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The periodic law
Chapter 5
1
Why do we need a table?


To organize the elements
To show trends
2
Periodic

A repeating pattern
3
Mendeleev’s table
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1869 – Dmitri Mendeleev – Russian
Arranged the elements in order of
increasing mass and noticed that
chemical properties were periodic
Put the elements into groups
according to properties
4
Mendeleev vs. Meyer


1860s Mendeleev and German
Lothar Meyer each made an eight
column table.
Mendeleev left some blanks in his
table in order for all the columns to
have similar properties – he
predicted elements that hadn’t been
discovered yet.
5
Why similar properties?

Why did they group according to
properties and mass and not atomic
number or number of outer level
electrons?
6
Germanium
Mendeleev’s
predictions
Atomic mass = 72
Atomic mass = 72.60
High melting point
Melting point = 958 °C
Density = 5.5 g/cm3
Dark gray metal

Actual element
Density = 5.36 g/cm3
Gray metal
Mendeleev’s blank spots and his
ability to predict future elements
helped his table win acceptance.
7
Mendeleev’s table



Elements arranged in order of
increasing mass.
Properties are repeated in an
orderly, periodic, fashion.
Mendeleev’s periodic law – the
properties of the elements are a
periodic function of their masses.
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9
Mass mistakes?


In order for Mendeleev to arrange
his elements by properties, he had
to put tellurium and iodine in the
wrong order.
He explained this by assuming that
their masses hadn’t been measured
very accurately.
10
More mass mistakes?
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Nickel and cobalt
Argon and potassium
Better mass measurements just
confirmed the discrepancy
11
Explanation
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1913 – Henry Moseley
X-ray experiments revealed the
atomic number was the number of
protons
Modern periodic law – the
properties of the elements are a
periodic function of their atomic
numbers
12
Modern periodic table

An arrangement of the elements in
order of their atomic numbers so
that elements with similar
properties fall in the same column
or group.
13
Noble gases
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Not discovered on Earth until 1894 1900.
Group 18 was added to the table
14
Lanthanides
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Hard to separate
All have similar properties
Added to the table in the early
1900s
15
Actinides
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Discovered later
Also all have similar properties
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Periodicity

Elements in the same group
(column) have similar properties.
17
Chemical properties of an element

Are governed by the electron
configuration of an atom’s highest
energy level
18
Period length


Determined by the number of
electrons than can occupy the
sublevels being filled in that period.
Table 5-1
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Full periodic table

Table with f-block in place
20
1st period


1s sublevel being filled
1s can hold 2 electrons, so there
are 2 elements in the 1st period.
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2nd and 3rd periods


2s and 2p or 3s and 3p being filled
s and p sublevels can hold 8 total,
so there are eight elements in these
periods
22
4th and 5th periods


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Add d sublevels, which can hold 10
electrons
Need to fill 4s, 3d, and 4p – 18
electrons
18 elements in each period
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6th and 7th periods


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Add f-block, which holds 14
electrons
Fill 6s, 5d, 4f, 6p
Need 32 electrons
32 elements in each period
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Figure 5-5

Shows blocks
25
Electron configurations


Elements in columns 1, 2, and 1318 have their last electron added in
an s or p orbital.
Elements in columns 3-12 have
their last electron added in a d
level.
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The s-block elements: Groups 1 and 2


Chemically reactive metals
Group 1

Have 1 electron in outer s orbital

Coefficient represents period

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Row 2: 2s1, Row 3: 3s1, etc. (ns1)
Group 2

Have 2 electrons in outer s orbital

Coefficient represents period

Row 2: 2s2, Row 3: 3s2, etc. (ns2)
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Alkali metals
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Metals in group 1
Have silvery appearance
Soft enough to cut with a knife
Not found alone in nature
React violently with nonmetals
Melting point decreases as you go
down the table
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Alkaline-earth metals
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
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Group 2
Harder, denser, and stronger than
alkali metals
Higher melting points than alkalis
Less reactive
Not found alone in nature
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Hydrogen and helium

Hydrogen
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

Located above group 1 because of its
electron configuration
Not really in group 1, because its
properties don’t match
Helium


Has an electron configuration like
group 2 elements
In group 18 because it is unreactive
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Discuss
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
Page 133
Sample problem 5-1 and practice
problems
31
Discuss

Without looking at the periodic table, give
the group, period, and block in which the
element with the electron configuration
[Rn] 7s1 is located.
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
Group 1, 7th period, s block
Without looking at the periodic table, give
the group, period, and block in which the
element with the electron configuration
[He]2s2 is located.

Group 2, second period, s block
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d-block elements: Groups 3-12

End in d1 to d10.

Coefficients are one less than the
period

Example: Fe is in the 6th column of
transition elements in the 4th period, ends
in 3d6
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Transition elements
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Groups 3-12
Typical metallic properties

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Good conductors
High luster
Less reactive than alkalis and
alkaline-earths
Some are unreactive enough to
appear in nature
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p-block elements: groups 13-18

End in p1 to p6.

