The periodic table is a tabular representation of all the chemical elements, arranged in order of increasing atomic number, with their corresponding electron configurations and properties. The modern periodic table is organized into rows (periods) and columns (groups) based on the electronic structure of the elements. The arrangement of the periodic table allows for the prediction of the chemical and physical properties of the elements based on their position in the table. This article will provide an in-depth look at the periodic table and the periodic properties of the elements.
History of the Periodic Table
The periodic table has a long and interesting history. The concept of periodicity, or the recurring nature of certain properties of the elements, was first recognized by early chemists such as John Dalton and Johann Wolfgang Döbereiner. However, the first attempt at organizing the elements into a periodic table was made by Dmitri Mendeleev in 1869.
Mendeleev arranged the elements in order of increasing atomic weight and observed that certain properties, such as valence and electronegativity, repeated themselves periodically. He also left gaps in the table for elements that were yet to be discovered, and predicted their properties based on the properties of the surrounding elements.
Mendeleev’s periodic table was successful in predicting the properties of previously unknown elements, such as gallium and germanium, which were discovered shortly after his publication. However, his periodic table had some inconsistencies, such as the placement of hydrogen, and the fact that some elements appeared to be out of place based on their chemical properties.
The modern periodic table, which is based on the electronic structure of the elements, was developed independently by several scientists, including Julius Lothar Meyer and Dmitri Mendeleev himself. The modern periodic table was first proposed by Henry Moseley in 1913, who arranged the elements in order of increasing atomic number, which provided a more accurate reflection of the electronic structure of the elements.
Quantum Numbers
Quantum numbers are used to describe the electron configuration of an atom. They are used to specify the energy, orbital shape, and orientation of electrons in an atom. There are four types of quantum numbers that are used to describe the electronic configuration of an atom:
Principal Quantum Number (n)
The principal quantum number, denoted by ‘n’, describes the energy level or shell in which an electron is present. The value of n can be any positive integer starting from 1. The energy of an electron in an atom increases as the value of n increases. Each energy level or shell can accommodate a certain maximum number of electrons, given by the formula 2n².
Azimuthal Quantum Number (l)
The azimuthal quantum number, denoted by ‘l’, describes the shape of the orbital in which an electron is present. The value of l can be any integer between 0 and (n-1). Each value of l corresponds to a specific shape of the orbital. The shape of the orbital is determined by the angular momentum of the electron.
Magnetic Quantum Number (m)
The magnetic quantum number, denoted by ‘m’, describes the orientation of the orbital in which an electron is present. The value of m can be any integer between -l and +l, including 0. Each value of m corresponds to a specific orientation of the orbital in space.
Spin Quantum Number (s)
The spin quantum number, denoted by ‘s’, describes the intrinsic spin of the electron. The value of s can be either +1/2 or –1/2. The spin of an electron is a fundamental property of the electron and cannot be changed.
These four quantum numbers together determine the electronic configuration of an atom.
The Pauli Exclusion Principle
The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers. This means that each electron in an atom has a unique electronic configuration described by its set of quantum numbers.
The electronic configuration of an atom is often written in shorthand notation, using the following format:
n, l, m, s
where n is the principal quantum number, l is the azimuthal quantum number, and m is magnetic quantum number, and s is the spin quantum number.
For example, the electronic configuration of a hydrogen atom is 1s1, which means that it has one electron in the 1s orbital, with n=1, l=0, and m = 0, s=+1/2.
The electronic configuration of an atom determines its chemical properties and reactivity. The arrangement of electrons in an atom determines the atom’s ability to form chemical bonds and participate in chemical reactions.
Aufbau Principle
The Aufbau principle states that electrons are filled into atomic orbitals in order of increasing energy. This means that the lowest energy orbital is filled first, followed by the next lowest energy orbital, and so on. The order of filling is determined by the value of the principal quantum number (n) and the azimuthal quantum number (l).
Hund’s Rule
Hund’s rule states that when filling degenerate orbitals (orbitals with the same energy), electrons will first occupy different orbitals with the same spin before pairing up. This means that electrons will fill up each orbital in a subshell with the same spin before pairing up in the same orbital.
