Download Unit 7 Periodic Properties of the Elements in the Periodic Table

Unit 7-1
Unit 7 Periodic Properties of the Elements in the Periodic Table
Section 7.1
Periodic Variation in Physical Properties of the
Elements H to Ar
( 1 ) Periodic properties, structure and bonding of the elements H to Ar
Chemical periodicity refers to the regular occurrence of a set of properties of an element or its
compound in the Periodic Table. Because the Periodic Table lists the elements in order of atomic number,
these periodic properties are related to the electronic configuration of the element.
On moving to the right across the Periodic Table, the number of electrons in the outer shell increases
from one to eight. Therefore, it is to be expected that any physical properties connected with electronic
arrangement will also exhibit this periodicity. Such properties include atomic radii, ionization enthalpies
and electronegativities. Several of the physical properties of the elements, such as melting point and
density, depend upon their structure and bonding. However, the structure and bonding of the elements are
in turn related to atomic properties such as electron structures, ionization enthalpies and atomic radii.
The electronic configuration, structure and bonding of the elements in period 2 and period 3 are
shown in the following tables :
Period 2
Electronic
configuration
Outer-shell
configuration
Type of
element
Structure
Bonding
Period 3
Electronic
configuration
Outer-shell
configuration
Type of
element
Structure
Bonding
Li
2,1
Be
2,2
B
2,3
C
2,4
N
2,5
O
2,6
F
2,7
Ne
2,8
Metal
Metal
Metalloid
Metalloid
Non-m
etal
Nonmetal
Non-m
etal
Non-me
tal
Metallic : strong
forces of attraction
of positive ions for
mobile outer
electrons
Covalent : very strong Covalent within the molecule,
forces of attraction
van der Waals’ forces between
between atoms due to molecules
the attraction of nuclei
for shared electrons
Na
2,8,1
Mg
2,8,2
Al
2,8,3
Si
2,8,4
P
2,8,5
S
2,8,6
Metal
Metal
Metal
Metalloid
Non-me
tal
Non-m
etal
Covalent
Covalent within the molecule,
van der Waals’ forces between
molecules
Metallic
Cl
2,8,7
Ar
2,8,8
Non-m Non-met
etal
al
Unit 7-2
Structure and bonding in elements of period 2 and period 3 :
Across a period, the elements change from metals through metalloids to non-metals. In period 3,
sodium, the left-hand element, is a very reactive metal, whereas chlorine, next to the extreme right, is a very
reactive non-metal. These periodic changes in properties of the elements across the table are reflected in a
periodic change in structure. The structure of the elements varies from giant metallic, through giant
molecular in the metalloids to simple molecular structures in the non-metals.
Unit 7-3
( 2 ) Variation in atomic radii, first ionization enthalpies, electronegativities and
melting points
1. Atomic radii
Period 2
Atomic radius (nm)
Ionic radius (nm)
Li
0.123
0.060
Be
0.089
0.031
B
0.080
0.020
C
0.077
N
0.075
0.171
O
0.073
0.140
F
0.072
0.136
Ne
0.065
Period 3
Atomic radius (nm)
Ionic radius (nm)
Na
0.157
0.095
Mg
0.136
0.065
Al
0.125
0.050
Si
0.117
P
0.110
0.212
S
0.104
0.184
Cl
0.099
0.181
Ar
0.095
Notice that along each period there is a gradual decrease in atomic size as the outer electron shell
is being filled.
Moving from one element to the next across a period, electrons are being added to the same shell at
about the same distance from the nucleus. At the same time, protons are being added to the nucleus.
Therefore, the electrons are attracted and pulled towards the nucleus with an increasing positive charge. So
the radius of the atom decreases. However, the rate of decrease in the radius becomes smaller as the atoms
get heavier. The addition of one more proton to the 11 already present in sodium causes a greater
proportional change in nuclear attractive power than the addition of one more proton to the 16 already
present in sulphur. Thus, the atomic radius falls by 0.021 nm from Na to Mg, but only 0.005 nm from S to
Cl.
For metallic elements, the ionic radii are smaller than the corresponding atomic radii. The reason is
that removal of an electron leads to contraction of the electron cloud which is then pulled closer to the
positive nucleus.
For the non-metallic elements, the ionic radii are greater than the corresponding atomic radii. This
is because addition of an electron causes a greater repulsion by the electron cloud leading to a greater ionic
radius.
