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TOPIC 4:
The Periodic Table and Some
Atomic Properties
Contents
1. Classifying the Elements:The Periodic Law and the
Periodic Table
2. Description of a modern Periodic Table-The Long
Form
3. Electron Configurations and The Periodic Table
4. Metals and Nonmetals and Their Ions
5. The Sizes of Atoms and Ions
6. Ionization Energy
7. Electron affinity
8. Magnetic Properties
9. Periodic Properties of The Elements
2
Alkali metals
Periodic Table
Alkaline earth metals
Halogenes
Noble Gases
Main Group
Transition Elements
Main Group
Lanthanides and actinides
3
Main Group Elements (IA – VIIIA)
1s
1s
2s
3s
2p
Transition Elements (IB – VIIIB)
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
Inner Transition Elements
4f
5f
4
METALS, NONMETALS AND THEİR IONS
METALS –In the periodic table there are 110 elements, The
total number of metal elements present is 85 of which only
15 elements are defined as artificial.
-Metals are good conductors of heat and electricity. They
easily lose e- and form + ions.
-They possess metalic bonds and are lustrous(shiny). They
can form alloys among each other.
-They are ductile. Most metals can be drawn out into thin
wires.
-They are present always at solid phase except for Hg, Ga,
Cs ve Fr .None of them is volatile except for Hg. They have
high melting and boling points
-Active metals (s1) react easily with water , half-active metals
(s2) react with water rather difficult, inactive (noble) metals
do not react with water. The active metals form the bases.
5
METALS, NONMETALS AND THEİR IONS
• NONMETALS – Except for H , all of them are the elements of
the p orbital element are totally in number 16.
• -They do not conduct heat and electricity. They gain or share
e- easily. They are not shiny.
While F-, O= can only (–) charged ions(Anions) , the others
can have –/+ ions (Anion/Cation)
• -They build up covalent bonds among each other, and form
ionic bonds with metals and metalloids
• -They have low melting and boiling points, most of them are
present at gas state at normal conditions. Br2 is liquid ; P, S, I2,
As are solid.
• -They can not combine with the acids without O. Their
Hydrides (Compunds with H) are acids .Nonmetaloxides are
acid anhydrides (They react with water and form acids).
6
METALLOIDS
• METALLOIDS are elements that look like
metals and in some way behave like metals,
but that have also have some nonmetallic
properties. They are present at solid form and
have high boiling points.
7
Metals lose e- to achieve noble gas electron configuration
Nonmetals gain e- to have noble gas electron configuration.
8
Nonmetal Ions:
The atoms of Groups 7A and 6A- the most active nonmetals,
have one and two electrons fewer than the noble gas at the end
of the period. These atoms can acquire the electron
configurations of noble gas atoms by gaining the appropriate
number of electrons.
Cl ([Ne]3s23p5) + e-  Cl- ([Ar])
S ([Ne]3s23p4) + 2e-  S-2 ([Ar])
In most cases a nonmetal atom will gain a single electron
spontaneously, but energy is required to force it to accept
additional electrons. Often other processes occuring
simultaneously supply the necessary energy (such as an
attraction to positive ions).
9
• Transition Metal Ions:
• Only a few transition metal atoms acquire noble-gas
electron configuration by losing electrons (such as in the
loss of 3 electrons of Scandium, Sc  Sc+3). Most
transition metal atoms, however, do not acquire a noblegas configuration when they ionize. Furthermore, a
common feature of transition metal is the ability to form
more than a single type of ion (Example for this
situation:Iron).
• Fe ([Ar]3d64s2)  Fe+2 ([Ar]3d6) + 2e• Fe ([Ar]3d64s2)  Fe+3 ([Ar]3d5) + 3e10
THE SIZES OF ATOMS AND IONS
To understand the physical and chemical properties,
we need to know something about atomic sizes.
Atomic Radius: The probability of finding an
electron decreases with increasing distance from the
nucleus, but nowhere does the probability fall to zero.
For this reason, there is no precise outer boundary to
an atom and it is hard to determine the atomic
radius. It is the distance between the nuclei of atoms
which can be measured effectively. We define atomic
radius in term of internuclear distance (covalent
radius, ionic radius, metallic radius).
11
The covalent radius of Br is 1,14 Å.
The length of C – C covalent bond is 1,54 Å .