Calculations from Quantum Number
The position of an element on the periodic table can be determined by its electron configuration, which is determined by the values of its quantum numbers. Specifically, the principal quantum number (n) indicates the energy level or shell that the electron occupies, while the azimuthal quantum number (l) indicates the subshell (s, p, d, f) and the magnetic quantum number (m) indicates the orbital within that subshell.
Here are five examples of elements and their corresponding quantum numbers:
- Carbon (C) – Carbon has six electrons, with the electron configuration of 1s2 2s2 2p2. The principal quantum number for all six electrons is 2, indicating they are in the second energy level. The first two electrons fill the 1s subshell (l=0), and the next two fill the 2s subshell (l=0). The remaining two electrons fill the 2p subshell (l=1), with one electron in the 2px orbital (m=-1), and the other in the 2py orbital (m=0).
- Sodium (Na) – Sodium has 11 electrons, with the electron configuration of 1s2 2s2 2p6 3s1. The first 10 electrons follow the same pattern as carbon, filling the 1s, 2s, and 2p subshells. The 11th electron is in the 3s subshell (l=0) of the third energy level (n=3).
- Chlorine (Cl) – Chlorine has 17 electrons, with the electron configuration of 1s2 2s2 2p6 3s2 3p5. The first 16 electrons follow the same pattern as sodium, filling the 1s, 2s, 2p, and 3s subshells. The remaining five electrons fill the 3p subshell (l=1), with one electron in each of the three 3p orbitals (m=-1, 0, 1).
- Iron (Fe) – Iron has 26 electrons, with the electron configuration of 1s2 2s2 2p6 3s2 3p6 4s2 3d6. The first 18 electrons follow the same pattern as chlorine, filling the 1s, 2s, 2p, 3s, 3p subshells. The remaining eight electrons fill the 4s subshell (l=0) and the 3d subshell (l=2). The 4s subshell is filled before the 3d subshell due to the lower energy of the 4s orbital.
- Gold (Au) – Gold has 79 electrons, with the electron configuration of 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s1. The first 54 electrons follow the same pattern as iron, filling the 1s, 2s, 2p, 3s, 3p, 4s, 3d, and 4p subshells. The remaining 25 electrons fill the 5s subshell (l=0) and the 6s subshell (l=0). The 5s subshell is filled before the 4d subshell due to its lower energy, and the 6s subshell is filled last as it has the highest principal quantum number.
The electronic configuration of an element can provide useful information about its position on the periodic table, including its group and period. Here are some steps to follow when using electronic configuration to deduce the group and period of an element:
Determine the number of energy levels (or shells) occupied by electrons in the element’s electronic configuration. This will correspond to the element’s period number.
Look at the outermost energy level (highest value of n) occupied by electrons in the element’s electronic configuration. The subshell (s, p, d, or f) that contains the highest value of l in that energy level corresponds to the element’s group number.
Read Also: Determining Atomic Number using Spectroscopy
Identifying the Element’s Position on the Periodic Table
Identify the element’s position on the periodic table based on its group and period.
