The Periodic Law Notes

1. The Periodic Law Notes
(Chapter 5)

2. I. History of the Periodic Table
About 70 elements were known by 1850
(no noble gases) but there didn’t appear
to be a good way of arranging or
relating them to study.

3. A. Mendeleev and Chemical
Periodicity
Mendeleev placed known information of
elements on cards (atomic mass,
density, etc…). He arranged them in
order of increasing atomic masses,
certain similarities in their chemical
properties appeared at regular intervals.
Such a repeating pattern is referred to
as periodic.

4. Dmitri Mendeleev 1869

5. http://www.chem.msu.su/eng/misc/mendeleev/welcome.html

6. 1. Mendeleev’s table was published in 1869.
2. He left blanks in his periodic table for
undiscovered elements and he predicted their
properties. Later elements were discovered
with properties he predicted!
3. Problems with his table – a few elements
did not fit – the atomic mass arrangement did
not match with other similar properties.
4. Recognition – Mendeleev never received
the Nobel Prize – the importance of the
Periodic Table was not realized in his lifetime!

7. B. Moseley and the Periodic
Law
1. Moseley, with Rutherford, in 1911 discovered a
new pattern. The positive charge of the nucleus
increased by one unit from one element to the next
when elements are arranged as they are now on
the Periodic Table
2. This discovery led to definition of atomic
number and the reorganization of the Periodic
Table based on atomic number not atomic mass.
3. The Periodic Law – the physical and chemical
properties of the elements are periodic functions of
their atomic numbers.
4. Moseley died at the age of 28 – victim of WWI

8. C. The Modern Periodic
Table
1. An arrangement of the elements in order
of their atomic numbers so that elements
with similar properties fall in the same
column (or group).
Groups: vertical columns (#1-18)
Periods: horizontal rows (# 1-7)
2. Periodicity – the similarities of the
elements in the same group is explained by
the arrangement of the electrons around the
nucleus.

9. Periodic Table Expanded View
The way the periodic
table is usually seen is a
compressed view, placing
the Lanthanides and
Actinides at the bottom of
the table.

10. Layout of the Periodic Table: Metals vs.
Nonmetals
1
IA
18
VIIIA
1
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8 9
VIIIB
10 11
IB
12
IIB
4
5
6
7
Metals
Nonmetals

11. Copyright © Houghton Mifflin Company. All rights reserved. 8 | 11
II. Electron Configurations and the
Periodic Table

12. A. The s-block Elements:
Groups 1 and 2
1. Group 1: Alkali metals
-soft silvery metals
-most reactive of all metals, never found free in
nature
-reacts with water to form alkaline or basic
solutions – store under kerosene
-whenever you mix Li, Na, K, Rb, Cs, or Fr with
water it will explode and produce an alkaline
solution
-ns1 (ending of all electron configurations for this
group) http://www.teachertube.com/members/viewVideo.php?title=Group_1_Alkali_metals_and_water&video_id=51068

13. 2. Group 2: Alkaline earth
metals
-less reactive than Alkali, but still react in water
to produce an alkaline solution
-never found free in nature
-harder, denser, stronger than alkali
- ns2 (ending of all electron configurations for
this group), because they have 2 electrons in
the s sublevel, this makes them a little less
reactive then the Alkali metals in group 1

14. B. The d-Block Elements:
Groups 3-12
-are all metals with metallic properties
(malleability, luster, good conductors, etc…);
are referred to as the Transition Metals
-Harder and denser than alkali or alkaline
-Less reactive than alkali or alkaline
-For the most part their outermost electrons are
in a d sublevel
-Exceptions to the electron configuration are
found in these groups (Ex: Ni, Pd, Pt)

15. Irregular Electron configurations
– sometimes the electron configuration is NOT
what we would predict it to be. Sometimes
electrons are moved because (l) it will result in
greater stability for that atom or (2) for some
unknown reason??
It is very important to define “stable” here.
STABLE means:
all degenerate (equal energy) orbital’s are FULL
all degenerate orbital’s are half-full
all degenerate orbital’s are totally empty.

16. Examples – draw the orbital notation (lines and
arrows) for the predicted electron configuration for
Cr #24:
However, the real E. C. is [Ar]4s13d5. The 4s1 electron
has been moved to achieve greater stability.
The elements with “irregular” electron configurations are
Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, La, Pt, Au, Ac, Gd, Th,
Pa, U, Np, and Cm. We will also use the actual electron
configuration (found on the periodic table at the back of
your book) instead of the predicted configuration for these
elements. You do NOT have to memorize these, they will
be highlighted or marked on your periodic table.

17. C. The p – Block Elements:
Groups 13 – 18
-Contain metals and nonmetals
-Metalloids, along zigzag line, have
characteristics of both metals and
nonmetals (many are good conductors
but are brittle). The metalloids are
boron, silicon, germanium, arsenic,
antimony, and tellurium.

18. 1. Group 17 - Halogens –
most reactive nonmetals
-7 electrons in outermost (s and p) energy
levels (that is why so reactive – only need
one electron to have 8)
-called the salt formers (they react vigorously
with metals to form salts) A salt is a metal
and a nonmetal bonded together.
-most are gases

19. 2. Group 18 - Noble gases –
unreactive
- 8 electrons in outermost s and p
energy levels
- all are gases
The s and p blocks are called the main
group or representative elements!

20. D. The f-Block Elements:
Inner Transition Metals
-final electrons fill an f sublevel
1. Lanthanides – shiny reactive metals;
Ce-Lu (fill the 4f sublevel)
2. Actinides – unstable and radioactive;
Th-Lr (fill the 5f sublevel)

21. Hydrogen and Helium - Oddballs
-Hydrogen is NOT an Alkali metal, it is a very
reactive gas. It is placed with the Alkali
metals because 1s1 is its electron
configuration.
-Helium is a Noble gas, it is unreactive, but it
does not have 8 electrons in outermost
energy level, because it only has 2 total
electrons!

