The periodic table of elements, by far, stands among the most notable discoveries in the history of chemistry. Nowadays, probably everyone who has ever heard about chemistry heard about the periodic table of elements. Its implications on the scientific world are so enormous that it has even leaked out to the general public and have become associated with the chemistry itself. This is not surprising considering the scientific strength of this instrument. Using the periodic table, many characteristic properties of the elements can be predicted and many phenomena can be explained. This discovery put so much scientific knowledge together, organized and systematized that it can deservedly be called one of the greatest discoveries in chemistry. As a matter of fact, many scientists contributed to the emergence and development of the table, but Dmitri Mendeleev is credited as its inventor. In reality, he was not alone to make advancements in classifying chemical elements and putting them together in the form of the table, but he made the revolutionizing discovery, the one that other chemists failed to notice. The periodicity of properties of the chemical elements has been noticed before Mendeleev, however, he was the first to leave the gaps in it to account for the unknown elements. Another important distinction of his table was that he placed elements according to their chemical families sometimes disregarding the atomic weights (tellurium and iodine). Mendeleev’s table was published in 1869 and predicted the existence of two elements which were not known at his time (gallium and germanium) and their atomic weights (Scerri 110). He also made some corrections to the known data, and all of his predictions turned out to be true. Since Mendeleev's time, a lot of scientific discoveries were made, and many new trends in the periodic table were found which I will further focus on.
The trends in the periodic table of elements which consider the chemical properties of the elements are among the strongest and most studied. One of such trends is the electronegativity trend. Electronegativity can be defined as a chemical property of an element describing its tendency to attract electrons. The higher is electronegativity; the stronger is the atom's ability to attract electrons. Although there is no standardized method of measuring electronegativity, the most common and accepted is the scale proposed by Linus Pauling. The periodic table with electronegativity values according to this scale is presented in Fig.1.
Figure 1. Periodic Table with Electronegativity
The electronegativity of the known elements ranges from 0.7 to 3.98 with cesium having the lowest value and fluorine having the highest value among known and studied chemical elements. The electronegativity itself exists because of the specific electron configurations of the atoms. Following the octet rule, when forming chemical bonds, atoms strive to achieve the perfect combination of 8 electrons on the outer layer. The elements which are situated at the right of the table and have more than half of the outer layer filled tend to gain electrons and finish their valence shell. Meanwhile, the elements at the left of the table tend to lose the electrons and gain the electron structure of the noble gas of the previous period. The general trends in the periodic table are that across the period from left to right, electronegativity increases. If the outer electron layer is more than half-full, the element is more energy-efficient in gaining the electron than in losing it. The converse is also true. In the groups from the top to the bottom, the electronegativity decreases. Electronegativity depends on the distance between the charged nucleus and the electrons in the valence shell. And when we go down the group, the atomic number increases, which means that the atomic radius also increases thus decreasing electronegativity. There are, however, some important exceptions to this trend. Obviously, noble gasses do not have a higher electronegativity than halogens (in fact, they do not have it at all) since they possess complete valence shell and normally neither gain nor lose the electrons. Lanthanoids and Actinoids also do not fall into this trend because they have more complex chemistry, and actually their valence shells are equally filled. Electronegativity is a rather important property of an atom which shows itself when the atoms interact to form a compound. The difference in electronegativity between the atoms in a molecule defines the nature of a bond between them. If the difference in electronegativity is lower than 0.4, then a nonpolar bond is formed. If its value lies between 0.4 and 1.7, then the bond is classified as polar. And finally, if the difference is higher than 1.7 then the bond between them is ionic. The nature of the bond between the two elements defines both chemical and physical properties of the substance. The larger is the difference between the elements’ electronegativities, the more likely will a reaction between them occur. For example, the reaction of fluorine with hydrogen is violent under normal conditions and is noticeable even at temperatures below 0 °C:
H2 + F2 → 2HF
while the reaction between carbon and boron does not occur at normal conditions and only starts to be noticeable beyond 1000 °C:
2B + N2 → 2BN.
The same tendency applies within a group. With the increase in the atomic number in the group, the electronegativity drops, and so drops the strength of the bond between the atoms. For example, in the row of hydrohalides, the lowest possible formation temperature increases from the fluorine to the iodine. The general reaction can be expressed as follows:
Hal2 + H2 → 2HHal
The fluorine reacts violently at room temperature while iodine shows significant reaction speeds only when irradiated with light. The same pattern is exhibited by the redox properties of these hydrohalides. The hydrogen iodide is a strong reducing agent which is easily oxidized even by the air oxygen while the hydrogen fluoride cannot even be considered a reducing agent due to fluoride's extremely strong oxidizing power.