| 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
| 1 | H 1 1.008 |
He 2 4.003 |
||||||||||||||||
| 2 | Li 3 6.941 |
Be 4 9.012 |
B 5 10.81 |
C 6 12.01 |
N 7 14.01 |
O 8 16.00 |
F 9 19.00 |
Ne 10 20.18 |
||||||||||
| 3 | Na 11 22.99 |
Mg 12 24.31 |
Al 13 26.98 |
Si 14 28.09 |
P 15 30.97 |
S 16 32.06 |
Cl 17 35.45 |
Ar 18 39.95 |
||||||||||
| 4 | K 19 39.10 |
Ca 20 40.08 |
Sc 21 44.96 |
Ti 22 47.87 |
V 23 50.94 |
Cr 24 52.00 |
Mn 25 54.94 |
Fe 26 55.85 |
Co 27 58.93 |
Ni 28 58.69 |
Cu 29 63.55 |
Zn 30 65.38 |
Ga 31 69.72 |
Ge 32 72.63 |
As 33 74.92 |
Se 34 78.96 |
Br 35 79.90 |
Kr 36 83.80 |
| 5 | Rb 37 85.47 |
Sr 38 87.62 |
Y 39 88.91 |
Zr 40 91.22 |
Nb 41 92.91 |
Mo 42 95.94 |
Tc 43 (98.00) |
Ru 44 101.07 |
Rh 45 102.91 |
Pd 46 106.42 |
Ag 47 107.87 |
Cd 48 112.41 |
In 49 114.82 |
Sn 50 118.71 |
Sb 51 121.76 |
Te 52 127.60 |
I 53 126.90 |
Xe 54 131.29 |
| 6 | Cs 55 132.91 |
Ba 56 137.33 |
La – Lu 57 – 71 138.91 – 174.97 |
Hf 72 178.49 |
Ta 73 180.95 |
W 74 183.84 |
Re 75 186.21 |
Os 76 190.23 |
Ir 77 192.22 |
Pt 78 195.08 |
Au 79 196.97 |
Hg 80 200.59 |
Tl 81 204.38 |
Pb 82 207.2 |
Bi 83 208.98 |
Po 84 (209.00) |
At 85 (210.00) |
Rn 86 (222.00) |
| 7 | Fr 87 (223.00) |
Ra 88 (226.03) |
Ac – Lr 89 -103 (227.03) – (262.11) |
Rf 104 (267.12) |
Db 105 (270.13) |
Sg 106 (271.13) |
Bh 107 (270.13) |
Hs 108 (277.15) |
Mt 109 (276.15) |
Ds 110 (281.16) |
Rg 111 (280.16) |
Cn 112 (285.17) |
Nh 113 (284.18) |
Fl 114 (289.19) |
Mc 115 (288.19) |
Lv 116 (293.20) |
Ts 117 (294.21) |
Og 118 (294.21) |
| La 57 138.91 |
Ce 58 140.12 |
Pr 59 140.91 |
Nd 60 144.24 |
Pm 61 (145.00) |
Sm 62 150.36 |
Eu 63 151.96 |
Gd 64 157.25 |
Tb 65 158.93 |
Dy 66 162.50 |
Ho 67 164.93 |
Er 68 167.26 |
Tm 69 168.93 |
Yb 70 173.05 |
Lu 71 174.97 |
||||
| Ac* 89 (227.03) |
Th 90 232.04 |
Pa 91 231.04 |
U 92 238.03 |
Np 93 (237.05) |
Pu 94 (244.06) |
Am 95 (243.06) |
Cm 96 (247.07) |
Bk 97 (247.07) |
Cf 98 (251.08) |
Es 99 (252.08) |
Fm 100 (257.10) |
Md 101 (258.10) |
No 102 (259.10) |
Lr 103 (262.11) |
Here are some specific examples:
Sodium (Na): The electronic configuration of sodium is 1s22s2 2p63s1. The highest value of n in this configuration is 3, so sodium is in period 3. The highest value of l in the 3s subshell is 0, so sodium is in group 1 (also known as the alkali metals). Therefore, sodium is located in the first column of period 3 on the periodic table (see the Periodic Table above).
Oxygen (O): The electronic configuration of oxygen is 1s22s2 2p4 . The highest value of n in this configuration is 2, so oxygen is in period 2. The highest value of l in the 2p subshell is 1, so oxygen is in group 16 (also known as the chalcogens). Therefore, oxygen is located in the sixth column of period 2 on the periodic table (see the Periodic Table above).
Iron (Fe): The electronic configuration of iron is 1s22s2 2p63s23p64s23d6. The highest value of n in this configuration is 4, so iron is in period 4. The highest value of l in the 3d subshell is 2, so iron is in group 8 (also known as the iron group). Therefore, iron is located in the eighth column of period 4 on the periodic table (see the Periodic Table above).
By following these steps, we can use the electronic configuration of an element to deduce its group and period on the periodic table.
