CHAPTER 3.0

CHEMISTRY UNIT, KMNS

CHAPTER 3
PERIODIC TABLE

3.1
Classification Of

Elements

3

Introduction
The periodic table (PT), is a tabular display of the chemical elements, which are
arranged by atomic number, electron configuration, and recurring chemical properties.
The structure of the table shows periodic trends. The seven rows of the table, called
periods, generally have metals on the left and nonmetals on the right. The columns,
called groups, contain elements with similar chemical behaviors.

Period

 Seven horizontal rows of elements, which are numbered 1 to 7.
 The elements are arranged in order of the increasing proton number from left to

right of PT.
 The successive periods from top to bottom of the PT contains 2,8,8,18,18,32 and 32

elements.
 The valence shells of elements in the same period have the same principal

quantum number.

4

Group

 Vertical columns of elements in the periodic table.
 The groups in the Periodic Table are numbered from 1 to 18.
 Elements placed in the same group, have

i. similar chemical properties
ii. same number of valence electrons

Main Groups in the Periodic Table

Group 1 Alkali metals (except H) 5
Group 2 Alkaline earth metals
Group 3 - Group 12 Transition metals
Group 16 Chalcogens
Group 17 Halogens
Group 18 Inert/noble gases

Blocks

 A block is a set of elements unified by the orbitals their valence electrons or vacancies
lie in. Each block is named after its characteristic orbital:
s, p, d, and f-blocks.

s-block p-block

d-block
f-block

6

f block d block p block s block Group 1 & 2
Electronic configuration of valence electron: ns1 to ns2
Eg. 12Mg = 1s2 2s2 2p6 3s2
consists of alkaline metals and alkaline earth metal.

Elements of Group 13 – 18
Electronic configuration of valence electron: ns2 np1 to ns2 np6
Eg. 15P: 1s2 2s2 2p6 3s2 3p3
consists of metals, metalloids & non-metals elements.

Elements in Group 3 – 12
Consist of the transition elements. Electronic configuration of
valence electron: (n-1)d1 ns2 to (n-1)d10 ns2
Eg. 23V: 1s2 2s2 2p6 3s2 3p6 3d3 4s2

Elements in Series of lanthanides (4f) & actinides (5f).
Final electron filled into a f subshell. f-Block elements mostly radioactive

7

Metals

 Properties: Dense, malleable, ductile, shiny, good conductor of heat and electricity.

8

Metalloids

 Known as semi-metals and have properties between those of metals and nonmetals
tend to be semiconductors and widely used in manufacturing sector

9

Nonmetals

 Further across the period towards the right, elements gradually lose their metallic
character and gained nonmetallic features. Generally gases which poor conductor.

nonmetals

10

Determining Position of Elements Based On Electronic Configuration

Example: Phosphorus 15P : 1s2 2s2 2p6 3s2 3p3

Period: Highest principal quantum number, n :3

Block: Subshell occupied by electron of highest energy : p

No of valence electrons (v.e.) :2+3=5

Group: p block (10 + number of v.e.) : 10 + 5 =15

Period 3,
p block
Group15

11

 For main group elements (s block and p block), valence electrons are those in the
electronic shell of highest principal quantum number, n.

Example 1: Example 2:
3Li 1s2 2s1 4Be 1s2 2s2

11Na 1s2 2s2 2p6 3s1 12Mg 1s2 2s2 2p6 3s2

19K 1s2 2s2 2p6 3s2 3p6 4s1 19K 1s2 2s2 2p6 3s2 3p6 4s2

 Highest principal quantum no: 3.  Highest principal quantum no: 2.
 Occupied by only 1e.  Occupied by 2e.
 Hence, no. of valence ē = 1.  Hence, no. of valence ē = 2.
 Valence electronic configuration, 3s1.  Valence electronic configuration, 2s2.

12

Example 3: • Highest principal quantum no: 2.
10Ne 1s2 2s2 2p6 • Occupied by 8e.
• Hence, no. of valence ē = 8.
• Valence electronic configuration 2s2 2p6.

 For d-block elements, valence electrons are those in the outermost shell, n and
(n-1)d orbital in inner shell.

Example 4:

27Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7 • No. of valence ē = 9.
• Valence electronic configuration:

4s2 3d7

13

3.2
Periodicity

PERIODICITY

 Periodicity is the periodic trend in properties of elements.

Periodic Trends In The Size Of Atom (Atomic Radii)

 The size /radius of atom is difficult to be defined exactly because the electron cloud
has no clear boundary.

