Periodic Table and Atomic Properties, Chemistry tutorial


The chemistry of the elements is hugely varied. But amidst that variety there are patterns, and the best recognized and most helpful is chemical periodicity: if the elements are laid out in order of atomic number, alike elements take place at regular intervals.

The properties of the elements exhibit trends. Such trends can be expected using the periodic table and can be illustrated and understood through analyzing the electron configurations of the elements. Components tend to gain or lose valence electrons to attain stable octet formation. Stable octets are seen in the inert gases, or noble gases, of Group VIII of the periodic table. In addition to this activity, there are 2 other significant trends:

  • 1st, electrons are added one at a time moving from left to right across a period. As this take place, the electrons of the outermost shell knowledge increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more strongly leap to it. 
  • 2nd, moving down a column in the periodic table, the outermost electrons happen to less forcefully bound to the nucleus. These take place because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) amplifies downward within each group. Such trends clarify the periodicity examined in the elemental properties of atomic radius, ionisation energy, electron affinity, and electronegativity.

Organization of the Periodic Table

From the table highlighted in Module 1 that shows the long form of table with the "block" structure emphasized. You will recall that the two f blocks are written at the bottom merely to keep the table from becoming inconveniently wide; these two blocks actually go in between La-Hf and Ac-Db, respectively, in the d block.

Table: The Periodic Table of Elements

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To comprehend how the periodic table is organized, visualize that we write down a long horizontal list of the components in order of their increasing atomic number. It would start this way:

H  He  Li  Be  B  C  N  O  F  Ne  Na  Mg  Al  Si  P  S  Cl  Ar  K  Ca...

Now if we look at the diverse physical and chemical properties of such elements, we would discover that their values tend to amplify or reduce throughout Z in a manner that reveals a repeating pattern-that is, a periodicity. For the elements listed above, such breaks can be indicated via the vertical bars following in colour:

H  He | Li   Be   B   C   N   O F   Ne | Na Mg Al Si P S  Cl Ar |Ca... 


To construct the table, we put each sequence in a divide row, which is recognized as a period. The rows are aligned in such a way that the elements in each vertical column possess definite similarities. Therefore the 1st short-period elements H and He are chemically similar to the elements Li and Ne at the starting and end of the 2nd period. The 1st period is dividing in order to place H above Li and He above Ne. 

The 'block' nomenclature following above terms to the sub-orbital kind (quantum number l, or s-p-d-f categorization) of the highest-energy orbitals that are occupied in a given element. For  n = 1  there is no  p block, and the s block is split so that helium is situated in the similar group as the other inert gases, which it resembles chemically. For the 2nd period (n = 2), there is a p block but no d block; in the common 'long form' of the periodic table it is customary to leave a gap between such 2 blocks in order to accommodate the  d  blocks that occur at  n  = 3 and above. At n = 6 we introduce an f block, but in order to hold the table to reasonable dimensions the  f blocks are situated below the major body of the table.


Each column of the periodic table is known as a group. The elements belonging to a following group bear a tough similarity in their chemical behaviors. 

In the past, 2 dissimilar systems of Roman numerals and letters were utilized to indicate the different groups. North Americans added the letter B to signify the d-block groups and A for the others; this is the system following in the table above. The rest of the world utilized A for the d-block elements and B for the others. In the year 1985, a new international system was adopted in that the columns were simply labelled 1-18. Even though this system has met enough opposition in North America to sluggish its incorporation into textbooks, it seems likely that the 'one to eighteen' system will slowly take over as older professors (the main hold-outs!) retire.


Chemists have long established it convenient to refer to the elements of dissimilar groups, and in several cases of spans of groups through the names indicated in Table. The 2 of such that are most significant for us to know are the noble gases and the transition metals. 