Coefficients are the same as the period
ns2np1
 Always have a full s-sublevel

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p-block elements
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Properties vary greatly
Includes all nonmetals except hydrogen
and helium


Includes all the metalloids
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

Solids, liquids and gases
Between metals and nonmetals
Brittle solids
Semiconductors – can conduct under certain
conditions
Includes some metals

Less reactive than alkalis and alkaline-earths
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Halogens
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Group 17
Most reactive nonmetals
Form compounds called salts
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f-block elements
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Lanthanides and actinides
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
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Endings are f1 to f14
Coefficients are two less than the
period
All actinides are radioactive
Those after neptunium are synthetic
38
Discuss
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Sample problems and practice
problems on pages 136, 138, and
139
With your group first, then join with
another group.
Do you have any questions?
39
Atomic radius

Ideally, the distance from the
center of the atom to the edge of
it’s orbital.


But, atoms are “fuzzy”, not clearly
defined.
Defined as one-half the distance
between the nuclei of identical
atoms that are bonded together.
40
Period trends – see figure 5-13
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As we move from left to right across
the table, we gain protons.
There is a greater positive charge
on the nucleus.
This greater charge pulls harder on
the outer electrons, pulling them in
closer.
The atom gets smaller.
41
Group trends
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As we move down the table, the
principle quantum number
increases.
When the principle quantum
number increases, the electron
cloud gets bigger.
The size of the atoms gets bigger.
42
Discuss

Which of the elements Li, Rb, K,
and Na has the smallest atomic
radius? Why?


Li, it is highest on the table
Which of the elements Zr, Rb, Mo,
and Ru has the largest atomic
radius? Why?

Rb, it is farthest to the left on the table
43
Ion

An atom or group of bonded atoms
that has a positive or negative
charge
44
Ionization

Any process that makes ions
45
Ionization energy (IE)


First ionization energy (IE1) – the
energy required to remove the most
loosely held electron.
Measured in kJ/mol
46
Ionization energy – see figure 5-15

Experimentally determined.

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Tends to increase as you move across a
row from left to right



From isolated atoms in the gas phase
Why group 1 is most reactive
Caused by higher charge
Tends to decrease as you move down a
column


Electrons are farther from nucleus
Shielding from inner electrons
47
Other Ionization Energies – see Table
5-3
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Energy required to remove other
electrons from positive ions.
IE2, IE3, etc
Get higher as you remove more
electrons

Less shielding
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Noble Gases
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Have High ionization energies
When a positive ion of another
element reaches a noble gas
configuration, its ionization energy
goes up.



Example: When K loses one electron,
it has Ar’s electron configuration
This makes it stable
Its IE2 is much higher than its IE1
49
Discuss

State in words the general trends in
ionization energies down a group
and across a period of the periodic
table.
50
Electron affinity

The energy change that occurs
when an electron is gained by a
neutral atom

Most atoms release energy


Represented by a negative number
Some atoms gain energy
Represented by a positive number
 These ions will be unstable


KJ/mol
51
Period trends – see figure 5-17

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Group 17 has most negative
electron affinity.
Tends to get more negative (release
more energy) as we move to the
right
Exceptions:

groups with full or half-full sublevels
are more stable
52
Group trends
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
Not as regular
Usually, electrons add with greater
difficulty as we move down
53
Adding additional electrons
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Second electron affinities are all
positive because it is more difficult
to add electrons to a negative ion.
If a noble gas configuration has
been reached, it is even more
difficult.
54
Discuss

State in words the general trends in
electron affinities down a group and
across a period of the periodic
table.
55
Ionic Radii

Cation – a positive ion


Ionic radius smaller than atomic radius
Anion – a negative ion

Ionic radius is larger
56
Period Trends – see figure 5-19

Metals form cations by losing
electrons



Ions are smaller
Radius decreases as we move across
Nonmetals form anions by gaining
electrons


Ions are larger
Radius decreases as we move across
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Group trends

Ionic radius increases as you go
down the table
58
Valence electrons
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Available to be lost, gained or
shared in the formation of chemical
compounds
In highest energy levels
For s-block, the group number is
the same as the number of valence
electrons
For the p-block, the group number
is 10 more than the number of
valence electrons
59
Electronegativity

The measure of the ability of an
atom in a compound to attract
electrons

The atom with higher electronegativity
pulls the electrons closer to itself
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Electronegativity trends (figure 5-20)


Increases left to right across the
rows
Decreases down the columns
61
Discuss

Explain why elements with high
(more negative) electron affinities
are also the most electronegative.
62
d- and f-block elements


Properties vary less and with less
regularity than others
Atomic radii

d-block


Usual patterns
f-block (unusual)
Increase across periods
 Decrease down groups

63
d- and f-block elements
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Ionization energy

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Ionic radii
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Increase across periods
d-block increases down groups
(unusual)
f-block decreases down groups
Cations have smaller radii
Electronegativity


d-block follows normal rules
f-block all have similar
electronegativities
64
Discuss


Among the main-group elements,
what is the relationship between
group number and the number of
valence electrons?
In general, how do the periodic
properties of the d-block elements
compare with those of the maingroup elements?
65
Prelab notes
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Precipitate – solid that falls out of a
solution
The formation of a precipitate
indicates there has been a chemical
change.
This means that there were ions
present that were free to react.
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