 To solve this, distance between the 2 nuclei in a molecule is determined.
 Radius, r = half of the distance between the nuclei of two adjacent identical atoms.

Radius, r = a/2

a 15

Periodic Trends In The Size Of Atom (Atomic Radii)

Factors that determine atomic radii are:
1. Effective nuclear charge (Zeff) felt by the valence electrons.
2. The shielding effect. Value of the n of the valence electrons.
3. Mutual repulsion of electrons in valence shell important only if factor 1and 2 is the

same.

1. Effective nuclear charge (Zeff) Zeff = Z – S
 The residual nett charge felt by
Z = no. of protons
the valence electrons. Those S = shielding factor
electrons closer to the nucleus (≈ no. of inner shell electrons)
experience a greater force than
those that are further away. 2. The Shielding @ Screening Effect
 is the result of the mutual repulsion

between the electrons in the inner
shell with those in the outer shell.

16

Trends in atomic size Increasing atomic radius Decreasing atomic radius

1. Decreases across a 17
period from left to
right.

2. Increases going
down a group.

3. Relatively constant
across the first
series of transition
metals.

1. Across Period 3 ELEMENT Z Electronic S Zeff Atomic
Configuration Radius
 As the nuclear charge Na 11
increase, Left to right 1s2 2s2 2p6 3s1 10 +1 186

 The attraction between Mg 12 1s2 2s2 2p6 3s2 10 +2 160
the nucleus and the
valence e- become Al 13 1s2 2s2 2p6 3s2 3p1 10 +3 143
stronger (Zeff increases),
Si 14 1s2 2s2 2p6 3s2 3p2 10 +4 117
 As a result, valence
electrons are pulled P 15 1s2 2s2 2p6 3s2 3p3 10 +5 110
closer to the nucleus, the
sizes of atom decreases S 16 1s2 2s2 2p6 3s2 3p4 10 +6 104

Cl 17 1s2 2s2 2p6 3s2 3p5 10 +7 99

Ar 18 1s2 2s2 2p6 3s2 3p6 10 +8 94

18

Li 1s2 2s1 2. Down Group 1

Na 1s2 2s2 2p6 3s1  As the number of shell

increases,
 Shielding effect of inner

K 1s2 2s2 2p6 3s2 3p6 4s1 electrons increases.
 The attraction between

the nucleus and the

Rb 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s1 valence electrons

become weaker.

 As a result, valence

Cs 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s2 5p6 6s1 electron cloud expands
and size of atom

increases.

Fr 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s2 5p6 6s2 6p6 7s1

19

3. Across First Row Of Transition Elements

 The transition metals are the elements which have incompletely filled d subshells.
 Across the period of the first row of transition metal: Size of atoms do not change

significantly.

 Atomic Radii Of First Row of Transition Elements

Element Sc Ti V Cr Mn Fe Co Ni Cu Zn

Atomic radius (pm) 162 147 134 128 127 126 125 124 128 132

 Explanation :
 As proton number increases, the electron will be added to the unoccupied inner 3d

subshell.
 This will increase the shielding effect to the 4s electrons. Therefore, the increase in

nuclear charge is cancelled by the increase in shielding effect.
 Therefore, there is no significant change in atomic size of the transition metals.

20

Element Atomic Z Electronic configuration Zeff = Z - S
radius (pm)
25 1s2 2s2 2p6 3s2 3p6 3d5 4s2 25 - 18 = +7
Mn 127 26 1s2 2s2 2p6 3s2 3p6 3d6 4s2 26 - 18 = +8
27 1s2 2s2 2p6 3s2 3p6 3d7 4s2 27 - 18 = +9
Fe 126 28 1s2 2s2 2p6 3s2 3p6 3d8 4s2 28 - 18 = +10

Co 125

Ni 124

 The table above illustrates how the increase in nuclear charge is offset
by the increase in shielding effect. Atomic size of the first row of transition
metals are relatively constant.

21

Size Of Atoms And Their Corresponding Ions

In general,

 a cation is smaller than its atom.

 while an anion is larger than its
atom.

 When an atom loses an electron,
there will be less repulsion
between the remaining electrons,
so the electron cloud will occupy
less space and the size of ion will
be smaller than its atom (and vice
versa).

atomic radius (pm) 22

Size of Sodium Atom & Sodium ion Na+
Loses an electron

Na

( r = 186 pm) ( r = 95 pm)

Element Atomic radius (pm) Z Electronic configuration Zeff = Z - S
11 - 10 = +1
Na 186 11 1s2 2s2 2p6 3s1 11 - 2 = +9
Na+ 95 11 1s2 2s2 2p6

 When Na atom loss@ donate electron from the valence shell to form Na+ ion.
 Na+ ion has fewer number of electron shells.
 Shielding effect decreases.
 Electron-electron repulsion decreases.
 Attraction of the nucleus towards remaining electron increases.
 Hence Na+ ion is smaller than Na atom.