Table: Groups of Elements in the Periodic Table 

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The Shell Model of the Atom

The properties of an atom based eventually on the number of electrons in the numerous orbitals, and on the nuclear charge that measures the compactness of the orbitals. In order to relate the properties of the elements to their locations in the periodic table, it is frequently convenient to build utilize of a simplified view of the atom in that the nucleus is enclosed through one or more concentric spherical "shells", each of which consists of the highest-principal quantum number orbitals (always s- and p-orbitals) that surround at least one electron. The shell model (as through any scientific model) is less a description of the world than a simplified way of looking at it that assists us to understand and correlate diverse phenomena. The principal simplification here is that it deals only through the major group elements of the s- and p-blocks, skipping the d- and f-block elements whose properties tend to be less directly tied to their group numbers. 

Table: Shell Model of the Atom showing Valence Shells of the First Eighteen Elements

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The electrons (indicated via the dots) in the outer-most shell of an atom are the ones that relate most eagerly through other atoms, and therefore play a main role in governing the chemistry of an element. As we know utilize of noble-gas symbols to simplify the electron-configuration notation. In exacting, the number of outer-shell electrons (that is specified via the rightmost digit in the group number) is a main determinant of an element's 'combining power', or valence. The common trend is for an atom to increase or lose electrons, either straight (leading to formation of ions) or via sharing electrons through other atoms so as to attain an outer-shell configuration of s2p6. This configuration, recognized as an octet, corresponds to that of one of the noble-gas elements of Group 18. 

  • The elements in Groups 1, 2 and 13 tend to provide up their valence electrons to form positive ions these as Na+, Mg2+and Al3+, as well as compounds NaH, MgH2 and AlH3. The outer-shell configurations of the metal atoms in such species correspond to that of neon.
  • Elements in Groups 15-17 tend to obtain electrons, forming ions such as P3-, S2-and Cl- or compound these as PH3, H2S and HCl. The outer-shell configurations of such elements correspond to that of argon.
  • The Group 14 elements don't usually form ions at all, but share electrons through other elements in tetravalent compounds such as CH4.

The above diagram shows the 1st three rows of what are recognized as the representative elements- that is, the s- and p-block elements only. As we move farther down (into the fourth row and below), the presence of d-electrons exerts a complicating influence that permits elements to exhibit numerous valances. This consequence is particularly evident in the transition-metal elements; this is the cause for not including the d-block through the ambassador elements at all. 

Effective nuclear charge

Those electrons in the outmost or valence shell are particularly significant since they are the ones that can engage in the sharing and swap that is dependable for chemical reactions; how strongly they are bound to the atom determines much of the chemistry of  the element. The degree of binding is the consequence of two opposing forces:

  • The attraction between the electron and the nucleus
  • The repulsions between the electron in question and all the other electrons in the atom.

All that matters is the net force, the dissimilarity between the nuclear attraction and the totality of the electron-electron repulsions. We can make simpler the shell model even additional through visualize that the valence shell electrons are the only electrons in the atom, and that the nuclear charge has whatever value that would be needed to bind such electrons as tightly as it is examined experimentally. Since the number of electrons in this model is less than the atomic number Z, the required nuclear charge will as well be smaller; this is recognized as the effective nuclear charge. Effective nuclear charge is basically the positive charge that a valence electron 'sees'.

Part of the difference between Z and Zeffective is due to other electrons in the valence shell, but this is generally only a minor contributor since such electrons tend to act as if they are extends out in a diffuse spherical shell of larger radius. The major actors here are the electrons in the much more compact inner shells which enclose the nucleus and exert what is often termed a shielding or "screening" consequence on the valence electrons. 

Table: Effective Nuclear Charges of the first 12 Elements

2302_Effective Nuclear Charges of the first 12 Elements.jpg

The formula for calculating effective nuclear charge is not very complicated, but we will skip a discussion of it here. An even simpler although rather crude procedure is to just subtract the number of inner-shell electrons from the nuclear charge; the result is a form of effective nuclear charge which is called the core charge of the atom.


