What is the first words commonly spoken by the student if there is a question, "what do you know about chemistry?". Probably some of students will answer periodic table. Chemistry without learning the periodic table is non sense. Here is the newest version of periodic table. (Source : ptable.com).

- The transition metals are the elements in Groups 1B and 3B through 8B, which have incompletely filled  d  sub-shells, or readily produce cations with incompletely filled  d  sub-shells. (These metals are sometimes referred to as the  d -block transition elements.) 

- The lanthanides and actinides are sometimes called  f -block transition elements because they have incompletely filled  f  sub-shells. 

The chemical reactivity of the elements is largely determined by their valence electrons, which are  the outermost electrons. For the representative elements, the valence electrons are those in the highest occupied  n  shell.  All non-valence electrons in an atom  are referred to as core electrons.  But today we are not discuss about that, it would be discussed in Quantum Theory.

Physical Properties

A. The    effective nuclear charge (Z eff )  is the nuclear charge felt by an electron when both the actual nuclear charge (Z) and the repulsive effects (shielding) of the other electrons are taken into account.  The effective nuclear charge increases from left to right a cross a period and top to bottom in a group. Why it can be happened? because the added electron is a valence electron and valence electrons do not shield each other well, the net effect of moving across the period is a greater effective nuclear charge felt by the valence electrons and the valence electrons are added to increasingly large shells as  n  increases, the electrostatic attraction between the nucleus and the valence electrons actually decreases. 

B. Atomic Radius is one-half the distance between the two nuclei in two adjacent metal atoms or in a diatomic molecule. Periodic trends are clearly evident, consider the atomic radii of these elements against their atomic numbers. So it will increases from right to left in a period and top to bottom in a group. 

C.  Ionic radius    is  the radius of a cation or an anion. Ionic radius affects the physical and chemical properties of an ionic compound. If the atom forms an anion, its size (or radius) increases, because the nuclear charge remains the same but the repulsion resulting from the additional electron(s) enlarges the domain of the electron cloud. On the other hand, removing one or more electrons from an atom reduces electron-electron repulsion but the nuclear charge remains the same, so the electron cloud shrinks, and the cation is smaller than the atom, so ionic radius has parallel trends with atomic radius.  For ions derived from elements in different groups, a size comparison is meaningful only if the ions are isoelectronic. For example N3- and F-, which one is larger? Both of them containing 10 electrons, because   N3-   has only seven protons and   F-   has nine, the smaller attraction exerted by the nucleus on the electrons results in a larger   N3-   ion.

D.    Ionization energy    is  the minimum energy (in kJ/mol) required to remove an electron from a gaseous atom in its ground state. In other words, ionization energy is the amount of energy in kilo-joules needed to strip 1 mole of electrons from 1 mole of gaseous atoms. The magnitude of ionization energy is a measure of how “tightly” the electron is held in the atom. The higher the ionization energy, the more difficult it is to remove the electron.  For a many-electron atom, the amount of energy required to remove the first electron from the atom in its ground state is called the first ionization energy, and so on. The trends is it increases from left to right in a period and bottom to up in a group. The patterns also increases from the first and the second ionization energy regarding to the removal of electron subsequently.  

Although the general trend in the periodic table is for fi rst ionization energies to increase from left to right, some irregularities do exist. The first exception occurs between Group 2A and 3A elements in the same period (for example, between Be and B and between Mg and Al). The Group 3A elements have lower fi rst ionization energies than 2A elements because they all have a single electron in the outermost p  sub-shell ( ns  2  np  1) , which is well shielded by the inner electrons and the  ns  2  electrons. Therefore, less energy is needed to remove a single  p  electron than to remove an s  electron from the same principal energy level. The second irregularity occurs between Groups 5A and 6A (for example, between N and O and between P and S). In the Group 5A elements ( ns  2  np  3 ), the  p  electrons are in three separate orbitals according to Hund’s rule. In Group 6A ( ns  2  np  4) , the additional electron must be paired with one of the three  p  electrons. The proximity of two electrons in the same orbital results in greater electrostatic repulsion, which makes it easier to ionize an atom of the Group 6A element, even though the nuclear charge has increased by one unit. Thus, the ionization energies for Group 6A elements are lower than those for Group 5A elements in the same period.

