline earth metals form 2+
ions.
3rd IE
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Using 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!!!
Trends in Electron Affinity
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The energy change associated with
adding an electron to a gaseous
atom.
Easiest to add to group 7A.
Gets them to full energy level.
Increase from left to right: atoms
become smaller, with greater nuclear
charge.
Decrease as we go down a group.
Trends in Ionic Size
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Cations form by losing electrons.
Cations are smaller than the
atom they come from.
Metals form cations.
Cations of representative
elements have noble gas
configuration.
Ionic size
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Anions form by gaining
electrons.
Anions are bigger that the atom
they come from.
Nonmetals form anions.
Anions of representative
elements have noble gas
configuration.
Configuration of Ions
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Ions always have noble gas
configuration.
Na is: 1s22s22p63s1
Forms a 1+ ion: 1s22s22p6
Same configuration as neon.
Metals form ions with the
configuration of the noble gas
before them - they lose electrons.
Configuration of Ions
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Non-metals form ions by gaining
electrons to achieve noble gas
configuration.
They end up with the
configuration of the noble gas
after them.
Group trends
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Adding energy level
Ions get bigger as
you go down.
Li1+
Na1+
K1+
Rb1+
Cs1+
Periodic Trends
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Across the period, nuclear
charge increases so they get
smaller.
Energy level changes between
anions and cations.
3Li1+
B3+
Be2+
C4+
N
O2-
F1-
Size of Isoelectronic ions
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Iso- means the same
Iso electronic ions have the
same # of electrons
Al3+ Mg2+ Na1+ Ne F1- O2- and
N3all have 10 electrons
all have the configuration:
1s22s22p6
Size of Isoelectronic ions
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Positive ions that have more
protons would be smaller.
Al3+
Na1+
Mg2+
Ne
F1-
2O
N3-
Electronegativity
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The tendency for an atom to attract
electrons to itself when it is
chemically combined with another
element.
How fair is the sharing?
Big electronegativity means it pulls
the electron toward it.
Atoms with large negative electron
affinity have larger electronegativity.
Electronegativity Video
Group Trend
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The further down a group, the
farther the electron is away, and
the more electrons an atom has.
More willing to share.
Low electronegativity.
Periodic Trend
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Metals are at the left of the
table.
They let their electrons go
easily
Low electronegativity
At the right end are the
nonmetals.
They want more electrons.
Try to take them away from
Summary: Fig. 12.10, p.285
Look at the Families
Group IA - The Alkali Metals
(Li, Na, K, Rb, Cs, Fr)
Highly
colored in
flames =
fireworks
Group IIA - The Alkaline Earth Metals
(Be, Mg, Ca, Sr, Ba, Ra)
Lose 2 valence electrons
Also react with H2O to
form an alkaline solution
(basic), and hydrogen gas,
but less violently
Ca(s) + 2H2O(l) 
Ca(OH)2(aq) + H2(g
Group IIA - The Alkaline Earth Metals
(Be, Mg, Ca, Sr, Ba, Ra)
Various forms of CaCO3
Group IIA - The Alkaline Earth Metals
(Be, Mg, Ca, Sr, Ba, Ra)
Strong reaction of
magnesium with oxygen to
produce magnesium oxide
Mg(s) + O2(g)  MgO(s)
flashbulbs
Group IIIA - The Boron Family
(B, Al, Ga, In, Tl)
Boron is mined in the
form of Borax, and is used
in laundry soap
Laboratory glassware
contains borosilicates
Group IIIA - The Boron Family
(B, Al, Ga, In, Tl)
Aluminum metal
is the most
abundant metal
in the earth’s
crust and has
many uses
Group IIIA - The Boron Family
(B, Al, Ga, In, Tl)
Gallium Arsenide is used
in the manufacture of
computer chips
Group IVA - The Carbon Family
(C, Si, Ge, Sn, Pb)
Carbon is essential for
life and is found in all
organic molecules
Group IVA - The Carbon Family
(C, Si, Ge, Sn, Pb)
Carbon is found in different structures or
ALLOTROPES - graphite, one of the softest
substances known, and diamond, the hardest
Group IVA - The Carbon Family
(C, Si, Ge, Sn, Pb)
Quartz or SiO2
Elemental Si is used in
the semiconductor
industry
Group VA - The Nitrogen Family
(N, P, As, Sb, Bi)
There are two
varieties of P,
red and white.
Group VIA - The Oxygen Family
(O, S, Se, Te, Po)
Stratospheric ozone shields us from harmful UV
radiation. Ozone is destroyed by Cl-containing
molecules used in refrigeration
The ozone
“hole” over
Antarctica
Group VIIA - The Halogens
(F, Cl, Br, I, At)
Br2 and I2
Halogen mean “salt-former”. Here sodium
metal reacts vigorously with Cl2(g)
Group VIIIA - The Noble (Inert) Gases
He, Ne, Ar, Kr, Xe
The Lights of Las Vegas
Helium-Neon lasers
The Transition Elements
An important use of transition elements is as pigments in
paints and glasses