Short Exercises
- Determine the electron configuration and position on the periodic table for an element with 13 electrons.
Answer: The electron configuration is 1s22s2 2p63s23p1. This element is aluminum (Al) and belongs in period 3 and group 13 (also known as group IIIA).
2. Determine the electron configuration and position on the periodic table for an element with 35 electrons.
Answer: The electron configuration is 1s22s2 2p63s23p64s23d104p5. This element is bromine (Br) and belongs in period 4 and group 17 (also known as group VIIA or the halogens).
3. Determine the electron configuration and position on the periodic table for an element with 56 electrons.
Answer: The electron configuration is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6. This element is barium (Ba) and belongs in period 6 and group 2 (also known as group IIA or the alkaline earth metals).
4. Determine the electron configuration and position on the periodic table for an element with 3 electrons.
Answer: The electron configuration is 1s2 2s1. This element is lithium (Li) and belongs in period 2 and group 1 (also known as group IA or the alkali metals).
5. Determine the electron configuration and position on the periodic table for an element with 92 electrons.
Answer: The electron configuration is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f2. This element is uranium (U) and belongs in period 7 and group 3 (also known as group IIIB).
Read Also: Exciting Atomic Structure and Periodicity
Structure of the Periodic Table
The modern periodic table is arranged in rows and columns, with the rows representing periods and the columns representing groups. Each period corresponds to the filling of a new electron shell, and each group corresponds to the number of valence electrons in the outermost shell.
The periodic table has element that belong to different groups with some of them with their names:
Alkali Metals (Group 1)
Alkali metals are a group of highly reactive metals that occupy the first column of the periodic table. The group includes lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). Alkali metals have 1 valence electron and tend to lose electrons in reactions.
Li (metal, low reactivity) | Na (metal, moderate reactivity) | K (metal, high reactivity) | Rb (metal, high reactivity) | Cs (metal, high reactivity) | Fr (metal, highly radioactive)
Alkaline Earth Metals (Group 2)
Alkaline earth metals are a group of reactive metals that occupy the second column of the periodic table. The group includes beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). Alkaline earth metals have 2 valence electrons and tend to lose electrons in reactions.
Be (metal, low reactivity) | Mg (metal, moderate reactivity) | Ca (metal, high reactivity) | Sr (metal, high reactivity) | Ba (metal, high reactivity) | Ra (metal, highly radioactive)
Transition Metals
Transition metals are a group of metals that occupy the middle of the periodic table.
The group includes:
Scandium (Sc), titanium (Ti), vanadium (V), chromium (Cr), manganese (Mn), iron (Fe), cobalt (Co), nickel (Ni), copper (Cu), zinc (Zn), yttrium (Y), zirconium (Zr), niobium (Nb), molybdenum (Mo), technetium (Tc), ruthenium (Ru), rhodium (Rh), palladium (Pd), silver (Ag), cadmium (Cd), hafnium (Hf), tantalum (Ta), tungsten (W), rhenium (Re), osmium (Os), iridium (Ir), platinum (Pt), gold (Au), mercury (Hg), and rutherfordium (Rf).
Transition metals have 1 or 2 valence electrons and tend to lose electrons in reactions.
Sc (metal, low reactivity) | Ti (metal, low reactivity) | V (metal, moderate reactivity) |
Cr (metal, moderate reactivity) | Mn (metal, moderate reactivity) | Fe (metal, moderate reactivity) |
Co (metal, moderate reactivity) | Ni (metal, moderate reactivity) | Cu (metal, moderate reactivity) |
Zn (metal, low reactivity) | Y (metal, low reactivity) | Zr (metal, low reactivity) | Nb (metal, low reactivity) |
Mo (metal, moderate reactivity) | Tc (metal, radioactive) | Ru (metal, low reactivity) | Rh (metal, low reactivity) | Pd (metal, low reactivity) | Ag (metal, low reactivity) | Cd (metal, moderate reactivity) | Hf (metal, low reactivity) | Ta (metal, low reactivity) | W (metal, moderate reactivity) | Re (metal, moderate reactivity) |
Os (metal, low reactivity) | Ir (metal, low reactivity) | Pt (metal, low reactivity) | Au (metal, low reactivity) |
Hg (metal, moderate reactivity) | Rf (metal, radioactive).
