1 I. STRUCTURE OF SUBSTANCES I.3. The order of filling orbitals

1

I. STRUCTURE OF SUBSTANCESI.3. The order of filling orbitals

Element 1s 2s 2px 2py 2pz Electron configuration H (Z = 1) 1s1 He (Z = 2) 1s2 (stable configuration) 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)

2

I. STRUCTURE OF SUBSTANCES

• 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)

I.3. The order of filling orbitals

3

I. STRUCTURE OF SUBSTANCESI.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.

4***Ununtrium was first detected in 2003 in the decay of ununpentium and was synthesized directly in 2004. 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.

5


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.

6


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

7

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)





Elements of group 18 have an outer shell full of electrons = stable

configuration.

8

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)


9

I. STRUCTURE OF SUBSTANCESI.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.

10


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.

11


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:

The sulfur atom becomes sulfide anion (S2-). Similarly, the chlorine

atom becomes chloride anion (Cl-)

S + 2e- S2-

[Ne] 3s23p4 [Ar]

Cl + e- Cl-

[Ne] 3s23p6 [Ar]

S Z = 16

Cl Z = 17

12


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.

13


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]

14


• 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.

15


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.

16


The relative sizes of the cations

17


The relative sizes of the anions

18


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

I.5. Periodic properties of the elements

19



N3- O2- F-

No. of protons 7 8 9

No. of electrons 10 10 10

Ionic radius [Å] 1.71 1.40 1.36

20



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.



21

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:



22The first ionization energy of the elements

23

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.


24

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

H2,20

He

Li0,98

Be1,57

B2,04

C2,55

N3,04

O3,44

F3,98

Ne

Na0,93

Mg1,31

Al1,61

Si1,90

P2,19

S2,58

Cl3,16

Ar

K0,82

Ca1,00

Sc1,36

Ti1,54

V1,63

Cr1,66

Mn1,55

Fe1,83

Co1,88

Ni1,91

Cu1,90

Zn1,65

Ga1,81

Ge2,01

As2,18

Se2,55

Br2,96

Kr3,0

Rb0,82

Sr0,95

Y1,22

Zr1,33

Nb1,6

Mo2,16

Tc1,9

Ru2,2

Rh2,28

Pd2,2

Ag1,93

Cd1,69

In1,78

Sn1,96

Sb2,05

Te2,1

I2,66

Xe2,6

Cs0,79

Ba0,89

La1,27

Hf1,3

Ta1,5

W2,36

Re1,9

Os2,2

Ir2,20

Pt2,28

Au2,54

Hg2,0

Tl1,62

Pb2,33

Bi2.02

Po2,0

At2,2

Rn

Fr0,7

Ra0,9

Ac1,10

Pauling’s electronegativities of the elements

As a rough rule, most metals have electronegativities of about

1.7 or less; semi-metals about 2 and nonmetals greater than 2.

Documents

1 I. STRUCTURE OF SUBSTANCES I.3. The order of filling orbitals