In a series of ions with the same number of electrons (an isoelectronic series), the ionic radius
decreases as the atomic number increases. The nuclear charge increases along the isoelectronic series. This
makes the electron clouds contract, because it is pulled in more effectively by an increasing positive charge.
The following figure shows the ionic radii of six isoelectronic ions all have ten electrons :
0.18
Ionic
radius (nm)
0.16
0.14
0.12
0.1
0.08
0.06
0.04
0.02
0
N(7)
O(8)
F(9)
Na(11) Mg(12) Al(13)
Atomic Number
Notice that hydride ion ( H- ) has an extraordinary large ionic radius : 0.208 nm. This is because
there is only one proton in a hydride ion to attract two electrons in the full-filled 1s orbital. The great
repulsion of the electron cloud results in a large ionic radius.
Unit 7-4
2. First ionization enthalpies
Period 2
First ionization
enthalpy (kJ mol-1)
Li
520
Be
900
B
800
C
1090
N
1400
O
1310
F
1680
Ne
2080
Period 3
First ionization
enthalpy (kJ mol-1)
Na
500
Mg
740
Al
580
Si
790
P
1010
S
1000
Cl
1260
Ar
1520
First Ionization
Enthalpy (kJmol-1)
First Ionization
Enthalpy (kJmol-1)
Period 2
Period 3
2500
2000
1500
1000
500
0
Li
Be
B
C
N
O
F
Ne
1600
1400
1200
1000
800
600
400
200
0
Na
Mg
Al
Si
P
S
Cl
Ar
Across the period, the first ionization enthalpy generally increases because the effective nuclear
charge increases giving rise to a decrease in atomic radius and thus a decrease in the distance of the electron
cloud. As the electron cloud becomes closer to the nucleus, it is more difficult to remove an electron from
an atom.
There are two discrepancies in both period 2 and period 3.
For example, in period 3 :
a. The first ionization enthalpy of aluminium is lower than that of magnesium. The electronic
configuration of aluminium is 1s22s22p63s23p1 but that of magnesium is 1s22s22p63s2 . Magnesium
atom has a full-filled subshell which acquires extra-stability. It is easier to remove the 3p electron of
aluminium from a higher energy and partially filled subshell (3p) than an electron from a full-filled
subshell (3s) of magnesium.
b. The first ionization energy of sulphur is less than that of Phosphorus. The electronic configuration
of sulphur is 1s22s22p63s23p4 but that of phosphorus is 1s22s22p63s23p3. Phosphorus atom has a
half-filled p subshell. This half-filled p subshell is more stable arrangement than the configuration
in sulphur. It is more difficult to remove an electron from phosphorus than from sulphur.
Unit 7-5
3. Electronegativities
H
2.1
Li
1.0
Na
0.9
Be
1.5
Mg
1.2
B
2.0
Al
1.5
C
2.5
Si
1.8
N
3.0
P
2.1
O
3.5
S
2.5
F
4.0
Cl
3.0
He
Ne
Ar
-
Pauling defined the electronegativity of an atom as the power of that atom in a molecule to attract
electrons. He obtained values of electronegativity by considering the strengths of the bonds between atoms
in molecules. Non-metals with a strong desire to gain electrons have the highest values of electronegativity.
Metals have low values.
Notice that electronegativities decrease down a group, but increase across a period. As
expected, the most electronegative elements are the reactive non-metals in the top right-hand corner of the
periodic table. In contrast, the least electronegative elements are the reactive metals in the bottom left-hand
corner.
Moving from one element to the next across a period, the nuclear charge increases by one unit and
one electron is added to the outer shell. As the positive charge on the nucleus rises, the atom has an
increasing electron-attracting power and therefore an increasing electronegativity.
Electronegativity generally decreases down a group since both increasing atomic size and the
screening effect of inner electrons reduce the attraction of the nucleus for electrons.
As a result of these two trends, there are some considerable diagonal similarities between elements
in different groups :
Group I
Group II
Group III
Group IV
Li
Be
B
C
Na
Mg
Al
Si
Electronegativity values can be used to estimate the polarity of different bonds. The bonds between
elements of widely different electronegativities (i.e. between a metal and non-metal) will be ionic. The
bonds between elements of similar electronegativity will be non-polar or slightly polar. If the two elements
are non-metals the bonding will be covalent. As a result of differing electronegativities, all ionic bonds have
some covalent character and most covalent bonds have some ionic character.