E. Electron Affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. The more positive is the electron affinity of an element, the greater is the affinity of an atom of the element to accept an electron and  a large positive electron affinity means that the negative ion is very stable (that is, the atom has a great tendency to accept an electron), just as a high ionization energy of an atom means that the electron in the atom is very stable. Electron affinity is positive if the reaction is exothermic and negative if the reaction is endothermic. Electron affinity is positive if the reaction is exothermic and negative if the reaction is endothermic.  

Chemical Properties

Diagonal relationships are similarities between pairs of  elements in different groups and periods of the periodic table.  

Hydrogen,  there is no totally suitable position for hydrogen in the periodic table, even sometimes it is placed in group 1A, but it really has own class. Hydrogen can be a uni-positive ion (hydrated in solution) and uni-negative ion (in ionic compound e.g NaH).

Group 1A elements, also called as alkali metal. They all have tendency to release one electron to form uni-positive ion and have low ionization energies. These elements are so reactive, as you can see in this video : 

When exposed to air, they gradually lose their shiny appearance as they combine with oxygen gas to form oxides. The other alkali metals all form oxides and  peroxides.  Potassium, rubidium, and cesium also form  super-oxides.  The reason that different types of oxides are formed when alkali metals react with oxygen has to do with the stability of the oxides in the solid state. Because these oxides are all ionic compounds, their stability depends on how strongly the cations and anions attract one another.  

Group 2A elements, alkali earth metals are less reactive than alkali metals. They have tendency to release two electrons to form 2+ ions, hence the metallic character increases from top to bottom.  The reactivities of alkaline earth metals with water vary quite markedly. Beryllium does not react with water; magnesium reacts slowly with steam; calcium, strontium, and barium are reactive enough to attack cold water.  The reactivities of the alkaline earth metals toward oxygen also increase from Be to Ba.  Beryllium and magnesium form oxides (BeO and MgO) only at elevated temperatures, whereas CaO, SrO, and BaO form at room temperature. They react with aqueous acid solutions to produce hydrogen gas. However, because these metals also attack water, two different reactions will occur simultaneously.   

(Strontium-90, a radioactive isotope, is a major product of an atomic bomb explosion. If an atomic bomb is exploded in the atmosphere, the strontium-90 formed will eventually settle on land and water, and it will reach our bodies via a relatively short food chain. For example, if cows eat contaminated grass and drink contaminated water, they will pass along strontium-90 in their milk. Because calcium and strontium are chemically similar,  Sr2+ ions can replace   Ca2+ ions in our bones. Constant exposure of the body to the high-energy radiation emitted by the strontium-90 isotopes can lead to anemia, leukemia, and other chronic illnesses). 

Group 3A elements, the first element if this group is boron, it is a metalloid; the rest are metals. Boron does not form binary ionic compounds and is nonreactive toward oxygen gas and water.  The next element, aluminum, readily forms aluminum oxide when exposed to air, which is less reactive than the elemental aluminum.  The other Group 3A metallic elements form both uni-positive and tri-positive ions. Moving down the group, we find that the uni-positive ion becomes more stable than the tri-positive ion. 

Group 4A elements,  The first member of Group 4A, carbon, is a nonmetal, and the next two members, silicon and germanium, are metalloids. The metallic elements of this group, tin and lead, do not react with water, but they do react with acids (hydrochloric acid, for example) to liberate hydrogen gas.  The Group 4A elements form compounds in both the +2 and +4   oxidation states,  as we move down the group, however, the trend in stability is reversed. 