Hence, it can be claimed that C has a radius of
0,77 Å .
12
Variation of Atomic Radii Within a Group of the
Periodic Table
The more electronic shells in an atom, the larger the atom
is. Atomic radius increases from top to bottom through a
group of elements.
The atomic radius decreases from left to right through a
period of elements, whereas atomic radii do not change very
much within a transition series.
13
14
Ionic Radius
When a metal atom loses one or more electrons to form a
positive ion, there is an excess of nuclear charge over the
number of electrons in the resulting cation. The nucleus
draws the electrons in closer, and as a consequence
• Cations are smaller than the atoms from which they are
formed
• For isoelectronic(having equal number of electrons)
cations, the more positive the ionic charge is , the
smaller is the ionic radius.
• When a nonmetal atom gains one or more electrons to
form a negative ion(anion),the nuclear charge remains
constant , but Zeff is reduced because of the additional
electrons.
15
Repulsions among the electrons increase and they are
not held tightly.
Anions are larger than the atoms from which they
are formed.
For isoelectronic anions, the more negative the ionic
charge is, the larger is the ionic radius.
Practice: Refer only to the periodic table and arrange
the following species in order of increasing size:
Ar, K+, Ca+2, Cl-, S-2
16
IONIZATION ENERGY
Atoms do not eject electrons spontaneously. Electrons are
attracted to the positive charge on the nucleus of an atom,
and energy is needed to overcome that attraction.
• The more easily an atom loses its electrons, the more it
tends to have a metallic character.
• Ionization Energy (I) is the quantity of energy a
gaseous atom must absorb so that an electron is stripped
from the atom. The electron is the one most loosely held.
• First ionization energy (I1), is the energy required to strip
one electron from a neutral gaseous atom. The second
ionization energy, I2 is the energy to strip an electron
from a gaseous ion with a charge of+1 . Further
ionization energies are (I3, I4, …)
•
Na(g)  Na+(g) + 1e17
Ionization Energy (kJ/mol)
Element
I1
I2
I3
Na
496
4560
Mg
738
1450
7730
Al
577
1816
2744
I4
11.600
Ionization energies decrease as atomic radii
increase.
The required energy to ionize the remaining electrons from
the ion of an atom (after an electron from the neutral atom
was stripped) is higher than the first ionization energy.
The smaller energy levels , the higher ionization energy
required to strip an electron from that level.
18
Example: Arrange the following in order of
increasing I1 (As, Sn, Br, Sr)?
19
Electron Affinity:
Ionization energy concerns the loss of electrons .
Electron affinity (EA), is a measure of the energy change that
occurs when a gaseous atom gains an electron .
When most of the positively charged ions or neutral atoms
gain an electron, heat is given off to the environment
(exothermic reaction). The EA is indicated with a – sign in this
case.
[Cl(g) + e-  Cl-(g) EA = - 328 kJ/mol ]. The EA of Chlorine is 328 kJ/mol.
Li(g) + e-  Li-(g) EA = - 59,8 kJ/mol (1s22s1  1s22s2 –
example for the metal atom)
20
Electron Affinity
•
In order that the anions and some neutral atoms gain
an electron, the energy of heat must be absorbed from
the environment. This event is an endothermic reaction
and has a positive value of EA .
• Ne(g) + e-  Ne-(g) EA = + 29 kJ/mol
1s22s22p63s1)
(1s22s22p6 
• O(g) + e-  O-(g) EA = - 141,4 kJ/mol
• O-(g) + e-  O-2(g) EA = + 880 kJ/mol
21
The smaller atoms to the right of the periodic
table,namely the halogenes(Group VII A) have large
negative electron affinities.
For example , in becoming Cl- , a chlorine atom
acquires the very stable electron configuration of the
noble gas Argon(Ar)
For the IIA ve VIIIA group elements, the electron
gained must enter the higher energy orbital since all
the other s and p orbitals of the atoms are filled. This
situation requires an endothermic process.
22
• Magnetic Properties
• The behaviour of atoms and ions in a magnetic field is also
helpful in establishing electron configurations. A spinning
electron is an electric charge in motion. It induces magnetic
field.
• In a diamagnetic atom or ion all electrons are paired, and
these individual magnetic effects cancel out. A diamagnetic
species is weakly repelled by a magnetic field.