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Fig: Calculation of Effective Nuclear Charge

Sizes of Atoms and Ions

The concept of 'size' is somewhat uncertain when applied to the scale of atoms and molecules. The reason for this is apparent when we remind that an atom has no specific boundary; there is a finite (but extremely small) probability of discovering the electron of a hydrogen atom, for instance, 1 cm, or even 1 km from the nucleus. It isn't possible to identify a definite value for the radius of an isolated atom; the best we can do is to describe a spherical shell within whose radius several arbitrary percentage of the electron density can be originate.

When an atom is joined through other atoms in a solid element or compound, an effective radius can be verified through observing the distances between adjacent rows of atoms in such solids. This is most generally carried out by X-ray scattering experiments. Since of the different ways in which atoms can aggregate together, numerous different types of atomic radii can be defined. 

Distances on the atomic scale have traditionally been expressed in Ångstrom (Å) units (1 Å=10-8  cm = 10-10 m); but nowadays the picometer is preferred:

1 pm = 10-12 m = 10-10cm = 10-2 Å, or 1Å = 100 pm. The radii of atoms and ions are typically in the range of 70 - 400 pm.

A rough idea of the size of a metallic atom can be attained simply via measuring the density of an example of the metal. This provides us the number of atoms per unit volume of the solid. The atoms are supposed to be spheres of radius r in contact through each other, each of which sits in a cubic box of edge length 2r. The volume of each box is just the total volume of the solid separated via the no of atoms in that mass of the solid; the atomic radius is the cube root of r.

Even though the radius of an atom or ion can't be calculated directly, in most cases it can be inferred from measurements of the distance between adjacent nuclei in a crystalline solid. This is most generally carried out via X-ray scattering experiments. Since these solids fall into numerous diverse classes, numerous types of atomic radius are described.

Many atoms have numerous different radii; for instance, sodium forms a metallic solid and therefore has a metallic radius, it forms a gaseous molecule Na2 in the vapour phase (covalent radius), and of course, it shapes ionic solids such as NaCl.

Metallic radius is half the distance between nuclei in a metallic crystal.

Covalent radius is half the distance between like atoms that are bonded together in a molecule.

Van der Waals radius is the effective radius of adjacent atoms which aren't chemically linked in a solid, but are presumably in 'contact'. 

An instance would be the distance between the iodine atoms of adjacent I2 molecules in crystalline iodine.

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Fig.: Sizes of Atoms    

Ionic radius is the effective radius of ions in solids these as NaCl. It is simple sufficient to determine the distance between adjacent rows of Na+ and Cl-ions in such a crystal, but there is no definite way to decide what portions of this distance are attributable to each ion. The best one can do is creating estimations depend on studies of numerous different ionic solids (LiI, KI, NaI, for instance) that enclose one ion in common. Many such estimates have been made, and they turn out to be remarkably consistent.

The lithium ion is suitably small that in LI, the iodide ions are in contact, so I-I distances are twice the ionic radius of I-. This isn't true for KI, but in this solid, adjacent potassium and iodide ions are in contact, allowing estimation of the K+ ionic radius.

189_Ionic Radius of I.jpg

Fig: Ionic Radius of I-

Many atoms have numerous different radii; for instance, sodium forms a metallic solid and therefore has a metallic radius. It as well shapes a gaseous molecule Na2 in the vapour phase (covalent radius), and of course it forms ionic solids as stated above. 

Periodic Trends in Atomic Properties

Atomic size

We would suppose the size of an atom to depend mostly on the principal quantum number of the maximum occupied orbital; in other terms, on the 'number of occupied electron shells'. Because each row in the periodic table corresponds to an increase in n, atomic radius raises as we move down a column. The other significant feature is the nuclear charge; the higher the atomic number, the more strongly will the electrons be drawn toward the nucleus, and the smaller the atom. This consequence is dependable for the contraction we examine as we move across the periodic table from left to right.

