Download I.5. Periodic properties of the elements

I. STRUCTURE OF SUBSTANCES
I.3. The order of filling orbitals
Element
1s
2s
2px 2py 2pz
Electron configuration
H (Z = 1)
1s1
He (Z = 2)
1s2
Li (Z = 3)
1s2 2s1
Be (Z = 4)
1s2 2s2
B (Z = 5)
1s2 2s22p1
C (Z = 6)
1s2 2s22p2
N (Z = 7)
1s2 2s22p3
O (Z = 8)
1s2 2s22p4
F (Z = 9)
1s2 2s22p5
Ne (Z = 10)
1s2 2s22p6 (stable configuration)
(stable configuration)
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I. STRUCTURE OF SUBSTANCES
I.3. The order of filling orbitals
• The configuration of an element differs from the previous
element only by an electron named “differentiating electron”.
• The differentiating electron is placed on the highest energy
orbital.
Examples: Write the electron configuration for following
elements:
Cl (Z = 17); Ca (Z = 20); Mn (Z = 25); Al (Z = 13)
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I. STRUCTURE OF SUBSTANCES
I.4. Periodic table of the elements
The periodic table was developed based on three fundamental
ideas that have developed over time:
• the tendency to find a natural classification of the elements
• the certainty that there is a relationship between a fundamental
property characteristic of each element and the chemical behavior of
that element
• the existence of a periodicity of the properties of elements
A classification scheme of the elements, similar to that used
today, was discovered independently by Dmitri Mendeleev and Luther
Meyer in 1869.
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***Ununtrium was first detected in 2003 in the decay of ununpentium and was synthesized directly in 2004.
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Only fourteen atoms of ununtrium have been observed to date. The longest-lived isotope known is 286Uut
with a
half-life of ~20 s, allowing first chemical experiments to study its chemistry.
I. STRUCTURE OF SUBSTANCES
I.4. Periodic table of the elements
In the periodic table, the elements are presented in the order of
increasing atomic number. The periodic table contains 18 columns called
“groups” and 7 rows called “periods”.
Periods:
1st period: consists of only two elements: hydrogen (H) and helium (He).
2nd and 3rd period: have eight elements each
4th and 5th period: have 18 elements each
6th period: contains 32 elements. From this period 14 elements are
extracted and placed at the bottom of the table. This series of 14
elements, which fits between lanthanum (La, Z=57) and hafnium (Hf,
Z=72) is called the lanthanides or rare earth series.
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I. STRUCTURE OF SUBSTANCES
I.4. Periodic table of the elements
7th period: is incomplete for the moment, but is believed to be as
long as the sixth one. A series of 14 elements, extracted from the
7th period and placed at the bottom of the table is called the
actinide series.
Groups:
• Group 1: the atoms of the elements in group 1 have a single
outer-shell electron placed in an s orbital. Elements of the first
group are called alkali-metals. Electron configuration ns1
• Group 2: the atoms of the elements from group 2 have 2
electrons in an outer shell (in an s orbital). These elements are
alkaline earth metals. Electron configuration ns2
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I. STRUCTURE OF SUBSTANCES
I.4. Periodic table of the elements
• Group 13: the elements of group 13 have 3 electrons in the outer
shell, two s electrons and one p electron. The p electron is the
differentiating electron. Electron configuration ns2np1
• Group 14: elements have 4 electrons in the outer shell (ns2np2)
• Group 15: elements have 5 electrons in the outer shell (ns2np3)
• Group 16: elements have 6 electrons in the outer shell (ns2np4)
• Group 17: elements have 7 electrons in the outer shell (ns2np5)
• Group 18: elements have 8 electrons in the outer shell (ns2np6)
Elements of group 18 have an outer shell full of electrons = stable
configuration.
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I. STRUCTURE OF SUBSTANCES
I.4. Periodic table of the elements
Group 1 + Group 2 = s block (their properties arise from the
presence of s electrons)
Group 13+14+15+16+17+18 = p block (their properties depend
on the presence of p electrons)
Group 3+4+5+6+7+8+9+10+11+12 = d block or transition
elements (their properties depend on the presence of d
electrons)
Lanthanides + Actinides = f block (their properties arise from
the presence of f electrons)
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
The elements of Group 18, rare gases, have the configuration
ns2 np6, except helium, whose configuration is 1s2. That means the
outer shells of the atoms are full. These prove to be very stable
configurations and they can be altered with great difficulty. As a
result, rare gases have a very low reactivity, they are also known as
noble gases.
The electron configuration of the elements of groups 1 and 2
differ from these of noble gases by only one or two electrons in the s
orbital of a new shell.
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
K
Z = 19: 1s22s22p63s23p64s1
or
[Ar] 4s1
Ca
Z = 20: 1s22s22p63s23p64s2
or
[Ar] 4s2
Except hydrogen, the elements of groups 1 and 2 are metals. The
characteristic chemical properties of metallic elements are based on
the ease of removal of one or more electrons from their atoms
to produce positive ions:
K → K+ + e -
Ca → Ca2+ + 2e-
Some physical properties of metals (ability to conduct heat and
electricity, ductility, malleability) also arise from these distinctive
electron configurations.
