CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES:
Nomenclature of Elements with atomic number greater than 100 |
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Elements with high atomic numbers are extremely unstable and hard to obtain that highly sophisticated and costly instruments are required for study of such elements. IUPAC has recommended rules for nomenclature of element whose discovery is not completely proven and name not officially recognized. The nomenclature includes names, which have abbreviation, assigned for each digit of the atomic number followed by the word “ium” at the end. When the existence of the new element is proven, permanent name and symbol are given by IUPAC representatives from each country.
Notation for IUPAC nomenclature is shown in table 1. Some examples for naming elements with atomic number above 100 are given in table 2. |
Electronic configurations of elements and the periodic Table
The distribution of electrons into orbitals of an atom is called electronic configuration.
Since periodic table contains periods and groups, we will study electronic configuration of periods and groups separately.
a. Electronic configurations in Periods :
Periods in the periodic table are arranged in such a way that each period reflects the valence shell that is being filled, the orbital in which the electrons are filled, and the number of electrons which is filled in each orbital. For example, for Hydrogen which has electronic configuration of (1s1), 1 means that 1st orbital is being filled, s tells us that s-orbital is being filled, and superscript 1 means that one electron is already filled in s-orbital, which can hold maximum of 2 electrons. Therefore, first period (n=1) starts with Hydrogen (1s1) and ends with Helium (1s2). The second period (n=2) starts with lithium (2s1; it is 2s1 because 1s orbital hold maximum of 2 electrons and therefore the third electron should be filled in the next shell which is 2s orbital) and ends with neon (2s22p6).The third period (n = 3) begins at sodium, and the added electron enters a 3s orbital.
Since periodic table contains periods and groups, we will study electronic configuration of periods and groups separately.
a. Electronic configurations in Periods :
Periods in the periodic table are arranged in such a way that each period reflects the valence shell that is being filled, the orbital in which the electrons are filled, and the number of electrons which is filled in each orbital. For example, for Hydrogen which has electronic configuration of (1s1), 1 means that 1st orbital is being filled, s tells us that s-orbital is being filled, and superscript 1 means that one electron is already filled in s-orbital, which can hold maximum of 2 electrons. Therefore, first period (n=1) starts with Hydrogen (1s1) and ends with Helium (1s2). The second period (n=2) starts with lithium (2s1; it is 2s1 because 1s orbital hold maximum of 2 electrons and therefore the third electron should be filled in the next shell which is 2s orbital) and ends with neon (2s22p6).The third period (n = 3) begins at sodium, and the added electron enters a 3s orbital.
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Successive filling of 3s and 3p orbitals gives rise to the third period of 8 elements from sodium to argon. The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s orbital. Now you may note that before the 4p orbital is filled, filling up of 3d orbitals becomes energetically favorable and we come across the so called 3d transition series of elements. This starts from scandium (Z = 21) which has the electronic configuration 3d14s2. The 3d orbitals are filled at zinc (Z=30) with electronic configuration 3d104s2. The fourth period ends at krypton with the filling up of the 4p orbitals. Altogether we have 18 elements in this fourth period. The fifth period (n = 5) beginning with rubidium is similar to the fourth period and contains the 4d transition series starting at yttrium (Z = 39). This period ends at xenon with the filling up of the 5p orbitals. The sixth period (n = 6) contains 32 elements and successive electrons enter 6s, 4f, 5d and 6p orbitals,
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in the order — filling up of the 4f orbitals begins with cerium (Z = 58) and ends at lutetium (Z = 71) to give the 4f-inner transition series which is called the lanthanoid series. The seventh period (n = 7) is similar to the sixth period with the successive filling up of the 7s, 5f, 6d and 7p orbitals and includes most of the man-made radioactive elements. This period will end at the element with atomic number 118 which would belong to the noble gas family. Filling up of the 5f orbitals after actinium (Z = 89) gives the 5f-inner transition series known as the actinoid series. The 4f- and 5f-inner transition series of elements are placed separately in the Periodic Table to maintain its structure and to preserve the principle of classification by keeping elements with similar properties in a single column.
b. Electronic configurations in groups:
- Elements in the same group (vertical column) have similar valence electronic configuration, and have similar chemical and physical properties.
- For example, valence shell electronic configuration for Group 1 elements (alkali metals) is shown below in table 3.
- Based on the type of valence atomic orbitals which are being filled, periodic table can be divided into four blocks, which are s-block, p-block, d-block, and f-block.
Electronic configurations and types of elements: s-, p-, d-, f- blocks
- Based on the type of valence atomic orbitals which are being filled, periodic table can be divided into four blocks, which are s-block, p-block, d-block, and f-block. Four blocks (s, p, d, f) are shown in figure 2.
