At the end of this chapter you should be able to sketch the periodic table showing the groups and periods; identify the metals, metalloids and non-metals in the periodic table. Identify the representative elements, the transition elements, the transition metals, the lanthanides and actinides in the periodic table. Also, give the electron configuration of cations and anions; determine the trends in the physical properties of elements in a group; describe and explain the trends in atomic properties in the periodic table; compare the properties of families and elements; predict the properties of individual elements based on their position in the periodic table; and perform exercises and collaborative work with peers.

1. PERIODIC
RELATIONSHIPS AMONG
THE ELEMENTS
Prepared by:
Eden C. Sanchez

2. 1. Sketch the periodic table showing the groups
and periods.
2. Identify the metals, metalloids and nonmetals
in the periodic table.
3. Identify the representative elements, the
transition metals, the lanthanides and actinides
in the periodic table.
Objectives

3. 4. Give the electron configuration of cations and
anions.
5. Determine the trends in the physical properties
of elements in a group
6. Describe and explain the trends in atomic
properties in theperiodic table
7. Compare the properties of families of elements
Objectives

4. 8. Predict the properties of individual elements
based on their position in the periodic table
9. Perform exercises and collaborative work
with peers.
Objectives

5. a. Periodic table
b. Metals
c. Non-metals
d. Metalloids
e. Alkali metals
f. Alkaline earth metals
g. Halogens
Keywords

6. Keywords
h. Noble gases
i. Representative elements/main group elements
j. Transition elements
k. Lanthanides
l. Actinides
m. Isoelectronic
n. Effective nuclear charge

7. o. Shielding or screening
p. Atomic radius
q. Ionic radius
r. Ionization Energy (First, Second, Third,…)
s. Electron affinity
Keywords

8. Development of the Periodic Table
 The arrangement of elements in the modern
periodic table was made possible through the
efforts of several chemists, such as; Dobereiner,
John Newlands, Dmitri Mendeleev, and Henry
Moseley.
 Dobereiner’s “Law of Triads” – relationship among
three elements where the atomic weight of the
middle element is nearly the same as average of the
atomic weights of other two elements.

9. Development of the Periodic Table
 John Newlands arranged the elements in what is
known as the “law of octaves”. He noted that
the eighth element has similar chemical
properties with the first element.
 Mendeleev prepared a tabulation of elements
based on equivalent weights (atomic mass) and
the regular recurrence of properties of the
elements.

10. Development of the Periodic Table
 In a few cases, the mass and the properties did
not go the same directions. But Mendeleev
rationalized that the properties were more
accurate than the masses since technology used
to determine the mass was still improving.

11. Development of the Periodic Table
 Henry Moseley discovered that each element
in Mendeleev’s table was arranged in an order
such that their integral positive charge
(atomic number) increased numerically from
left to right and top to bottom.

12. Development of the Periodic Table
 The present periodic table is arranged
according to increasing atomic number which
also equals the number of electrons.
 The electron configuration helps to predict
and explain the recurrence of chemical and
physical properties.

13. PERIODIC CLASSIFICATION OF
ELEMENTS

14. PERIODIC CLASSIFICATION OF
ELEMENTS
The periodic table is a chart in which elements
having similar chemical and physical properties
are grouped together.
The elements are arranged according to
increasing atomic number.
The rows are called periods.
The vertical columns are called groups or families

15. At present, it contains 118 elements; however,
elements 113 to 118 have just recently been
synthesized and naming is not yet fully complete.
There are 18 groups or families.
There are two conventions in designating the
groups: The Union of Pure and Applied
Chemistry (IUPAC) refers to the columns are
Groups 1-18, and Groups A and B convention such
as Groups 1A, 2A, 3B, and so on.

16. Majority of the elements are metals (good
conductors of electricity), followed by non-
metals and metalloids (have properties that are
ntermediate between metals and nonmetals)
Some groups have been given collective names.
Group 1A elements are called alkali metals;
Group 2A elements are referred to as alkaline
earth metals; Group 7A elements are called
halogens; Group 8A elements are known as
noble gases.