The next property which has a distinct trend in the periodic table of elements is the ionization energy. Ionization energy is the energy needed to remove one electron from a neutral atom in a gaseous phase. The higher is this energy, the more atom is likely to lose an electron and become a cation. Obviously, as it gets higher, the atom is less likely to give its electron and become a cation. Generally, the ionization energy trend follows the one with electronegativity – it increases within a period from left to right and decreases from top to bottom within a group. This happens due to two phenomena which are called electron shielding and valence shell stability. The former is exhibited when we go from top to bottom within a group – as the number of electron layers between the nucleus and the valence shell increases, each additional layer shields the valence shell from the nucleus and weakens the attraction between them. Valence shell stability is exhibited in the left to right movement within a period, the extra electrons in the valence shell stabilize the atom and make its radius smaller which strengthens the attraction between the nucleus and the electrons in the valence shell. The only difference between the ionization energy and electronegativity trends is that noble gasses have the ionization energy value, and it is very high. The ionization energy trend is well illustrated by the reaction of the alkali metals with water. This reaction in general from for all these metals can be written as
2Me + 2H2O → 2MeOH + H2
In the sequence from lithium to cesium, the speed of the reaction increases showing that each consequent metal loses its electron more easily. Lithium reacts with a moderate hissing and slowly dissolves in water, sodium reacts faster, potassium already starts to burn while cesium reacts with a violent explosion.
Another important trend in the periodic table is electron affinity. Electron affinity is the atom’s tendency to accept an electron. It may sound very alike to the electronegativity, but it has a quantitative expression and is measured as the energy change in a neutral atom in the gas phase when an electron is added to it. The electron affinity increases in the period from left to right, however, the up-down trend is not so straightforward as in electronegativity case. Fluorine, for example, has lower electron affinity than chlorine (Ramireddy, Zheng, and Nguyen).
The metallic character also has a trend in the periodic table of elements. It is by far the most tangible of the trends. This trend is generally determined by the two former trends as metals readily lose their electrons while nonmetals tend to gain them. The trend is caused by the change in atomic radius and the attraction between the nucleus and electrons in the valence shell. As the size increases, the attraction becomes weaker, and the atom more readily loses its electron and vice versa. Here, the general trends follow the pattern for the electronegativity but work in an opposite direction. The metallic character of an element increases in the right to left direction within the period and from top to bottom making fluorine the most nonmetallic and cesium most metallic. This defines the acid-base properties of the compounds the elements form. The more metallic character the element has, the more basic are its compounds. The nonmetals, in contrast, form acids. The elements which are situated in between, the ones having half-full valence shell, exhibit amphoteric properties. Alkali metals replace hydrogen in water and form lye while nonmetals form acids if their oxides react with water. Typical representatives of basic hydroxides are NaOH, KOH, BaOH while the typical acids are H2SO4 and HNO3. The reactions of acids and bases are called neutralization reactions. Amphoteric elements exhibit the properties of both acids and bases and react with both of them:
Al(OH)3 + 3HCl → AlCl3 + 3H2O
Al(OH)3 + NaOH → NaAlO2 + 2H2O
Accordingly, as the metallic properties increase, so do the basic properties of the element’s compounds. With the increase of nonmetallic properties, the acidic properties also rise.
In fact, the trend defining the metallic character and chemical properties of the elements is the structure of their valence shell. This is defined by two factors – the group and period in which the element resides. The group defines how many electrons an atom has on the outer shell and the period determines to which extent is the character of the element metallic. Although all of the previously discussed factors take part in determining chemical properties of a chemical element and trends between the elements in the table, the number of valent electrons is the hard limitation of these properties. For example, alkali metals having only one electron in the valence shell can only lose their electron and take the oxidation state +1. The more valent electrons the element has in the outer layer, the more flexible it is in its properties. Nitrogen, for example, can either lose or gain electrons, taking oxidation levels -3, 0, +1, +2, +3. +4, +5. Chlorine having seven electrons has even more possible oxidation states, however, the most stable still is -1. The metallic character determined by the period also defines the element’s chemical activity. For all the elements beside the alkali and alkaline earth metals, the elements having greater period number are less active. For example, H2SO4 is a highly stable and strong acid decomposing at 700 °C, selenic acid H2SeO4 decomposes at 200 °C (De 543-545), and telluric acid H2TeO4 is not known.