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
Elements of the groups 16 and 17 have an electron configuration
with two or one electron less that the corresponding noble gas. Atoms of
these elements can realize the electronic configuration of a noble gas by
gaining the appropriate number of electrons. For example, the
electron configuration of S becomes that of Ar by gaining two electrons:
S
Z = 16
S
+
2e-
[Ne] 3s23p4
 S2[Ar]
The sulfur atom becomes sulfide anion (S2-). Similarly, the chlorine
atom becomes chloride anion (Cl-)
Cl
Z = 17
Cl
+ e-
[Ne] 3s23p6
 Cl[Ar]
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
These elements whose atoms can acquire a noble gas
configuration by a small number of electrons are non – metals. Non
– metals are H from group 1, C from group 14, N and P from group
15, O, S and Se from group 16 and F, Cl, Br and I from group 17.
B (13), Si, Ge (14), As, Sb (15), Tc, Po (16) and At (17) are
metalloids or semi – metals.
18th group is a special family of elements, but noble gases
may be considered non – metals. The rest of the elements, including
of course the lanthanides and actinides are metals.
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
1) Atomic radius – is determined by the number of electronic shells.
The atomic radius increases
in
a
group
from
top
to
bottom and decreases in a
period from left to right.
Atomic radii [in pm]
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
• in a group, the atomic radius increases because the number of
electronic shells increases. The outer shell electrons are further and
further from the nucleus, therefore less attracted by the positive
charge of the nucleus
the atoms get larger.
• in a period, the atomic radius decreases from left to right, because
the charge of the nucleus (nr. of protons Z) increases but the
electrons are still filling the same shell. The outer shell electrons are
attracted more strongly by the nucleus and, as a result, the atomic
radius decreases from left to right through a period.
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
2) Ionic radius
• When electrons are removed from a metal atom a positive ion
(cation) is formed. Cations have a smaller ionic radius than the
corresponding atom.
• When the atoms of a non-metal accept electrons a negative ion
(anion) is formed. The anions have larger ionic radius than the
corresponding atom.
• in a group, the ionic radius increases from top to bottom.
• in a period, the ionic radius decreases from left to right.
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
The relative sizes of the cations
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
The relative sizes of the anions
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
For example, in the series of cations Na+, Mg2+, Al3+ the
number of electrons is the same (10), while the number of protons
increases together with the atomic number Z.
Al3+ is smaller that Mg2+ because the electrostatic force
between the 10 electrons and the nuclear charge of Al (+13) is
more powerful than that between the 10 electrons and the nuclear
charge of Mg (+12).
Na+
Mg2+
Al3+
No. of protons
11
12
13
No. of electrons
10
10
10
Ionic radius [Å]
0.95
0.65
0.50
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
N3-
O2-
F-
No. of protons
7
8
9
No. of electrons
10
10
10
Ionic radius [Å]
1.71
1.40
1.36
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
3) Ionization energy – Is the energy required to remove one
electron from an individual atom in the gaseous phase. This is the
first ionization energy.
In case of metals, which have a small number of electrons in the
outer shell, a small amount of energy is needed to remove an
electron, that is metals have low ionization energies.
Inside a group, the ionization energy tends to decrease from
top to bottom because the attraction force of the nucleus decreases
in the same way and the electron is more easily removed.
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I. STRUCTURE OF SUBSTANCES
I.5. Periodic properties of the elements
Nonmetals have large ionization energies because they
have a large number of electrons in the outer shell.
Nonmetals
tend to gain, not to lose electrons. Ionization energies tend to
increase from left to right along a period of the periodic table.
In general, the elements that appear in the lower left
region of the periodic table have the lowest ionization energies
and are therefore the most chemically active metals. On the other
hand, the elements with the highest ionization energies occur in
the upper right hand region of the periodic table.
The first ionization energy of the elements is a function of
atomic number Z:
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The first ionization energy of the elements
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I. STRUCTURE OF SUBSTANCES
A property used to describe the type of bond that
results when atoms combine is electronegativity.
Electronegativity describes the ability of an atom to
attract electrons towards itself. The most widely used
electronegativity scale was devised by Linus Pauling.
Pauling’s
electronegativities
are
dimensionless
numbers ranging from about 1 for very active metals to
4.0 for fluorine, the most active nonmetal.
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1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
H
2,20
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He
Li
0,98
Be
1,57
B
2,04
C
2,55
N
3,04
O
3,44
F
3,98
Ne
Na
0,93
Mg
1,31
Al
1,61
Si
1,90
P
2,19
S
2,58
Cl
3,16
Ar
K
0,82
Ca
1,00
Sc
1,36
Ti
1,54
V
1,63
Cr
1,66
Mn
1,55
Fe
1,83
Co
1,88
Ni
1,91
Cu
1,90
Zn
1,65
Ga
1,81
Ge
2,01
As
2,18
Se
2,55
Br
2,96
Kr
3,0
Rb
0,82
Sr
0,95
Y
1,22
Zr
1,33
Nb
1,6
Mo
2,16
Tc
1,9
Ru
2,2
Rh
2,28
Pd
2,2
Ag
1,93
Cd
1,69
In
1,78
Sn
1,96
Sb
2,05
Te
2,1
I
2,66
Xe
2,6
Cs
0,79
Ba
0,89
La
1,27
Hf
1,3
Ta
1,5
W
2,36
Re
1,9
Os
2,2
Ir
2,20
Pt
2,28
Au
2,54
Hg
2,0
Tl
1,62
Pb
2,33
Bi
2.02
Po
2,0
At
2,2
Rn
Fr
0,7
Ra
0,9
Ac
1,10
Pauling’s electronegativities of the elements
As a rough rule, most metals have electronegativities of about
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1.7 or less; semi-metals about 2 and nonmetals greater than 2.