Size of Chlorine Atom & Chloride ion Cl–

Gain an electron

Cl

(r = 99 pm) (r = 181 pm)

Element Atomic radius (pm) Z Electronic configuration Zeff = Z - S
17-10 = +7
Cl 99 17 1s2 2s2 2p6 3s2 3p5 17-10 = +7
Cl- 181 17 1s2 2s2 2p6 3s2 3p6

 When Cl atom gain@ accept electron into the valence shell to form Cl- ion.
 Electron-electron repulsion increases@ repulsion between valence electrons

incrases.
 Attraction of the nucleus towards outermost electrons@ valence electrons are

weaker
 Hence Cl- ion is bigger than Cl atom.

Isoelectronic

Isoelectronic species
 species that have the same number of electrons and hence the same
electron configuration.

Example: Having the
Same
18Ar : 1s2 2s2 2p6 3s2 3p6 number of
17Cl– : 1s2 2s2 2p6 3s2 3p6 electrons
20Ca2+ : 1s2 2s2 2p6 3s2 3p6

25

Compare The Radius Of Isoelectronic Species

E Na, Mg and Al are elements located in period 3. The formation of cations from neutral

atom of representative element occur when one or more electrons are removed from
the outermost shell.

Example:

Neutral atoms Cations
Na: 1s2 2s2 2p6 3s1 Na+: 1s2 2s2 2p6 = [Ne]
Mg: 1s2 2s2 2p6 3s2
Al: 1s2 2s2 2p6 3s2 3p1 Mg2+: 1s2 2s2 2p6
Al3+: 1s2 2s2 2p6

 These cations have stable electronic configuration.

26

Compare The Radius Of Isoelectronic Species

E F, O and N are period 2 elements .The formation of anions from neutral atom of

representative element occur when one or more electrons are added to the outermost
shell.

Example:

Neutral atoms Anions
F: 1s2 2s2 2p5 F-: 1s2 2s2 2p6 = [Ne]
O: 1s2 2s2 2p4
N: 1s2 2s2 2p3 O2-: 1s2 2s2 2p6
N3-: 1s2 2s2 2p6

 These anions also have stable electronic configuration.

27

Compare The Radius Of Isoelectronic Species

ExampEle: Anions
F-: 1s2 2s2 2p6
Cations O2-: 1s2 2s2 2p6
Na+: 1s2 2s2 2p6 N3-: 1s2 2s2 2p6
Mg2+: 1s2 2s2 2p6
Al3+: 1s2 2s2 2p6

 All the cations and anions above are known as isoelectronic ions because they have
same electronic configuration. Figure 3.2.1 shows some of the series of
isoelectronic ions.

28

All ions
have 10
electrons

A series of
isoelectronic ions

Figure 3.2.1: Series of isoelectronic ions

29

Figure 3.2.2: Variation in Radius of Isoelectronic Species 30

Ion N3– O2– F– Na+ Mg2+ Al3+

No. of 10 10 10 10 10 10
electrons No. of electrons constant +12 +13
Nuclear +10 +11
+7 +8 +9 +11
charge Increasing nuclear charge 0.66 0.51
Radius DECREASES
Zeff +5 +6 +7 +9
1.71 Increasing Zeff
Radius (Å)
1.40 1.33 0.97

 For isoelectronic ions,
cations are always
SMALLER than anions.

 Figure 3.2.2 shows size comparison of isoelectronic ions.

 For isoelectronic ions, cations

are always smaller than anions.

 Both Na+ and F- ions Eg: Na+ is smaller than F-.

has 10 electrons but

Na (Z = 11) has more 02 01  Due to stronger nucleus
protons than attraction, the remaining
F (Z = 9). The larger electrons of Al3+ is pulled
effective nuclear
charge of Na+ results inward more than in
in smaller radius. 04 Mg2+. Therefore ionic

radii of Al3+ is smaller

 For isoelectronic cations, Al3+ has the 03 than Mg2+ . The smaller
same number of electrons as Mg2+ radius of Mg2+ compared
(10 electrons) but Al (Z = 13) has with Na+ can be similarly
explained

one more proton than Mg (Z = 12). 31

 Figure 3.2.2 shows size comparison of isoelectronic ions.

 For isoelectronic anions, ionic
radius decreases from N3-, O2-
and F-.