22. III. Electron Configuration
and Periodic Properties
A. Valence Electrons – electrons in the
outermost s and p orbital’s; electrons
available to be lost, gained, or shard in
the formation of chemical compounds
When all of the valence orbital’s are full
(have 8 electrons), the atom tends to be
unreactive (like the Noble Gases)

23. # of Valence
Electrons Group # Ending Configuration
1 1 ns1 very reactive
2 2 ns2
3 13 np1
4 14 np2
5 15 np3
6 16 np4
7 17 np5 very reactive
8 18 np6 very unreactive

24. B. Atomic Radii
1. The size of an atomic radius
cannot be measured exactly
because it does not have a
sharply defined boundary.
However the atomic radius
can be thought of as ½ the
distance between the nuclei
of identical atoms joined in a
molecule
Atomic Radii Trend

25. 2. Period trend – atomic radii decrease as you
move across a period.
As you move across a period, from left to right,
the size of an atom decreases because the
nucleus is getting larger and more positive,
but it is still pulling on the same number of
energy levels. Which atom is larger?
(A) Zr or Sn (B) Li or Cs

26. 3. Group trend - atomic radii decrease as you move up a
group (or increase as you move down a group).
Shielding effect - an invisible barrier made of core electrons
serve to decrease the pull of the nucleus on the outer
(valence) electrons.
Shielding increases as you go down a group because there
are more core electrons.
Shielding is considered to be constant as you move across a
period because the number of energy levels is staying the
same.
Which atom has more shielding? (A) K or Ca (B) Na or K
Which atom is smaller? (A) N or P (B) Li or K

27. C. Ionization Energy (IE)
– a measure of the ease with which
an electron can be removed
from a gaseous atom or ion
(sometimes called ionization
potential).
1. First ionization energy – the
energy required to remove
one outermost electron from
a gaseous atom. Second
ionization energy – the
additional amount of energy
needed to remove an
outermost electron from the
gaseous +1 ion.
IE Trend

28. 2. Period trend – ionization energy increases as you move across
the period.
As you move across a period the ionization energy increases
because the atoms get smaller. Another way to think of it: the
number of valence electrons increases (the amount of energy
needed to remove one electron is less then what is needed to
remove 7 or 8 electrons).
3. Group trend – ionization energy increases as you move up a
group (or decreases as you move down a group).
In general, as you do down a group the ionization energy
decreases because the size of the atom is increasing and the
outermost electrons are further from the nucleus.
1. Which atom has the higher first ionization energy?
(A) Hf or Pt (B) Cl or Ar

29. D. Electron Affinity (EA)
– the energy change that
accompanies the addition
of an electron to a gaseous
atom. The desire of an
atom to have another
electron. Chlorine for
example would like to
have another electron and
therefore has a very high
electron affinity.
EA Trend

30. 1. Period trend – electron affinity increases as you move
across a period because atoms become smaller and the
nuclear charge increases. This means there is a
greater pull
from the nucleus.
2. Group trend – electron affinity increases as you move
up a group (or decreases as you move down a group)
because the size of the atom increases and shielding
increases. This means there is less pull from the
nucleus.
Example: Which element has the greater electron
affinity? Pb or Sn

31. E. Ionic Radii
1. Period trend – The size of ions
decreases as you move across a
period because you have more
protons pulling on the same number
of energy levels.
2. Group trend – the size of the ions
decrease as you move up a group (or
increase as you move down a group)
because the number of energy levels
increases.
Example: Which would be larger?
K+1 or Ge +4
Which is larger? P ion or Cl ion
Ionic Radii Trend

32. 3. Metals – metal ions (cations) always have a smaller radii than
its corresponding atom because:
a. it loses its outer energy level electrons (valence electrons)
b. the proton to electron ratio is greater in the ion than in the atom
sodium atom 11p+/11e- = 1.0
sodium ion 11p+/10e- = 1.1
The value of p+ ro e- ratio varies inversely to the size of the ion.
The sodium ion is smaller because it has a larger proton to
electron ratio.
Which would be larger? Potassium atom or potassium ion

33. 4. Nonmetals - nonmetal ions are always larger
than their corresponding atoms because:
a. repulsion between electrons
b. the p+ to e- ratio is less in the ion
chlorine atom 17p+/17e- = 1.0
chlorine ion 17p+/18e- = 0.94
The chlorine ion is larger because it has a smaller
proton to electron ratio.
Which would be larger? Sulfur atom or sulfur ion

34. F. Electronegativity (EN)
– the tendency of an
atom to attract
electrons to itself
when it is chemically
combined (bonded) to
another element.
Electronegativity Trend

35. In general, metals have low EN and nonmetals
have high EN. The actual amount of EN an atom
has is indicated by a number of the Pauling
Electronegativity Scale that goes from 0 to 4. Dr.
Linus Pauling set up this scale and gave the
element having the greatest EN an arbitrary
number of 4, and he assigned numbers to the
others relative to this element.
Flourine is the most electronegative element at 4.
(3.98) and Francium is the least electronegative at
0.7.

36. 1. Period trend - EN increases as you go across a period
(excluding the noble gases) because size decreases.
2. Group trend - EN increases as you go up a group (or
decreases as you go down a group) because there is less
pull from the nucleus as the electrons get further away.
Which would have the greater EN?
(A) Ca or Se (B) Be or O
Electronegativity enables us to predict what type of bond
will be formed when two elements combine.