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.

***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.

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.

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

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.

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)

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.

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.

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 + 2e-  S2- [Ne] 3s23p4 [Ar] S Z = 16 The sulfur atom becomes sulfide anion (S2-). Similarly, the chlorine atom becomes chloride anion (Cl-) Cl + e-  Cl- [Ne] 3s23p6 [Ar] Cl Z = 17

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.

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]

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.

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.

The relative sizes of the cations I. STRUCTURE OF SUBSTANCES I.5. Periodic properties of the elements The relative sizes of the cations

I. STRUCTURE OF SUBSTANCES I.5. Periodic properties of the elements The relative sizes of the anions

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 Ionic radius [Å] 0.95 0.65 0.50

I. STRUCTURE OF SUBSTANCES I.5. Periodic properties of the elements N3- O2- F- No. of protons 7 8 9 No. of electrons 10 Ionic radius [Å] 1.71 1.40 1.36

I. STRUCTURE OF SUBSTANCES I. STRUCTURE OF SUBSTANCES I.5. Periodic properties of the elements 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.

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:

The first ionization energy of the elements

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.

Pauling’s electronegativities of the elements 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 H 2,20 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 Zn 1,65 Ga 1,81 Ge 2,01 As 2,18 Se Br 2,96 Kr 3,0 Rb 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 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 Os Ir Pt Au 2,54 Hg 2,0 Tl 1,62 Pb 2,33 Bi 2.02 Po At 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 1.7 or less; semi-metals about 2 and nonmetals greater than 2.