• A paramagnetic atom or ion has unpaired electrons, and the
individual magnetic effects do not cancel out. The unpaired
electrons induce a magnetic field that causes the atom or ion
to be attracted into an external magnetic field. The more
unpaired electrons present, the stronger the attraction.
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Periodic Properties of the elements:
Physical properties (melting points, boiling points, etc. )
Chemical properties (reduction,oxidation)
Metal Elements
Luster
Malleable
Good conductors of heat
and electricity
Nonmetal Elements
Dull appearance
Brittle, high degree of
hardness
Poor conductors of heat and
electricity
Their ionic compounds with
Their oxides have acid
the oxides have base
character
character
They are the anions or oxiThey are the cations in the
anions in the aqueous
aqueous solutions
solutions.
24
25
Group 1A (Alkali Metals):
1A
3
Li
(Lithium)
11
Na
(Sodium)
19
K
(Potassium)
37
Rb
(Rubidium)
55
Cs
- Each of them has 1 electron in the s orbital.
- FROM TOP TO DOWN, the melting point
decreases.
- The density rises.
- Atomic radius rises
- First ionization energy lowers (I1)
- They are the group having the smallest I1
among the other groups of elements in the
periodic table.
- They are extremely reactant since their ability to
gain or lose electron and to build up ions is high.
Therefore, they are present in nature in the form
of compounds. (M  M+ + e- )
(Cesium)
87
Fr
(Francium)
26
• 2M(s) + H2(g)  2MH(s) They react with hydrogen to form
hydrides.
• They react with sulphur to form sulphides:
• 2M(s) + S(s)  M2S(s)
• They react with chlorine to form chlorides:
• 2M(s) + Cl2(g)  2MCl(s)
• When they react with water, hydrogen gas and the
hydroxide of the alkali metal are produced (exothermic
reaction):
• 2M(k) + 2H2O(s)  2MOH(aq) + H2(g)
• Their reactions with oxygen is more complex:
• 4Li(s) + O2(g)  2Li2O(s) (lithium oxide)
• 2Na(s) + O2(g)  Na2O2(s) (sodium peroxide)
• K, Rb ve Cs may form superoxides
• K(s) + O2(g)  KO2(l) (potassium superoxide)
27
Group 2 A: Alkaline Earth Metals
2A
4
Be
- Compared to the Group 1A metals
- they have greater hardness
(Beryllium)
12
Mg
(Magnesium)
20
Ca
(Calcium)
38
Sr
(Strontium)
56
Ba
(Barium)
88
Ra
- higher densities
- higher melting points
- higher I1 values and that’s why they are less
reactant.
- Be and Mg are the least active earth metals.
- Calcium and the other elements placed below it
react with water at room temperature :
(Radium)
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
28
• GROUP 3A METALS
13
Al
(Aluminium)
31
Ga
(Gallium)
49
In
(Indium)
81
Tl
Aluminium is an excellent reducing
agent since it is easily oxidized to
+3 ions. It’s main use is lightweight alloys
Aluminium is such a good reducing
agent that it will extract oxygen
from metal oxides to produce
Aluminiumoxide
Fe 2O 3 (s) + 2 Al (s)
Al2 O 3 (s) + 2 Fe (l)
(Thallium)
29
It reacts with acids to produce H 2 (g):
2 Al (s) + 6 H + (aq)
2 Al 3+ (aq) + 3 H 2 (g)
Powered Aluminium is easily oxidized by air or other oxidants
in highly exothermic reactions used in rocket fuels and
explosives
Galium is used to make GaAs that can convert light into
electricity(photoconduction). This semi-conducting material
is also used in light emitting diodes(LED) and in solid state
devices such as transistors
Indium is a soft silvery metal used to make low-melting point
alloys.
Thallium and its compounds are extremely toxic. It’s use is in
high temperature superconductors
30
GROUP 6A: THE OXYGEN FAMILY
6A
- From top to bottom the elements tend to show
more metalic character
8
O
(Oxygen)
16
S
- Oxygen is in gaseous form , the other elements
are at solid state.
(Sulphur)
34
Se
(Selenium)
52
Te
(Tellurium)
84
Po
- Nonmetals: Oxygen, sulphur and selenium;
- Tellurium is metalloide , polonium is metal.