17. The Group A elements are classified as
representative elements or main group
elements – have unfilled or filled s and p
orbitals in the highest principal quantum
number.
The Group B elements are called the
transition elements where the d subshells are
being filled up.

18. The two separate rows at the bottom of the
periodic table are lanthanides and the actinides,
referred to as the f-block elements.

19. The Periodic Table: Representative,
Transition, and f-block Elements

20. 1. Write the electron configuration (using
noble gas notation) of the elements in
Group 1A.
2. Comment on the outermost electron
configuration of Group 1A elements.
3. How many valence electrons do Group 1A
elements have?
Seatwork:

21. Seatwork:
4. Write the electron configuration (using noble
gas notation) of the halogens.
5. Comment on the outermost electron
configuration of the halogens.
6. How many valence electrons do the halogens
have?
7. Comment on the arrangement of the
representative elements in the periodic table
with respect to their electron configuration.

22. • Ions derived from representative elements
In the formation of cations, the electrons are
removed from the outermost shell to achieve a
noble gas configuration.
In the formation of anions, electrons are added
to the highest partially filled n shell so that they
become isoelectronic (same number of
electrons) with the noble gas.
ELECTRON CONFIGURATION OF
CATIONS AND ANIONS

23. 1. Give the electron configuration of Na and
Na+.
2. Give the electron configuration of Ca and
Ca2+.
3. Give the electron configuration of F and F-.
4. the electron configuration of O and O2-.
Exercises

24. • Ions derived from transition elements
• Example: Mn

25. PERIODIC VARIATION IN PHYSICAL
PROPERTIES
1. The Effective Nuclear Charge
In many-electron atoms, the inner or core
electrons shield the outer electrons from the
nucleus reducing the electrostatic attractions
between the nucleus and the outer electron. The
effective nuclear charge, Zeff, is given by Zeff = Z - σ
where Z is the nuclear charge and σ is the shielding
constant.

26. Screening or shielding refers to how an outer
electron is blocked from the nuclear charge by
the inner electrons. It means that the attraction
of the outer electron to the nucleus is not felt
100% because of the effect of the inner electrons.
Electrons in the inner shells are very effective in
shielding the nucleus.

27. Zeff increases as you go from left to right
across a period.
3Li 4Be 5B 6C 7N 8O 9F 10Ne
Z 3 4 5 6 7 8 9 10
Zeff 1.28 1.91 2.42 3.14 3.83 4.45 5.10 5.76

28. 2. Atomic Radius
distance of the electron from the nucleus within
which 95% of the electron charge density is found.
A more specific way to get atomic radius values is to get
one-half the distance between two nuclei in adjacent
atoms (the internuclear distance) in a metal solid or in a
diatomic molecule.

29. The covalent radius is one-half the distance
between two identical atoms joined together
by a single bond.

30. The metallic radius – is one-half the distance
between the nuclei of the two atoms in contact in
the crystalline solid metal.

31. • The atomic radius decreases from left to right
through a period of elements for
representative elements. This corresponds to
the increase in Zeff across a period. With the
increase in Zeff, the outer electrons are pulled
in and attracted towards the nucleus resulting
in a decrease in the size of the atoms.

32. • The more electronic shells (n) in an atom, the
larger is the atom. Atomic radius increases
from top to bottom through a group of
elements.

33. Using the periodic table, arrange the
following atoms in order of increasing atomic
radius. Explain your reasoning.
a. C, Li, Be
b. As, I, S
c. P, Si, N
Exercises:

34. 3. Ionic Radius
Ionic radii are very difficult to measure with
certainty because they are affected by their
immediate environment
can be measured by x-ray diffraction
sizes vary depending on the environment

35. Atomic radius versus ionic radius
• Cations are smaller than the atoms from
which they are formed.
• For cations, the more positive the ionic charge,
the smaller the ionic radius.

36. • Anions are larger than the atoms from which they are
formed.
• The nuclear charge remains constant, but Zeff is reduced
because of the additional electrons. The additional
electrons results in increase repulsions among the
electrons in the outer shell, resulting the electrons to
spread out more, thus increasing the size of the anion.
• For isoelectronic anions, the more negative charge, the
larger is the ionic radius.

