Transcript Chapter 12 The Periodic Table

The Periodic Table
Chapter 5
Video clip
What Does Periodic Mean?
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To answer:
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What things occur periodically?
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Turn to the picture on pg. 133, Fig. 1
Moon phases, magazine publications
Keep these in mind as we learn about the
periodic table
History
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Russian scientist Dmitri Mendeleev taught
chemistry in terms of properties.
Mid 1800 - molar masses of elements were
known.
Wrote down the elements in order of
increasing mass.
Found a pattern of repeating properties.
The Table
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Compare this
table with our
modern table
The Table
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
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Example: Iodine after Tellurium
Found some gaps.
Must be undiscovered elements.
Predicted their properties before they were
found.
Interestingly…
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Mendeleev never won a Nobel Prize.
He was nominated shortly before his death,
but lost to Henri Moissan, who discovered
Fluorine in 1906
Periodic Law
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Why could most of the elements be arranged
in the order of increasing mass, but a few
could not?
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Henry Moseley- discovered that atomic number
should be the basis for the table, not mass
Example: I = 53, Te = 52
Physical and chemical properties of the
elements are periodic functions of their
atomic numbers
The Modern Table
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Elements are still grouped by properties.
Similar properties are in the same column.
Order is in 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.
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So a new group was formed
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Horizontal rows are called periods
There are 7 periods
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Vertical columns are called groups.
Elements are placed in columns by
similar properties.
Also called families
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1A
2A
The elements in the A groups are
called the representative elements
3A 4A 5A 6A 7A
8A
0
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
Review
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Rewrite the periodic law in your own words
What group are the Alkali Metals in?
Halogens?
Why do families have the same
properties?
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The part of the atom another atom “sees” is the
electron cloud.
More specifically the outside (valence) orbitals.
A “family’s” orbitals fill up in a regular pattern.
The outside orbital electron configuration
repeats.
The properties of atoms repeat.
H
Li
1
1s1
1s22s1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10
5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s
24f145d106p67s1
1s2 He
2
1s22s22p6 Ne
10
1s22s22p63s23p6 Ar
18
1s22s22p63s23p64s23d104p6 Kr
36
1s22s22p63s23p64s23d104p65s24d105p6
Xe
1s22s22p63s23p64s23d104p65s24d10 54
5p66s24f145d106p6 Rn
86
S- block: Reactive Metals
s1
s2
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Alkali metals (column 1) all end in s1
Alkaline earth metals (column 2) all
end in s2
really have to include He but it fits
better later.
He has the properties of the noble
gases.
Demo- Ca and Mg in test tube (pg.
142)
Transition Metals -d block
d1
d2
d3
d4
d5
d6
d7
d8
d9
d10
The P-block: p1 p2 p3
Main Group
Elements
p4
p5
p6
F - block
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f1
inner transition elements
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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4f
f orbitals start filling at 4f
5f
Review
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Create an acronym for remembering the
position of the s, p and d blocks of the table
The s-Block Elements
Group 1
Group 2
1 s electron
2 s electrons
Alkali metals
Alkaline-earth
metals
Extremely reactive
A little less reactive
silvery
vary
soft
Stronger, harder
Not found in nature
Not found in nature
Melt at lower temps
Melt at higher temps
Quantum Formula: ns1,2
Quantum Formula: ns1,2
Example problems
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Sample Problem A on pg. 143
Without looking at the periodic table, identify the
group, period, and block that [Xe] 6s2 is located
Answer: group 2, sixth period, s block
Write the electron config. For the Group 1 element in
the third period. Will it be more reactive or less?
Answer: Grp 1, third period= 1s22s22p63s1
Must be more reactive, because it’s in
group 1
Quantum Formula: ns0-2 (n-1)d1-10
The d-Block
Groups 3-12
5 orbitals, 10 e’ total
Transition metals
Extremely reactive
Conduct electricity
High luster
Less reactive
Not always the same outer
e’configuration
d block
Example Problem
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Sample Problem B on pg. 146
For d-block problems, identify group using
this formula: d + s
Without using the periodic table, identify the
period, block and group of [Kr] 4d55s1
Answer: 5 period, d-block, group 6
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Molybedenum
Quantum Formula: ns0-2(n-1)d1-10
p- Block
Groups 13-18
Group number – 10 for
electrons
Main group elements
Properties vary
Nonmetals, metalloids,
halogens
Not always the same
outer e’configuration
Quantum Formula: ns2np1-6
Example
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Sample Problem C on pg. 148
Write the outer electron configuration for the Group
14 element in the second period, name it and
identify it as a metal, nonmetal or metalloid
Group number is higher than 12- so it’s the P block,
second period makes n=2, group number -10 = 1410= 4
So, you have 2 left over for the p’s
2s22p2
Quantum Formula: ns2np1-6
Halogens (p-block)
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Group 17
Gases, mostly
Most reactive of nonmetals
Seven valence electrons
Metalloids (p-block)
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Groups btw metals and nonmetals
solids
Electrical conductivity
Metals (p-block)
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Harder and denser than alkaline-earth
metals, but softer than d-block
Found only in compounds
Quantum Formula: ns2np1-6
Problem Examples
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Without the table, write the outer electron
config. For Group 14 in second period.