- Oxygen can be found in two different molecular
forms:
O2 and O3 (ozon) (allotropic form).
3O2(g)  2O3(g) (+ 284,6 kJ)
(Polonium)
31
• Oxygen tends to strip electrons from the other elements,
and therefore it is a good oxidizing agent.
• - Oxygen generally is in the form of O-2
compounds with metals
when it forms
• - The other two forms of oxygen anions are O22- and O2• - Sulphur has plenty of allotropic forms of which the most
stable one is found as S8 in yellow color and at solid phase.
• - It makes up sulphides by gaining electron from other
elements (S-2).16Na(s) + S8(s)  8Na2S(s)
• - It is naturally abundant in the form of metal-sulphur
compound.
32
GROUP 7A: HALOGENES
7A
9
F
(Fluorine)
17
Cl
(Chlorine)
35
Br
(Bromine)
53
I
(Iodine)
85
At
(Astatine)
- Astatin is a radioactive element is
rarely found, so it’s properties are not
known.
- All halogenes are nonmetal.
- Under normal conditions each element
is composed of di-atomic molecules
- Fluorine: Pale yellow,
Yellowish green, bromine:
brown, iodine: Violet
chlorine:
Reddish
- Halogenes are the group with greatest
electron affinity, When they strip an
electron from other elements, this
reaction is exothermic.
33
 Fluorine and Chlorine are the most reactive ones, fluorine
can strip electrons from nearly all of the other elements.
• - Chlorine react with water and hydrochloric acid and
hypochlorous acid are produced :
• Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq)
• - Hypochlorous acid is an important desinfectant and used in
swimming pools(Chlorine is added into the water).
• - Halogenes react with plenty of metals to form metallic
halogenes
• Cl2(g) + 2Na(s)  2NaCl(s)
34
Grup 8A: THE NOBLE GASES (INERT GASES)
8A
2
He
(Helium)
10
Ne
(Neon)
18
Ar
(Argon)
36
Kr
(Krypton)
54
Xe
(Xenon)
86
Rn
(Radon)
- Nonmetal.
- At gaseous state at room temperature.
- Composed of monoatomic molecules
- Their “s” and “p” orbitals are totally filled in
- I1 is extremely high, falls down in the group from
top to bottom.
- Rn is pretty much reactive, some of it’s properties
are not known
- Very low reactivity rate
- In order to react with, they need the elements with
the highest electron affinity
Their compounds which we know until now:
XeF2 ,XeF4 ,XeF6 ,KrF2
He, Ne, ve Ar do not have any compounds and are
35
totally inert.
THE TRANSITION ELEMENTS
• d-block and f-block elements of the periodic table
• High melting points,good electrical conductivity,
moderate to extreme hardness
• Little variation among the atomic radii across the
first transition series(Exceptions:Sc and Ti).In the
series of elements in which the 4f subshell is
filled,atomic sizes actually decreases somewhat. This
phenomenon is called lanthanide
contraction(lanthanide series)
• Both ionic and covalent character( ionic character in
the compounds with the transition metal in lower
oxidation states; covalent character in higher
oxidation states
36
The Transition Elements
37
THE TRANSITION ELEMENTS
• Higher oxidation state in transition metals are
stabilized when oxide and/or fluoride ions are bound
to the metal
• Increasing stability of higher oxidation states for the
elements further down a group.
• Ionization energies are fairly constant across the first
transition series. Standard electrode potentials
increase in value across the series.All these
elements are more readily oxidized than
hydrogen(Exception: Cu)
• High catalytic activitity(ability to adsorb gaseous
species, some important catalysts:
Ni,Fe,Pt,Rh,V2O5,Cr2O3,MnO2,TiCl4
38
Different aqueous solutions of the transition
elements in various colors
39
THE TRANSITION ELEMENTS
• Most transition elements are paramagnetic.
• Fe, Co and Ni are ferromagnetic which means that
the magnetic moments are aligned into domains.The
ordering of domain persists when the object is
removed from the magnetic field and thus
permanent magnetism results
• Group 2B(Zn, Cd and Hg) is sometimes classified as
post-transition metals since these elements have the
electron configuration with no incomplete d-shell
(d10s2). They have lower melting points because their
full d-subshells prevent d-d bonding.
40
41
THE PERIODIC TABLE
42