Pnictogens (Group 15)
Pnictogens are a group of nonmetals and metalloids that occupy the fifteenth column of the periodic table. The group includes nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). Pnictogens have 5 valence electrons and tend to gain or share electrons in reactions.
N (nonmetal, low reactivity) | P (nonmetal, low reactivity) | As (metalloid, moderate reactivity) | Sb (metalloid, moderate reactivity) | Bi (metal, low reactivity)
Chalcogens (Group 16)
Chalcogens are a group of nonmetals and metalloids that occupy the sixteenth column of the periodic table. The group includes oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po). Chalcogens have 6 valence electrons and tend to gain or share electrons in reactions.
O (nonmetal, high reactivity) | S (nonmetal, moderate reactivity) | Se (nonmetal, low reactivity) | Te (metalloid, low reactivity) | Po (metal, low reactivity).
Halogens (Group 17)
Halogens are a group of highly reactive nonmetals that occupy the seventeenth column of the periodic table. The group includes fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Halogens have 7 valence electrons and tend to gain electrons in reactions.
F (highly reactive nonmetal) | Cl (highly reactive nonmetal) | Br (reactive nonmetal) | I (reactive nonmetal) | At (radioactive nonmetal)
Inert gases (Group 18)
Inert gases, also known as noble gases, are a group of non-metals that occupy the eighteenth column of the periodic table. The group includes helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Inert gases have a full outer shell of 8 valence electrons and are unreactive.
He (nonmetal, unreactive) | Ne (nonmetal, unreactive) | Ar (nonmetal, unreactive) | Kr (nonmetal, unreactive) | Xe (nonmetal, unreactive) | Rn (nonmetal, radioactive)
Lanthanoids (Lanthanide series)
The Lanthanoids are a group of elements that occupy the top row of the f-block of the periodic table. The group includes cerium (Ce), praseodymium (Pr), neodymium (Nd), promethium (Pm), europium (Eu), gadolinium (Gd), terbium (Tb), dysprosium (Dy), holmium (Ho), erbium (Er), thulium (Tm), ytterbium (Yb), and lutetium (Lu). Lanthanoids have varied reactivity and valence electrons, but they tend to lose electrons in reactions.
Ce (metal, low reactivity) | Pr (metal, moderate reactivity) | Nd (metal, moderate reactivity) | Pm (metal, radioactive) | Sm (metal, reactive) | Eu (metal, reactive) | Gd (metal, reactive) | Tb (metal, reactive) | Dy (metal, reactive) | Ho (metal, reactive) | Er (metal, reactive) | Tm (metal, reactive) | Yb (metal, reactive) | Lu (metal, reactive)
Actinoids (Actinide series)
The Actinoids are a group of elements that occupy the bottom row of the f-block of the periodic table. The group includes thorium (Th), protactinium (Pa), uranium (U), neptunium (Np), plutonium (Pu), americium (Am), curium (Cm), berkelium (Bk), californium (Cf), einsteinium (Es), fermium (Fm), mendelevium (Md), nobelium (No), lawrencium (Lr). Actinoids have varied reactivity and valence electrons, but they tend to lose electrons in reactions.
Th (metal, low reactivity) | Pa (metal, moderate reactivity) | U (metal, high reactivity) | Np (metal, high reactivity) | Pu (metal, high reactivity) | Am (metal, high reactivity) | Cm (metal, high reactivity) | Bk (metal, high reactivity) | Cf (metal, high reactivity) | Es (metal, radioactive) | Fm (metal, radioactive) | Md (metal, radioactive) | No (metal, radioactive) | Lr (metal, radioactive)
Note: The reactivity and valence electron information for each element is a general trend and can vary based on the specific chemical environment and bonding interactions. Additionally, the transition metals have a variety of reactivities, from low to high, depending on the specific element and its oxidation state.