Table: Covalent Radii of Elements

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Table illustrates a periodic table in that the sizes of the atoms are symbolized graphically. The apparent discontinuities in this diagram reflect the complexity of comparing the radii of atoms of metallic and non-metallic bonding kinds. Radii of the noble gas elements are estimates from those of nearby elements.

Ionic radii

A positive ion is always smaller than the neutral atom. This is due to the reduced electron-electron repulsion. If a 2nd electron is lost, the ion gets even smaller; for instance, the ionic radius of Fe2+is 76 pm, while that of Fe3+ is 65 pm. If formation of the ion involves complete emptying of the outer shell, then reduce in radius is especially great.

The hydrogen ion H+ is in a class through itself; having no electron cloud at all, its radius is that of the bare proton, or about 0.1 pm-a contraction of 99.999%. Since the unit positive charge is concentrated into such a small volume of space, the charge density of the hydrogen ion is extremely high; it interacts extremely robustly by other matter, including water molecules, and in aqueous solution, it exists only as the hydronium ion H3O+

Negative ions are always larger than the parent ion; the addition of one or more electrons to an existing shell amplifies electron-electron repulsion which consequences in a common expansion of the atom.

Table: Ionic Radii of Elements

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An isoelectronic series is a sequence of species all having the similar number of electrons (and hence, the similar amount of electron-electron repulsion) but varying in nuclear charge. Of course, only one member of such a sequence can be a neutral atom (for instance, neon in the series shown below.) The consequence of rising nuclear charge on the radius is visibly seen.

929_Isoelectronic Series.jpg

Fig: Isoelectronic Series

Periodic trends in ion formation Chemical reactions are based largely on the interactions between the most loosely bound electrons in atoms, so it is not surprising that the tendency of an atom to gain, lose or share electrons is one of its fundamental chemical properties.

Ionisation energy

This term refers to the formation of positive ions. In order to remove an electron from an atom, work must be done to overcome the electrostatic attraction between the electron and the nucleus; this work is termed the ionisation energy of the atom and corresponds to the exothermic process in which M(g) stands for any isolated (gaseous) atom.

M(g)  → M+(g) + e-

An atom has as many ionisation energies as it has electrons. Electrons are always removed from the highest-energy occupied orbital. An examination of the successive ionisation energies of the first ten elements (below) provides experimental confirmation that the binding of the 2 innermost electrons (1s orbital) is significantly diverse from that of the n = 2 electrons. Successive ionisation energies of an atom increase rapidly as the reduction in the electron-electron repulsion reasons the electron shells to contract; therefore binding the electrons even more tightly to the nucleus. 

Table: Successive Ionizations of the first Ten Elements

1041_Successive Ionisations of the first Ten Elements.jpg

Ionisation energies amplify through the nuclear charge Z as we move across the periodic table. They reduce as we move down the table since in each period; the electron is being eliminated from a shell one step farther from the nucleus than in the atom immediately above it.

This consequence in the familiar zigzag lines when the 1st ionisation energies are plotted as a function of Z.

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Fig: Ionisation Energies of the first Twenty Elements

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Fig: Plot of Ionisation Energies against Nuclear Charges of Elements-

This more detailed plot of the ionisation energies of the atoms of the 1st ten elements reveals several interesting irregularities that can be related to the  slightly lower energies (greater stabilities) of electrons in half-filled (spin-unpaired) relative to completely-filled sub-shells.

At last, a more comprehensive analysis of the ionisation energies of the chief group elements is following below:

Table: First Ionisation Energies of Elements

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Some points to note:

  • The noble gases have the maximum ionisation energies of any element in the period. This has nothing to do through any mysterious 'special stability' of the s2p6 electron configuration; it is basically a matter of the high nuclear charge acting on more contracted orbitals.
  • Ionisation energies (in addition to many other properties) tend not to differ greatly between the d-block elements. This reflects the fact that as the more-compact d orbitals are being filled, they exert a screening effect that partly offsets the increasing nuclear charge on the outermost s orbitals of higher principal quantum number.
  • Each of the Group 13 elements has lower first-ionisation energies than that of the element preceding it. The reversal of the ionisation energies trend in this group is often quality to the more easy elimination of the single outer-shell p electron compared to that of electrons contained in filled (and therefore spin-paired) s- and d-orbitals in the preceding elements.