Name it and classify it as a metal, nonmetal
or metalloid
Answer: 14= p block, 14-10= 4 so 2s22p2,
carbon= nonmetal
f- Block
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Sixth and seventh periods
La-Hf = Cerium-Lutetium
Ac-Rf = Thorium-Lawrencium
Mostly lab made (sythetic)
1
2
3
4
5
6
7
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4f
f orbitals start filling at 4f
5f
More Example Problems
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Name the block and group for each and identify as metal,
nonmetal or metalloid
[Xe]4f145d96s1
Answer: d-block, group 10, Pt, metal (period 6)
[Ne]3s23p6
Answer: p-block, group 18, Ar, nonmetal (period 3)
Group
#
Group
Config
Block Comments
1, 2
ns1,2
s
3-12
ns0-1(n-1)d1-10 d
Sum of electrons in ns and (n-1)d
equals group number
13-18
ns2np1-6
Number of electrons in np sublevel
equals group number +/- 10
p
One or two electrons in ns
Metals
Alkali
Alkali-Earth
Transition
Nonmetals
Metalloids Halogens
Nobel gases
Driving Force of Atoms
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Full Energy Levels are very stable
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas
configuration.
Atomic Size
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First problem where do you start measuring.
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.
What’s A Trend?
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Name some fashion trends
Color trends?
Behavior trends?
Trends in Atomic Size
Influenced by two factors.
 Energy Level
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Higher energy level is further away.
Charge on nucleus
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More charge pulls electrons in closer.
Periodic Trends
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As you go across a period the radius gets
smaller.
They have the same energy level, though.
More nuclear charge.
Outermost electrons are closer.
Na
Mg
Al
Si
P
S Cl Ar
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
Ionic Size
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Cations form by losing electrons.
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Form positive ions
Groups 1-3
Cations are smaller than the atom they come
from.
Metals form cations
Ionic size
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Anions form by gaining electrons.
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Form negative ions
Anions are bigger than the atom they come
from.
Nonmetals form anions.
Ionization Energy- Pg. 153
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The amount of energy required to remove an
electron from an atom (only deals with losing an e’)
Removing one electron makes a +1 ion.
The energy required is called the first ionization
energy.
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Measured in kilojoules per mole
If an atom has a low IE, it will release an electron
easier than one with a high IE (making them more
reactive)
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So, which group would have the highest IE?
 Nobel gases
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 of 2nd IE.
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
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
Group trends
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As you go down a group first IE
decreases because
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The electron is further away.
More shielding occurs
What’s shielding?
Shielding
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The electron on the outside
energy level has to look
through all the other
energy levels to see the
nucleus
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Inner shell electrons “shield”
nuclear charge from outer shell
electrons
Period trends
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All the atoms in the same period have the
same energy level.
Same shielding.
Increasing nuclear charge
So IE generally increases from left to right
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A higher charge will more strongly attract
electrons, holding them “hostage”
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
First Ionization energy
He
Li has lower IE than
H
 more shielding
 further away
 outweighs greater
nuclear charge
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H
Li
Atomic number
First Ionization energy
He
Be has higher IE
than Li
 same shielding
 greater nuclear
charge
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H
Be
Li
Atomic number
First Ionization energy
He
B has lower IE than
Be
 same shielding
 greater nuclear
charge
 By removing an
electron we make s
orbital half filled
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H
Be
B
Li
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
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N
H
C O
Be
Breaks the pattern
because removing
an electron gets to
1/2 filled p orbital
B
Li
Atomic number
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
Ne
First Ionization energy
He
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N F
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H
C O
Be
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Ne has a lower IE
than He
Both are full,
Ne has more
shielding
Greater distance
B
Li
Atomic number
Ne
First Ionization energy
He
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N F
Na has a lower
IE than Li
Both are s1
 Na has more
shielding
 Greater distance
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H
C O
Be
B
Li
Na
Atomic number
Atomic number
First Ionization energy
Electron Affinity- Pg 157
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The energy associated with adding an
electron to an atom (only deals with gaining
e’)
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Trends:
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Easiest to add to group 7A.
A highly negative number = a high EA, that
means the atom will gain electrons easily
EA increases from left to right because atoms
become smaller, with greater nuclear charge.
EA decrease as we go down a group.
Also measured in kJ/mol
Electronegativity- Pg. 161
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The tendency for an atom to attract electrons to
itself when it is chemically combined with another
element.
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How fair it shares the electron with the atom its bonding
with
Big electronegativity means it pulls the electron
toward it- it’s a bully!
Atoms with large negative electron affinity have
larger electronegativity.
Flouine is the boss!
So, Chlorine isn’t
sharing the
electron fairly
with sodium,
because it has
such a large
electronegativity
Valence Electron
Na
Cl
Electronegativity= 3.0
Electronegativity= 0.9
So, Chlorine isn’t
sharing the
electron fairly
with sodium,
because it has
such a large
electronegativity
Valence Electron
Na
Cl
Electronegativity= 3.0
Electronegativity= 0.9
Group Trend
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The further down a group, the farther the
electron is away from the nucleus because
there are more energy levels
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Therefore, atoms are more willing to share these
electrons.
Low electronegativity.
Periodic Trend
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Cations let their electrons go easily
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Low electronegativity
Anions want more electrons- try to steal them
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High electronegativity
Ionization energy, electronegativity
Electron affinity INCREASE
Atomic size increases,
shielding constant
Ionic size increases