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Chapter 5: Periodic Classification Of Elements
At present, 118 elements are known, each with different properties. As more elements were discovered, scientists found it increasingly difficult to organise and study their vast information individually. They sought patterns in the properties of elements to classify them into groups, making their study easier and more systematic. This led to the development of periodic classification.
Making Order Out Of Chaos – Early Attempts At The Classification Of Elements
Scientists made several early attempts to classify elements based on their properties, aiming to bring order to the growing list of known elements.
The earliest attempts simply classified elements into metals and non-metals. As knowledge increased, more sophisticated classifications were proposed.
Döbereiner’s Triads
In 1817, German chemist Johann Wolfgang Döbereiner attempted to group elements with similar properties into sets of three, which he called triads.
Döbereiner's Triad Principle: When the three elements in a triad were arranged in increasing order of their atomic masses, the atomic mass of the middle element was approximately the average of the atomic masses of the other two elements.
Example: Lithium (Li), Sodium (Na), Potassium (K)
- Atomic masses: Li = 6.9, Na = 23.0, K = 39.0
- Average atomic mass of Li and K = $\frac{6.9 + 39.0}{2} = \frac{45.9}{2} = 22.95$.
This average is very close to the atomic mass of Sodium (23.0).
Other examples of Döbereiner's Triads:
- Calcium (Ca), Strontium (Sr), Barium (Ba)
- Chlorine (Cl), Bromine (Br), Iodine (I)
Limitations of Döbereiner's Triads: Döbereiner could identify only three such triads from the elements known at that time. All the known elements could not be classified into triads. This system was not comprehensive enough to be widely useful.
Newlands’ Law Of Octaves
Encouraged by Döbereiner's work, other chemists explored relationships between atomic masses and properties. In 1866, English scientist John Newlands arranged the then known elements in increasing order of their atomic masses.
Newlands' Law of Octaves: Newlands observed that when elements were arranged in increasing order of atomic mass, the properties of every eighth element were similar to those of the first element, much like the repetition of notes in a musical octave (sa, re, ga, ma, pa, da, ni, sa...).
He started with Hydrogen (lowest atomic mass) and arranged the elements up to Thorium (56th element).
Example: Lithium (Li) properties were found to be similar to Sodium (Na), which is the eighth element after Li in his arrangement. Beryllium (Be) properties were similar to Magnesium (Mg).
| Notes | sa (do) | re (re) | ga (mi) | ma (fa) | pa (so) | da (la) | ni (ti) |
|---|---|---|---|---|---|---|---|
| H | Li | Be | B | C | N | O | |
| F | Na | Mg | Al | Si | P | S | |
| Cl | K | Ca | Cr | Ti | Mn | Fe | |
| Co and Ni | Cu | Zn | Y | In | As | Se | |
| Br | Rb | Sr | Ce and La | Zr | — | — |
Limitations of Newlands' Law of Octaves:
- The Law of Octaves was found to be applicable only up to Calcium. After Calcium, the properties of every eighth element were not similar to the first.
- Newlands assumed that only 56 elements existed and that no more elements would be discovered. This proved incorrect as many new elements were discovered later whose properties did not fit into his Law of Octaves.
- To fit existing elements into his table, Newlands sometimes placed two elements in the same slot (e.g., Cobalt and Nickel).
- He also placed unlike elements under the same note. For example, Cobalt and Nickel (metals with properties different from halogens) were placed in the same column as Fluorine, Chlorine, and Bromine (halogens). Iron, which has similar properties to Cobalt and Nickel, was placed far away from them.
- The discovery of noble gases later also challenged the Law of Octaves, as these new elements did not fit into Newlands' existing arrangement.
Newlands' Law of Octaves was successful only for lighter elements.
Making Order Out Of Chaos – Mendeléev’s Periodic Table
Despite the limitations of previous attempts, the search for a pattern correlating element properties with atomic masses continued. The most significant contribution in this early phase came from the Russian chemist Dmitri Ivanovich Mendeléev.
Mendeléev arranged elements based on their atomic mass and the similarity of their chemical properties. He primarily focused on the compounds elements formed with highly reactive elements like oxygen and hydrogen, using the formulae of oxides and hydrides as key indicators of chemical properties.
He arranged the 63 known elements in order of increasing atomic masses and observed that elements with similar physical and chemical properties recurred periodically. Based on these observations, Mendeléev formulated the Periodic Law.
Mendeléev's Periodic Law: The properties of elements are the periodic function of their atomic masses.
Mendeléev organised elements into a table with vertical columns called groups and horizontal rows called periods (Table 5.4).