One of the notable trends mentioned in the previous paragraphs is the atomic radius change trend. Atomic radius is half distance from the nucleus to the outermost electron orbital of the atom. Another definition specifies it as the half of the distance between two single bound nuclei in a diatomic compound. The trend in atomic radii is counterintuitive – the radii increase from right to left in the periods and from top to bottom in the groups. From the first sight, it is likely to seem that with the increase in electrons and protons in an atom (when going from left to right within a period) the radius of the atom should at least stay the same if not increase. However, the opposite actually happens – the atomic radius of sodium is greater than the atomic radius of argon. This is because within the same period the electrons are added to the same shell, and at the same time, protons are added to the nucleus. The additional protons create stronger attractional force with the electrons. As the valence shell electrons are pulled closer to the nucleus, the radius of the atom gets smaller with each additional electron until the valence shell is not completely filled. Accordingly, the noble gasses have the smallest atom radius in their respective period, and the alkali metals have the largest atom radius. The up-down trend is more easily explainable. As we move down the group, the new electron shells are added to the atom, and each new electron shell is significantly larger than the previous valence shell. Although, the horizontal and vertical shift in atomic radius are not equal. The largest difference occurs between the elements of the first and second group which is generally larger than the between period difference.
The most basic trend in the periodic table is the trend of electron configuration of the chemical elements. In its first from, Dmitri Mendeleev arranged the chemical elements in order of their atomic masses with some adjustments for the chemical properties. He switched places iodine and tellurium, to place them in their similar chemical families. He did not realize the significance of his arrangement but later, with further discoveries of the subatomic particles, the bigger picture became much clearer. Indeed, as Mendeleev made his mind, there appeared to be something more important for the order of the elements than the atomic mass. These were the atomic numbers. Atomic number is the numbers of protons in the nucleus of an atom of a given chemical element. And as the nucleus builds up, the structure of the electron shell grows accordingly. The structure of electron shell changes cyclically. Each new period adds another layer to the atom structure. In the same layer, electrons build up by consecutively filling vacant orbitals. The first orbital is the s orbital which can host two electrons. The second is the p orbital with six possible places for electrons. These two orbitals together from the possible octet of electrons (noble gasses). The larger d and f orbitals are filled much later. For example, 3d sublevel is filled after 4s and 4f only after 6s (Setti and Satake 28). The whole structure of the periodic table resides upon the repeatability of the electron shell pattern.
The melting point and boiling point are rather tangible and practical properties for us, however, in the periodic table there is no such hard, determined pattern for these characteristics as for the electronegativity or metallic character. The general tendency, if it may be called so, suggests that that the melting point of the elements is the highest in the middle of the period and drops towards its ends. In the groups, from top to bottom, the melting point generally gets higher. Although, there are a lot of outliers in this tendency. For example, the elements of the first two groups have a lower melting point with the period number increase; the same is true about the sixth group. The most notable elements from the table are the tungsten – the simple substance having the highest melting point under atmospheric pressure (graphite actually has a bit higher melting point, but it melts only under pressure, under regular conditions it sublimes), and two simple substances under normal conditions – bromine, and mercury. The boiling points of the elements mainly repeat the distribution of the melting points with some exceptions. The highest temperature boiling substance is neither graphite nor tungsten, it is rhenium (5590 °C).(“Periodic Table: Trends”)
The periodic table of elements is by far one of the most important consolidations of our knowledge of the external world. This table unites all the elements in a systematic way providing explanation and understanding for the structure and interactions of the building blocks of our universe. A number of trends about the elements in this table were researched, found, and explained. Such trends include the electronegativity, energy of ionization, electron affinity, metallic character, atomic radius, electron structure while such properties as melting and boiling points do not have the exact representation in the table. The periodic table explains the chemical properties of the elements and their compounds as well as serves as a strong prognostic utility. It is clear that this discovery brought up by Dmitri Mendeleev will continue to play its huge part in the overall scientific and specifically chemical science progress.
References
De, Anil K. A Textbook Of Inorganic Chemistry. New Delhi: Wiley Eastern, 1983. Print.
"Periodic Table: Trends". Rsc.org. N.p., 2016. Web. 24 May 2016.
Ramireddy, Swetha, Bingyao Zheng, and Emily Nguyen. "Periodic Trends". Chemwiki.ucdavis.edu. N.p., 2013. Web. 24 May 2016.
Scerri, Eric R. The Periodic Table. Oxford: Oxford University Press, 2007. Print.
Sethi, M. S and M Satake. Periodic Table & Periodic Properties. New Delhi: Discovery Publishing House, 1992. Print.