 This is due to proton 06 05
number increase from
N (Z = 7), S (Z = 8) and  Therefore, N3- has
F (Z = 9) . 08 biggest ionic radii

compared to O2- and F-.

 N3- has lowest effective nuclear 07
charge, therefore the attraction
between nucleus and valence
electrons of N3- is the weakest.

32

Example:
You are given Na+, Mg2+ , O2- and N3- ions. Arrange the following ions in descending
radius. Explain.

 Radius: N3- > O2- > Na+ > Mg2+

 N3- , O2- , Na+, Mg2+ are isoelectronic species with electronic configuration 1s2 2s2 2p6.

 From N to Mg, proton number increase with N(Z = 7), O (Z = 8), Na (Z = 11) and Mg
(Z = 12).

 As proton number increases, effective nuclear charge increase from N3- to Mg2+ .
N3- experience weakest nucleus attraction towards valence electron while Mg2+
experience the strongest nucleus attraction towards valence electron.
Therefore, ionic size decrease from N3- to Mg2+ .

33

Example:

15P3- 16S2- 17Cl-

No. of 15 16 17
proton

No. of 1s2 2s2 2p6 3s2 3p6
electron
18 18 18

Predict the size of the isoelectronic species. Explain.

 From P3- to Cl- ionic radii decrease.

Explanation:

 As proton number, Z increases from P3- to Cl-, effective nuclear charge, Zeff increases
which lead to a greater attraction between the nucleus and the valence electrons,

hence reduce the size. 34

Variation In The Ionic Radii Across Period 2 And Period 3

FVariation Of Ionic Radii Across Period 2

Proton Ion Electronic Principal
Number configuration quantum
Li+ number, n = 1
3 Be2+ 1s2
4 B3+ 1s2 Principal
5 C4+ 1s2 quantum
6 N3- 1s2 number, n = 2
7 O2- 1s2 2s2 2p6
8 1s2 2s2 2p6 35
9 F- 1s2 2s2 2p6

Example:
By referring to the previous slide, explain the variation of ionic radii across period 2.

Guidelines:
To solve this question, you have to divide your answer into three parts.
√ First : compare ionic radius among isoelectronic cations, from Li+ to C4+
√ Second: Compare ionic radius between cation C4+ and anion N3-
√ Third : compare ionic radius for all isoelectronic anions from N3- to F-

36

 Cations Li+, Be2+ , B3+ and C4+ are isoelectronic with electronic configuration 1s2.

 Ionic radii decreases from Li+ to C4+ . From Li (Z = 3) to C (Z = 6), proton number increase,
thus effective nuclear charge, Zeff increase within the same shell (n = 1). Attraction between
nucleus and valence electrons become stronger.

 There is a big jump in radius occurred from C4+ to N3- due to an increase in the number of
shell from n = 1 to n = 2 which means greater shielding effect. Attraction between nucleus and
valence electrons become weaker.

 Anions N3- to F- are isoelectronic with electronic configuration 1s2 2s2 2p6.

 From N3- to F-, ionic radii decreases due to increase in proton number from N (Z = 7) to
F (Z = 9). Effective nuclear charge, Zeff increase within the same shell (n = 2). Attraction
between nucleus and valence electrons become stronger.

 All anions (N3-, O2-, F-) have bigger ionic radius than cations (Li+, Be2+, B3+ and C4+) because
anions has one shell more than the cations.

37

Variation Of Ionic Radii Across Period 3

Proton Number Ion Electronic Principal
configuration quantum
11 Na+ number, n = 2
12 Mg2+ 1s2 2s2 2p6
13 Al3+ 1s2 2s2 2p6 Principal
14 Si4+ 1s2 2s2 2p6 quantum
15 P3- 1s2 2s2 2p6 number, n = 3
16 S2- 1s2 2s2 2p6 3s2 3p6
17 Cl- 1s2 2s2 2p6 3s2 3p6 38
1s2 2s2 2p6 3s2 3p6

39

Graph Of Ionic Radii vs Proton Number

The graph shows The size of cations From Si 4+ to P3- ionic radii
the decrease (from Na+ to Si4+) decreases of P3- increase due to an
in ionic radii for due to an increase in increase in the principal
isoelectronic effective nuclear charge, Zeff. quantum numbers from
cations Valence electrons are held n = 2 to n = 3 which lead to
(Na+ to Si 4+) and more tightly to the nucleus an increase in screening
isoelectronic anions and hence ionic radii effect/ shielding effect.
(P3- to Cl-). decreases. Attraction between nucleus
and valence electrons
becomes weaker.