Periodic Properties of the Elements
The periodic table is a useful tool for predicting the properties of the elements based on their position in the table. The properties of the elements can be divided into two main categories: physical properties and chemical properties.
Physical Properties
Atomic Radius
Atomic radius refers to the size of an atom, and it is the half of the distance between the nuclei of two adjacent atoms. Atomic radius increases as you move down a group in the periodic table, due to the addition of a new electron shell. Atomic radius decreases as you move from left to right across a period, due to an increase in effective nuclear charge.
For example, lithium has a larger atomic radius than sodium because it has one fewer electron shell. On the other hand, fluorine has a smaller atomic radius than oxygen because of its higher effective nuclear charge.
Ionic Radius
Ionic radius refers to the size of an ion, which is an atom that has gained or lost electrons. Cations, which are positively charged ions, are smaller than their parent atoms because they have lost one or more electrons. Anions, which are negatively charged ions, are larger than their parent atoms because they have gained one or more electrons.
For example, the ionization of calcium results in the formation of a cation with a smaller ionic radius than the neutral calcium atom. The addition of an electron to oxygen results in the formation of an anion with a larger ionic radius than the neutral oxygen atom.
Ionization Energy
Ionization energy refers to the amount of energy required to remove an electron from an atom. Ionization energy generally increases as you move from left to right across a period, due to an increase in effective nuclear charge. Ionization energy generally decreases as you move down a group, due to an increase in atomic radius.
For example, the ionization energy required to remove an electron from lithium is less than that required to remove an electron from beryllium, due to the increase in effective nuclear charge in beryllium. On the other hand, the ionization energy required to remove an electron from potassium is less than that required to remove an electron from sodium, due to the increase in atomic radius in potassium.
Electronegativity
Electronegativity refers to the ability of an atom to attract electrons in a chemical bond. Electronegativity generally increases as you move from left to right across a period, due to an increase in effective nuclear charge. Electronegativity generally decreases as you move down a group, due to an increase in atomic radius.
For example, fluorine is the most electronegative element in the periodic table because of its small atomic radius and high effective nuclear charge. On the other hand, cesium is the least electronegative element in the periodic table because of its large atomic radius and low effective nuclear charge.
Chemical Properties
Reactivity
Reactivity refers to the ability of an element to undergo a chemical reaction. Reactivity generally increases as you move from left to right across a period, due to an increase in effective nuclear charge. Reactivity generally decreases as you move down a group, due to an increase in atomic radius.
For example, the alkali metals are highly reactive because they have one valence electron that is easily removed. The noble gases are unreactive because they have a full valence shell.
Metallic and Nonmetallic Character
Metallic character refers to the ability of an element to form positive ions, while nonmetallic character refers to the ability of an element to form negative ions. Metallic character generally increases as you move from right to left across a period, due to a decrease in effective nuclear charge. Nonmetallic character generally increases as you move from left to right across a period, due to an increase in effective nuclear charge.
For example, the alkali metals are highly metallic because they have one valence electron that is easily removed to form a positive ion. The halogens are highly nonmetallic because they have a tendency to gain an electron to form a negative ion.
Conclusion
The periodic table is an essential tool for understanding the properties of elements and their behavior in chemical reactions. The periodic properties of elements, such as atomic radius, ionization energy, electronegativity, reactivity, metallic character, and nonmetallic character, provide a framework for predicting chemical behavior and reactions.
In addition to understanding the periodic properties of elements, it is also important to understand the organization of the periodic table. The periodic table is arranged in order of increasing atomic number, and elements are grouped together based on their electron configurations and chemical properties.
The main groups of the periodic table include the alkali metals, alkaline earth metals, transition metals, halogens, and noble gases. Each group has unique properties and behavior in chemical reactions.
The alkali metals, for example, are highly reactive and tend to lose one electron to form a +1 ion. The noble gases, on the other hand, are unreactive and tend to have a full valence shell. The transition metals are known for their variable oxidation states and ability to form complex ions.
Understanding the periodic table and periodic properties is essential for success in chemistry, as it provides a foundation for understanding chemical behavior and predicting chemical reactions.
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