(1) s-block:
- Elements of group 1 (alkali) and group 2 (alkali earth metals) which have valence shell electronic configuration of the form ns^1 and ns^2 belong to s-block elements.
(2) p-block:
- Elements of group 13 to group 18 which have valence shell electronic configuration that varies from ns2np1 to ns2np6 respectively belong to s-block elements.
- Group 18 (noble gas family) have ns2np6 valence shell configuration, and elements in this group are low in chemical reactivity because of stable electronic configuration.
- Group 16 and group 17 are called chalcogens and halogens respectively.
(3) d-block:
- Elements of group 3 to group 17 which have general valence shell electronic configuration as (n-1)d^(1-10)ns^(0-2) belong to d-block elements. They are called transitional elements.
(4) f-block:
- Two rows of elements at the bottom of the periodic table are collectively called f-block elements. First row, which has valence shell electronic configuration of the form (n-2)f1-14(n-1)d0-1ns2, is called Lanthanoids. Second row, which has valence shell electronic configuration of the form (n-2)f^(1-14)(n-1)d^(0-2)ns^2, is called Actinoids.
- Actinoid elements are radioactive.
Metals, Non-metals, and Metalloids
Elements in the periodic table can also be divided into metals, non-metals, and metalloids as shown in figure 2 given above.
Metals, which appear on the left side of the periodic table and non-metals, which appear on the top right side of the periodic table, are divided by semi-metals as shown in figure 2 given above.
Differences between properties of metals and non-metals are given in table 4 give below:
Metals, which appear on the left side of the periodic table and non-metals, which appear on the top right side of the periodic table, are divided by semi-metals as shown in figure 2 given above.
- Elements that has properties of both metals and non-metals are called Semi-metals or Metalloids.
- Metallic character decreases and non-metallic character increases as we move from left to right on periodic table.
- Metallic character increases as we move down in a group.
Differences between properties of metals and non-metals are given in table 4 give below:
Physical properties of elements in the periodic table
Each element in the periodic table has its own atomic and ionic radius, ionization enthalpy, electron gain enthalpy and electronegativity due to which elements have its own chemical and physical properties.
a. Atomic Radius:
- In the periodic table, atomic radii of elements decrease as we move from left to right in a period. This is because, with increase in atomic number across a period, electrons are added in the same valence shell due to which effective nuclear charge increases resulting in increased attraction of electrons to the nucleus.
- Figure 3 given below shows the variation of atomic radius with atomic number across the second period.
- Similarly, atomic radii of elements increase as we move down a group. This is because, with increase in atomic number down a group, the principal quantum number (n) also increases due to which the valence electrons are further away from the nucleus. The valence electrons are further away from the nucleus as we move down a group because of shielding effect provided by completely filled inner orbitals (shells) which reduces the attraction between the nucleus and valence electrons.
- Figure 4 given below shows the variation of atomic number with increase in atomic number for group 1 (alkali) and group 17 (halogen) metals.
b. Ionic radius:
- Removal of electrons from the parent atom will lead to the formation of ion called anion, while addition of electrons to the parent atom will lead to the formation of ion called cation. As cations and anions are formed, their atomic radii also change.
- An anion has larger ionic radius than its parent atom because addition of electron which result in increased repulsion among the electrons and decreased effective nuclear charge.
- A cation has smaller size than its parent atom because removal of electron will increase the attraction between the nucleus and electrons resulting in increased effective nuclear charge.
Isoelectronic species:
Atoms and ions which contain the same number of electrons are called isoelectronic species.
Example: O^2-, F^1-, Na^1+, and Mg^2+ are isoelectronic species.
Question: Arrange the above isoelectronic species based on increasing ionic radius.
Answer: Mg^2+ < Na^1+ < F^1- < O^2-
Example: O^2-, F^1-, Na^1+, and Mg^2+ are isoelectronic species.
Question: Arrange the above isoelectronic species based on increasing ionic radius.
Answer: Mg^2+ < Na^1+ < F^1- < O^2-
c. Ionization enthalpy:
- Energy required to remove an electrons from an element in its ground state is called ionization enthalpy. Ionization enthalpy is expressed in units of KJ mol-1 and has positive value at all times because energy is always required to remove and electrons from an atom.
- Figure 5 shows the variation of first ionization enthalpies with atomic number for elements from Z=1 to Z=60.
- Energy required to remove first electrons from an atom is called first ionization enthalpy or ionization energy. Energy required to remove second electron and third electron from an atom is called second ionization enthalpy and third ionization enthalpy respectively.
- The first ionization enthalpy generally increases as we go across a period because when atomic number increases, electrons are added to the same valence shell of elements in a period leading to increase in effective nuclear charge. Because of this increased effective nuclear charge, electrons in atoms with higher atomic numbers are more firmly attracted to the nucleus and thus, require more ionization enthalpy. Figure 6 shows graph of first ionization enthalpies of the second period as a function of atomic number.