Group 5A elements, nitrogen and phosphorus are nonmetals, arsenic and antimony are metalloids, and bismuth is a metal.  Elemental nitrogen is a diatomic gas (N2 ). It forms a number of oxides (NO, N2O, NO2 , N2O4 , and N2O5 ), of which only N2 O5  is a solid; the others are gases. Nitrogen has a tendency to accept three electrons to form the nitride ion. Phosphorus exists as P4  molecules. A rsenic, antimony, and bismuth have extensive three-dimensional structures. Bismuth is a far less reactive metal than those in the preceding groups. 

Group 6A elements,  he first three members of Group 6A (oxygen, sulfur, and selenium) are nonmetals, and the last two (tellurium and polonium) are metalloids. Oxygen and Sulfur are very famous, oxygen is diatomic gas that we inhale. Sulfuric acid is formed from sulfur trioxide reacts with water. Polonium as the last member, is a radioactive element that is difficult to study in the laboratory. Sulfur, selenium, and tellurium also form dinegative anions. 

Group 7A elements,  all the halogens are nonmetals with the general formula X2,  where X denotes a halogen element. Because of their great reactivity, the halogens are never found in the elemental form in nature. (The last member of Group 7A, astatine, is a radioactive element. Fluorine is so reactive that it attacks water to generate oxygen.  The halogens have high ionization energies and large positive electron affinities. Anions derived from the halogens (F2,Cl2,Br2,andI2  ) are called  halides . They are iso-electronic with the noble gases immediately to their right in the periodic table. The halogens also form many molecular compounds among themselves and with nonmetallic elements in other groups. The halogens react with hydrogen to form hydrogen halides. When this reaction involves fluorine, it is explosive, but it becomes less and less violent as we substitute chlorine, bromine, and iodine. The hydrogen halides dissolve in water to form hydrohalic acids. Hydrofluoric acid (HF) is a weak acid (that is, it is a weak electrolyte), but the other hydrohalic acids (HCl, HBr, and HI) are all strong acids (strong electrolytes). 

Group 8A elements,  ll noble gases exist as monatomic species. The Group 8A ionization energies are among the highest of all elements, and these gases have no tendency to accept extra electrons. For years they are known as inert gas until Neil Bartlett exposed xenon to platinum hexafluoride, a strong oxidizing agent. Since then,  a number of xenon compounds (XeF4 , XeO3 , XeO4 , XeOF4 ) and a few krypton compounds (KrF2 , for example) have been prepared.  No compounds of helium and neon are known.  

Published by Galuh Catur WP, on Monday, August 6, 2019.

Reference : Chang, R. (2010). Chemistry 10th edition ( Chapter 8 page 322-355 ). New York: McGraw-Hill.

Whether you have a test coming up or just want to learn something new, the periodic table of elements is a helpful tool to know. Memorizing all 118 elements may seem tricky, especially since each one has a unique symbol and atomic number. Fortunately, if you start early, you can learn a few elements every day. Mnemonic devices, phrases, and pictures will boost your memory while making studying enjoyable. If you’re ready to test your skills, try a few games or even draw a table completely from memory. 

1. Learn a few elements a day. Start with the first ten. Once you have mastered those, add in another ten. Keep reviewing the old elements even as you learn new ones. Start studying early so that you have time to memorize all 118 elements. 

2. Print out a copy of the periodic table. Wherever you go, it will go with you. It's advisable to print out more than one copy. Keep one on your desk, one in your backpack or purse, and one wherever else you might go. You can also use a digital version on a phone or tablet. 

3. Make flashcards for each element. On one side, put the element symbol, such as Ag, S, or Cu, as well as the atomic number. On the other side, put the full name of the element, such as Silver, Sulphur, or Copper. Use the cards to test yourself.

4. Break down the table into smaller sections. You could go by row, column, atomic weight, groups, or blocks.

5. Quiz yourself during your breaks and free time. Instead of cramming for several hours, try studying whenever you have a few minutes to spare. This may be on the bus, during lunch, or while you’re in line.

6. Using Mnemonic Devices.  Some teacher may teach you by using donkey bridge for those elements. 

7. Periodic Table Song.

8. Playing Periodic Table Games 

Here some videos about memorizing periodic table in easiest and fun way. Check it out !