Unit 7-6
4. Melting points
Period 2
Melting point (oC)
Li
180
Be
1280
B
2030
C
3700(G)
3550(D)
Period 3
Melting point (oC)
Na
98
Mg
650
Al
660
Si
1410
N2
-210
O2
-219
P4
S8
44
119
(white) (rhombic)
F2
-220
Ne
-250
Cl2
-101
Ar
-189
Melting point (oC)
4000
3500
3000
2500
2000
1500
1000
500
0
-5 0 0
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Moving across a period from left to right, the melting point rises through the metals and metalloids
and then drops abruptly to low values for the non-metals.
Giant metallic structures :
Metals usually have high melting points. In the metal lattice, each positively charged ion is attracted
to the ‘cloud’ of negative electrons. Since the mobile valence electrons are responsible for the bonding in
metals, it is not surprising that moving from sodium (one outer shell electron) through magnesium (two
outer shell electrons) to aluminium (three outer shell electrons) the bonding gets gradually stronger. Thus,
the melting point rises from Na to Mg to Al. Similar trends are observed in the second period from Li to Be.
Giant molecular (covalent) structures :
The metalloids (boron, graphite and silicon) and the non-metal diamond have giant molecular
structures. The covalent linking in these elements extends from one atom to the next through the whole
lattice forming a three-dimensional giant molecule. The strong covalent bonds hold each atom tightly in the
crystal and it is extremely difficult to break one atom away from the lattice. Thus, these elements have very
high melting points as most of their covalent bonds are broken at melting.
Simple molecular structures :
All the non-metals in periods 2 and 3 (except diamond) form molecular structure. Each of these
elements consists of separate, small molecules, i.e. N2 , O2 , F2 , Ne , P4 , S8 , Cl2 , Ar . There are only very
weak intermolecular forces (van der Waals forces) between the separate molecules. Consequently, the
small distinct molecules can be separated easily and these non-metals have low melting points. Moreover,
as van der Waals force is directly proportional to the size and number of electrons in a molecule, melting
points of period 3 non-metals decrease in the following order : S8 > P4 > Cl2 > Ar
Unit 7-7
Section 7.2
Periodic relationship among the Oxides of the elements
Li to Cl
( 1 ) Bonding and stoichiometric composition of the oxides
Stoichiometric composition of the oxides of elements in period 2 and period 3 :
Bonding and structure of the oxides of elements in period 2 :
Formula of oxide
En(O) - En(X)
State at s.t.p.
Conductivity of
liquid
Bonding
Structure
Li2O
2.5
Solid
Good
Ionic
BeO
2.0
Solid
Fairly
poor
Ionic
B2O3
1.5
Solid
Very
poor
Covalent
CO2
1.0
Gas
Nil
N2O4
0.5
Liquid
Nil
O2
0
Gas
Nil
F2O
-0.5
Gas
Nil
Covalent
Covalent
Covalent Covalent
Bonding and structure of the oxides of elements in period 3 :
Formula of oxide
En(O) - En(X)
State at s.t.p.
Conductivity of
liquid
Bonding
Structure
Na2O
2.6
Solid
Good
MgO
2.3
Solid
Good
Al2O3
2.0
Solid
Good
Ionic
Ionic
Ionic
Bonding and structure of the oxides :
SiO2
1.7
Solid
Very
poor
Covalent
P4O10
1.4
Solid
Nil
SO2
1.0
Gas
Nil
Cl2O
0.5
Gas
Nil
Covalent
Covalent
Covalent
Unit 7-8
In each period, the oxides of metals and metalloids have giant structures, whereas the oxides of
non-metals are composed of simple molecules.
Notice the gradations in structure and bond type across each period from ionic oxides to simple
molecular oxides. The graduation can be correlated with changes in electronegativity across the period
from low values on the left to high values on the right. Thus, atoms of low electronegativity, such as Na, Mg
and Li, form ionic oxides in which they have given up electrons to oxygen atoms. In contrast to this,
compounds formed between the more electronegative atoms (such as Si, N, S and F) and oxygen exist as
discrete molecules (e.g. NO2, SO2, F2O) or as giant covalent structure (e.g. SiO2).
Giant ionic structures :
Lithium oxide (Li2O), beryllium oxide (BeO), sodium oxide (Na2O), sodium peroxide (Na2O2),
magnesium oxide (MgO), and aluminium oxide (Al2O3), may be regarded as ionic structures.