37. 1. Compare the size of a neutral atom of Na and a
Na+ ion. Which is larger? Explain.
2. Compare the size of a Mg atom and a Mg2+
ion. Which is larger?
3. Compare the sizes of Na+, Mg2+, and Al3+.
Arrange according to increasing size.
4. Compare the size of a F atom and a F– ion.
Which is larger?
Example

38. 5. Which is larger, the O atom or the O2– ion? Which is
larger?
6. Compare the sizes of F–, O2–, and N3–. Arrange
according to increasing size.
7. Arrange the following set of ions and atoms in
increasing size and explain your answer
a. K+, Cl-, S2-, Ca2+
b. N, Cs, As, Mg2+, Br-

39. 4. Ionization Energy
Ionization energy (IE) is the minimum amount
of energy (in kJ/mol) required to remove an
electron from a gaseous atom in its ground
state.
Energy + X(g) —> X+ (g) + e–

40. • The energy required to remove the first electron
is called the first ionization energy. The first
ionization energy, IE1, has the lowest value.
• The second ionization energy, IE2, is the energy
required to strip the second electron from the
atom; it has higher energy value, and so on.
• IE1 < IE2 < IE3 < …

41. • Examples of ionization energies:
Al(g) —> Al+(g) + e– IE1 = 577.9 kJ/mol
Al+(g) —> Al2+(g) + e– IE2 = 1,820 kJ/mol
Al2+(g) —> Al3+(g) + e– IE3 = 2,750 kJ/mol
Al3+(g) —> Al4+(g) + e– IE4 = 11,600 kJ/mol
a. Explain why IE for Al increases from IE1 to IE2
to IE3 to IE4.
b. Why is there a drastic increase in energy from
IE3 to IE4 for Al?

42. • Ionization energies decrease as atomic radii
increases. The farther an electron is from the
nucleus, the easier it is to be released. Down a
group, as n increases and atomic size increases,
electrons are easily released.
• Ionization energy decreases from top to bottom
of a group. Across a period, as Zeff increases and
size decreases, ionization energy increases.

43. IE1 Li Be B C N O F Ne
kJ/
mol
520 899 801 1,086 1,400 1,314 1,680 2,080
The following tables provide some first
ionization energies for representative
elements.
IE1 Li Na K Rb Cs
kJ/mol 520 495.9 418.7 403.0 375.7

44. 5. Electron Affinity
The electron affinity of an atom may be defined
as the negative of the energy change that occurs
when a gaseous atom accepts an electron.
F(g) + e–—> F– (g) Energy involved = -328 kJ/mol
The electron affinity is
F– (g) —> F(g) + e– EA = +328 kJ/mol

45. • Electron affinity is sometimes defined as the
ionization energy of a negative ion.
• The more positive the electron affinity, the
greater the tendency to accept an electron and
form an ion.
• The electron affinity increases across a period
from left to right. The electron affinity generally
decreases going down a group.

46. Knowledge of atomic and ionic radii is used to vary
physical properties of materials.
a. Strengthening Glass. Normal glass windows
that contain Na+ and Ca2+ ions are brittle and
shatters easily. Replacing the Na+ ions with
bigger K + ions results in surfaces where
surface sites are being filled up leaving less
opportunity for cracking.
Some Applications of Metal Ions

47. b. Colors in gemstones. Pure Al2O3 is colorless.
Substituting Al3+ with a little amount of Cr3+
in Al2O3 gives a red color in ruby.

48. Fill up the blank periodic table with the Element as
described by each statement below:
1. Element A is the biggest in Group 1A.
2. Element B forms the biggest anion in period 2
3. Element C has complete d electrons in period 4
4. Element D is the most electronegative in
period 2.
Seatwork

49. 5. Element E will be isoelectronic with the noble
gas in period 3 when it loses two electrons.
6. Element F has the highest ionization energy in
period 4.
7. Element G has the least electron affinity in
group 6.
8. Element H has the 4f14 configuration
9. Element I is the first member of the actinide
series.

51. THANK YOU FOR
LISTENING!