- The first ionization enthalpy decreases as we go down in a group because the valence electrons are further away from the nucleus due to shielding of the valence electrons by completely filled orbitals. Figure 7 shows graph of first ionization enthalpies of the first group (alkali metals) as a function of atomic number.
Question: Why is the first ionization enthalpy of boron (Z=5) lower than the first ionization enthalpy of beryllium (Z=4).
Ans: The first ionization enthalpy of boron (Z=5) is lower than the first ionization enthalpy of beryllium (Z=4) because removal of electron from
2p-orbital in B (1s^22s^22p^1) is easier than the removal of electron from 2s-orbital in Be (1s^22s^2). In Boron, the electron in 2p-orbital is shielded by fully filled 1s and 2s-orbitals. Therefore, valence electron of boron in 2p-orbital is further away from the nucleus as compared to the valence electron in
2s-orbital of beryllium.
Ans: The first ionization enthalpy of boron (Z=5) is lower than the first ionization enthalpy of beryllium (Z=4) because removal of electron from
2p-orbital in B (1s^22s^22p^1) is easier than the removal of electron from 2s-orbital in Be (1s^22s^2). In Boron, the electron in 2p-orbital is shielded by fully filled 1s and 2s-orbitals. Therefore, valence electron of boron in 2p-orbital is further away from the nucleus as compared to the valence electron in
2s-orbital of beryllium.
Question: Why is the first ionization enthalpy of nitrogen (Z=7) higher than the first ionization enthalpy of oxygen (Z=8).
Ans: The first ionization enthalpy of nitrogen (Z=7) is higher than the first ionization enthalpy of oxygen (Z=8) because nitrogen has valence shell electronic electronic configuration as 2s^22p^3 which is more stable than valence shell electronic electronic configuration of oxygen which is 2s^22p^4. In oxygen, two electrons are accommodated in one of the p-suborbitals resulting in more electron-electron repulsion. Because of this, first electron can be removed more easily in oxygen as compared to nitrogen.
Ans: The first ionization enthalpy of nitrogen (Z=7) is higher than the first ionization enthalpy of oxygen (Z=8) because nitrogen has valence shell electronic electronic configuration as 2s^22p^3 which is more stable than valence shell electronic electronic configuration of oxygen which is 2s^22p^4. In oxygen, two electrons are accommodated in one of the p-suborbitals resulting in more electron-electron repulsion. Because of this, first electron can be removed more easily in oxygen as compared to nitrogen.
d. Electron Gain Enthalpy:
- The change in enthalpy when an electron is added to a neutral gaseous atom is called Electron Gain Enthalpy. Unlike ionization energy, electron gain enthalpy could be positive (exothermic) or negative (endothermic).
- Electron gain enthalpy generally increases (becomes more negative) as we move from left to right across a period because when we move from left to right, atomic size decreases (effective nuclear charge increases) and it becomes easier to add electron leading to release of more energy.
- Electron gain enthalpy generally decreases (becomes more positive) as we move down a group because of addition of more shell (increasing principle quantum number (n)) making it more difficult to add electron to an atom.
e. Electronegativity:
- The tendency of an atom in a chemical compound to attract shared electrons to itself is called electronegativity.
- Electronegativity of elements generally increases as we move from left to right across a period due to increase in effective nuclear charge.
- Electronegativity decreases as we move down a group because of addition of shells, which shield the valence electrons and decrease attraction between between electrons and the nucleus.
Figure 8 shows trend of atomic radius, ionization enthalpy, electron gain enthalpy, electronegativity, metallic, and nonmetallic character.
Periodicity of Valence state shown by second row elements
- Valence state, also called oxidation state, of an element is usually given by the number of electrons in the outer most orbital or the value of eight minus the number of electrons in the outermost orbitals as shown in figure 9. For Example, in OF2, each fluorine atom shares one electron to achieve octet, and one oxygen atom share two electrons with two fluorine atoms to attain octet. Therefore, oxygen has the oxidation state of +2 and each fluorine atom has the oxidation state of -1.
- Some elements, especially transitional metals and actinoids exhibit different oxidation states.
In the periodic table, chemical reactivity of elements is great at the two extremes. Chemical reactivity of elements at the center of periodic table is the least. Therefore, chemical reactivity of elements decrease and then increase as we move across a period. It is because electron gain enthalpy of elements at extreme right is the highest negative (means that a largest amount of energy is given off when an electron is added to the atom), and ionization enthalpy of element at the extreme left is the least (means that least amount of energy is required to remove an electron from the atom).
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