Consequently, they are solids at room temperature, with high melting points and boiling points. These ionic
oxides will conduct electricity in the molten state.
Giant molecular (covalent) structures :
The metalloids, boron and silicon, form oxides with giant molecular structures.
In boron oxide (B2O3), boron and oxygen atoms are arranged in layers. Strong covalent bonds link
one atom to the next in a giant sheet structure. The giant structures are therefore solids with high melting
points and boiling points. The melting point of B2O3 is 557oC and that of SiO2 is 1700oC.
Simple molecular structures :
The simple molecular oxides are much more volatile than the ionic metal oxides and the giant
molecular metalloids oxides. They have low melting points and low boiling points and they do not conduct
electricity in the liquid state.
( 2 ) Behaviour of the oxides with water, dilute acids and dilute alkalis
Unit 7-9
Formula of oxide
Action of water
Li2O
Reacts to
form
LiOH(aq)
alkaline
solution
BeO
Does not
react with
water
Na2O
Reacts to
form
NaOH(aq)
alkaline
solution
MgO
Reacts to
form
Mg(OH)2(aq)
weakly
alkaline
solution
B2O3
Reacts to
form
H3BO3,
a very
weak
acid
CO2
Reacts
to form
H2CO3,
a weak
acid
NO2
Reacts to
form an
acid
solution
of HNO3
and
HNO2
O2
F2O
Reacts
slowly
forming
O2 and
an
acidic
solution
of HF
SO2
Reacts
to form
H2SO3
acid
solution
Cl2O
Reacts
to form
HClO
acid
solution
Acid-base nature
Formula of oxide
Action of water
Al2O3
Does not
react with
water
SiO2
Does not
react with
water
P4O10
Reacts
to form
H3PO4
acid
solution
Acid-base nature
Moving from left to right across a period, the oxides of the elements change from ionic oxides of
metals, which are basic or amphoteric, to the oxides of metalloids with giant structure, which are weakly
acidic, and finally to the simple molecular oxides of non-metals, which are acidic.
Basic oxides :
The ionic oxides contain O2- ions in the crystal lattice. Thus, the oxides of Group I metals react
vigorously with water to form alkaline solutions.
O2-(s) + H2O(l) → 2 OH-(aq)
Li2O(s) + H2O(l) → 2 Li+(aq) + 2 OH-(aq)
Na2O(s) + H2O(l) → 2 Na+(aq) + 2 OH-(aq)
These oxides would react even more vigorously with acids.
Na2O(s) + 2 H+(aq) → 2 Na+(aq) + H2O(l)
The peroxides of Group I metals react with water to form alkaline solution containing hydrogen
peroxide.
O22-(s) + 2 H2O(l) → 2 OH-(aq) + H2O2(aq)
The oxides of Group II metals do not react so readily with water or acids since the large charge
density on the cation holds the O2- ions more firmly. Thus, MgO is only slightly soluble in water, although
it reacts readily with acids.
Amphoteric oxides :
Oxides which show both basic and acidic properties are called amphoteric oxides.
Unit 7-10
BeO is insoluble in water, but it shows basic properties by dissolving in acids to form Be2+ salts.
However, BeO also resembles acidic oxides by reacting with alkalis to form beryllate.
Aluminium oxide is also amphoteric. It does not react with water, but it will react with both H+ and
OH ions.
-
Neutral oxides :
CO, N2O , NO and ClO2 are neutral oxides which show neither acidic nor basic character.
Acidic oxides :
Boron (III) oxide (B2O3) reacts slightly with water to form boric (III) acid, H3BO3, a very weak acid.
CO2 dissolves in water and reacts slightly to form weak carbonic acid, H2CO3 .
CO2(aq) + H2O(l) – H2CO3(aq)
CO2 will react with the OH- ions in alkalis to form first hydrogencarbonate, HCO3- and then carbonate,
CO3 .
CO2(aq) + OH-(aq) – HCO3-(aq)
HCO3-(aq) + OH-(aq) – CO32-(aq) + H2O(l)
2-
Silicon (IV) oxide (SiO2) does not react with water, but it reacts with concentrated alkalis forming
silicate ions, SiO32- .
N2O3 reacts with water to form nitric (III) acid, HNO2 .
N2O5 reacts with water to form nitric (V) acid, HNO3 .
NO2 reacts with water to form a mixture of two acids, HNO2 and HNO3 .
The oxides of P, S and Cl react readily with water to form strong acids.