The size of anions
(from P3- to Cl-) decreases due to
an increase in effective nuclear
charge, Zeff. Valence electrons
are held more tightly to the
nucleus and hence ionic radii
decreases.

Review Questions

1. Determine the period, group and block for the following elements:
X: 1s2 2s2 2p2
Y: 1s2 2s2 2p6 3s2 3p6 4s1
Z: [Ar] 3d5 4s2

2. Rank each set of ions in order of decreasing radii and explain.
(a) Ca2+, Sr2+, Mg2+
(b) K+, S2–, Cl–

40

Question 1:

X: 1s2 2s2 2p2  period 2
 group 14
Valence shell, n = 2  p-block
Valence electrons = 4
Highest energy orbital = p

Y: 1s2 2s2 2p6 3s2 3p6 4s1

Valence shell, n = 4  period 4

Valence electron = 1  group 1

Highest energy orbital = s  s-block

41

Question 1:

Z: 1s2 2s2 2p6 3s2 3p6 4s2 3d5

Valence shell, n = 4  period 4
Valence electrons = 7  group 7
Highest energy orbital = d  d-block

42

Question 2: 43

(a) Radii: Sr2+ > Ca2+ > Mg2+ (b) Radii: S2- > Cl- > K+

From Sr2+, Ca2+ to Mg2+ S2-, Cl- and K+ are isoelectronic
species.
 Number of shells filled by From S2–, Cl– to K+
 Proton number increases
electrons decreases  Effective nuclear charge
 Shielding effect decreases
 Attraction between nucleus and increases
 Attraction between nucleus and
valence electrons increases
 Radius decreases valence electrons increases
 Radius decreases

First And Second Ionisation Energies 44

GG

Second Ionisation Energy, IE2
The energy required to remove one mole of electron from one mole of gaseous
unipositive ion.

X+ (g)  X2+ (g) + e– ∆H = IE2

45

FAFCaTcOtoRrSs IINnFflLuUenEcNinCgINIoGnisation Energy, IE Larger atomic radius,
IONISATION ENERGY  Lower IE.

ATOMIC
SIZE

EFFECTIVE IONISATION SHIELDING
NUCLEAR ENERGY EFFECT
CHARGE (IE)
Higher principal quantum
 Higher Zeff number (or number of shells)
 Higher IE  Lower IE.

DOWN THE GROUP ACROSS A PERIOD

Atomic size@radii Shielding Effect Effective Nuclear
 When atomic  When shielding Charge, Zeff
 When Zeff
radii increases, effect increases
 Attraction  Attraction increase
 attraction between
between nucleus between nucleus
nucleus and
and valence and valence valence electrons
become stronger
electrons electrons become  More energy is
required to
become weaker weaker remove the 1st
 Less energy is  Less energy is valence electron
electron.
required to required to  Higher IE1
remove the 1st remove the 1st

valence electron valence electron
 Lower IE1  Lower IE1

46

Example:

With reference to the table given below

Element 11Na 12Mg
496 738
Ionisation energy
(kJ mol-1)

(a) Define first ionisation energy

 The minimum energy required to remove one mole of electron from one
mole of gaseous atom to produce unipositive ion.

47

Example: 48

With reference to the table given below

Element 11Na 12Mg
496 738
Ionisation energy
(kJ mol-1)

(b) Why first ionization energy of sodium smaller than magnesium?

 Na : 1s2 2s2 2p6 3s1 Mg : 1s2 2s2 2p6 3s2

 Proton number for Na is less than Mg.

 Effective nuclear charge, Zeff for Na is lower than Mg.
 Attraction between nucleus and valence electrons in Na is weaker than Mg.

 Atomic size of Na is bigger than Mg.

 Less energy is required to remove the first electron in Na than Mg.

 Na has lower first ionisation energy, IE1 than Mg.

Example: 49

With reference to the table given below

Element 11Na 12Mg
496 738
Ionisation energy
(kJ mol-1)

(c) Which element has second ionisation energy?

Na+ : 1s2 2s2 2p6 Mg+ : 1s2 2s2 2p6 3s1

Na2+ : 1s2 2s2 2p4 Mg2+ : 1s2 2s2 2p6

 Na has higher second ionisation energy because the second electron is
remove from completely filled 2p orbital which is more stable and thus
require higher energy.

 In Mg, second electron is removed from partially filled 3s orbital which is
less stable.

50

Variations In The First Ionisation Energy Across A Period And Down A Group

Group 1 Group 2

Decreasing IE1

Increasing atomic radius
Increasing shielding effect