 Electron affinity

Formation of a negative ion takes place when an electron from numerous external sources enters the atom and happens to incorporated into the lowest energy orbital that possesses a vacancy. Because the entering electron is attracted to the positive nucleus, the formation of negative ions is commonly exothermic. The energy given off is the electron affinity of the atom. For numerous atoms, the electron similarity emerges to be slightly negative, proposing that electron-electron repulsion is the dominant factor in such instances.

In common, electron affinities tend to be much smaller than ionisation energies, suggesting that they are managed via opposing factors having alike magnitudes. Such 2 factors are, as stated before, the nuclear charge and electron-electron repulsion. But the latter that is only a minor actor in positive ion formation, is now much more important.

One cause for this is that the electrons enclosed in the inner shells of the atom exert a collective negative charge that partially cancels the charge of the nucleus, therefore exerting a so-termed shielding effect which reduces the tendency for negative ions to form.

Since of such opposing consequences, the periodic trends in electron affinities aren't as clear as are those of ionisation energies. This is particularly evident in the 1st few rows of the periodic table, in which small consequences tend to be magnified anyway since an added electron produces a huge percentage, amplify in the number of electrons in the atom.

In common, we can say that electron affinities become more exothermic as we move from left to right across a period (owing to raised nuclear charge and smaller atom size). There are numerous interesting irregularities, although: 

  • In the Group 2 elements, the filled 2s orbital apparently shields the nucleus so efficiently that the electron affinities are slightly endothermic.
  • The Group 15 elements have rather low values, due possibly to the need to place the added electron in a half-filled p orbital; why the electron affinity of nitrogen should be endothermic is not clear. The vertical trend is for electron affinity to become less exothermic in succeeding periods owing to better shielding of the nucleus via more inner shells and the superior size of the atom, but here as well there are several apparent anomalies.


When 2 elements are connected in a chemical bond, the element that attracts the divided electrons more powerfully is more electronegative. Elements through low electronegativity (the metallic elements) are said to be electropositive. It is significant to comprehend that electronegativity is properties of atoms that are chemically bound to each other; there is no way of measuring the electronegativity of an isolated atom.

Furthermore, the similar atom can exhibit different electronegativities in diverse chemical environments, so the 'electronegativity of an element' is only a common guide to its chemical behaviour rather than a precise requirement of its behaviour in a particular compound.

However, electronegativity is eminently helpful in summarizing the chemical behaviour of an element. We will build substantial utilize of electronegativity when we learn chemical bonding and the chemistry of the individual elements.

Since there is no single definition of electronegativity, any numerical scale for measuring it must of requirement be somewhat arbitrary. Most of these scales are themselves based on atomic properties that are directly measurable and which as well relate in one way or the other to electron-attracting propensity. The most widely used of these scales was devised through Linus Pauling and is related to ionisation energy and electron affinity. The Pauling scale runs from 0 to 4; the highest electron affinity, 4.0, is assigned to fluorine, while caesium has the smallest value of 0.7. Values less than about 2.2 are usually connected through electropositive or metallic character.

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Fig: Electronegativity Scale

In the illustration of the scale revealed in the above figure, the elements are organized in rows analogous to their locations in the periodic table. The correlation is obvious; electronegativity is connected through the higher rows and the rightmost columns.

The location of hydrogen on this scale reflects several of the important chemical properties of this element. Even though it acts like a metallic element in many respects (forming a positive ion, for example), it can as well shape hydride-ion (H-) solids via the more electropositive elements, and of course its ability to share electrons through carbon and other p-block elements provides increase to an extremely rich chemistry, including of course the millions of organic compounds.

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