In his table, the formulae of oxides and hydrides formed by the elements in each group were indicated (e.g., $\text{RH}_4$ for hydrides, $\text{RO}_2$ for oxides, where R represents any element in the group).
| Series \ Group | I | II | III | IV | V | VI | VII | VIII |
|---|---|---|---|---|---|---|---|---|
| $\text{R}_2\text{O}$ | RO | $\text{R}_2\text{O}_3$ | $\text{RO}_2$ | $\text{R}_2\text{O}_5$ | $\text{RO}_3$ | $\text{R}_2\text{O}_7$ | $\text{RO}_4$ | |
| RH | $\text{RH}_2$ | $\text{RH}_3$ | $\text{RH}_4$ | $\text{RH}_3$ | $\text{RH}_2$ | RH | ||
| 1 | H 1 | |||||||
| 2 | Li 7 | Be 9.4 | B 11 | C 12 | N 14 | O 16 | F 19 | |
| 3 | Na 23 | Mg 24 | Al 27.3 | Si 28 | P 31 | S 32 | Cl 35.5 | |
| 4 | K 39 | Ca 40 | — 44 | Ti 48 | V 51 | Cr 52 | Mn 55 | Fe 56, Co 59, Ni 59, Cu 63 |
| 5 | Cu 63 | Zn 65 | — 68 | — 72 | As 75 | Se 78 | Br 80 | |
| 6 | Rb 85 | Sr 87 | Yt 88 | Zr 90 | Nb 94 | Mo 96 | Rh 104, Ru 104, Pd 106, Ag 108 | |
| 7 | Ag 108 | Cd 112 | In 113 | Sn 118 | Sb 122 | Te 125 | I 127 | |
| 8 | Cs 133 | Ba 137 | Ce 140 | |||||
| 9 | — 180 | W 184 | Ir 193, Pt 195, Au 197 | |||||
| 10 | Au 197 | Hg 200 | Tl 204 | Pb 207 | Bi 208 | |||
| 11 | Th 232 | U 238 |
Achievements Of Mendeléev’s Periodic Table
Mendeléev's Periodic Table had several notable achievements that contributed significantly to the development of chemistry.
- Prediction of Undiscovered Elements: Mendeléev left gaps in his table, boldly predicting the existence and properties of elements that had not been discovered at that time. He named these predicted elements by prefixing 'Eka-' to the name of the preceding element in the same group (e.g., Eka-aluminium, Eka-silicon). Later discovered elements like Gallium (Ga), Scandium (Sc), and Germanium (Ge) had properties remarkably similar to those predicted for Eka-aluminium, Eka-boron, and Eka-silicon, respectively (Table 5.5). This provided strong support for his classification.
- Placement of Noble Gases: When noble gases (He, Ne, Ar, etc.), which are very inert and found in low concentrations, were discovered later, they could be placed in a new group without disturbing the existing arrangement of the table.
- Order based on Properties: Mendeléev sometimes placed an element with a slightly higher atomic mass before an element with a slightly lower atomic mass (e.g., Cobalt (58.9) before Nickel (58.7)) to ensure that elements with similar properties were grouped together. This indicated that similarity in properties was a more fundamental criterion than simply atomic mass.
| Property | Eka-aluminium (Predicted) | Gallium (Discovered) |
|---|---|---|
| Atomic Mass | 68 | 69.7 |
| Formula of Oxide | $\text{E}_2\text{O}_3$ | $\text{Ga}_2\text{O}_3$ |
| Formula of Chloride | $\text{ECl}_3$ | $\text{GaCl}_3$ |
These successes led to wide acceptance of Mendeléev's Periodic Table and established the concept of periodicity in element properties.
Limitations Of Mendeléev’s Classification
Despite its achievements, Mendeléev's classification had some limitations:
- Position of Hydrogen: Hydrogen's position was anomalous. Like alkali metals (Group 1), hydrogen forms compounds with halogens, oxygen, and sulphur with similar formulae (e.g., $\text{HCl}, \text{NaCl}$). Like halogens (Group 17), hydrogen exists as a diatomic molecule ($\text{H}_2$) and forms covalent compounds with metals and non-metals. Mendeléev could not assign a unique, correct position to hydrogen.
- Position of Isotopes: Isotopes (atoms of the same element with similar chemical properties but different atomic masses) were discovered after Mendeléev's work. According to his Periodic Law, isotopes of an element should be placed in different positions based on their atomic masses. However, placing them separately was inconsistent with their chemical similarity.
- Irregular Increase in Atomic Mass: The increase in atomic mass from one element to the next was not regular. This made it difficult to predict the exact number of elements that could be discovered between two known elements, especially for heavier elements.
Making Order Out Of Chaos – The Modern Periodic Table
The limitations of Mendeléev's classification highlighted the need for a more fundamental property as the basis for arranging elements. In 1913, Henry Moseley discovered that the atomic number of an element is a more fundamental property than its atomic mass.
The Modern Periodic Law is based on atomic number:
Modern Periodic Law: The properties of elements are a periodic function of their atomic number.
Arranging elements in increasing order of atomic number leads to the Modern Periodic Table (Table 5.6), which provides a more precise and accurate classification of elements.
The Modern Periodic Table successfully addressed the limitations of Mendeléev's table:
- The positions of Cobalt and Nickel were resolved, as they are placed in order of increasing atomic number (Nickel, Z=28, comes after Cobalt, Z=27).
- Isotopes of an element have the same atomic number and are therefore placed in the same position in the table, consistent with their similar chemical properties.
- Atomic number increases by one unit from one element to the next, providing a regular basis for arrangement.
The anomalous position of hydrogen is still a point of discussion, as it can be placed in Group 1 (due to its single valence electron, like alkali metals) or Group 17 (as it requires one electron to fill its shell, like halogens).
| Group | 1 | 2 | 3-12 | 13 | 14 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|---|---|
| Period 1 | H 1 |
He 2 |
|||||||
| Period 2 | Li 3 |
Be 4 |
Transition Metals | B 5 |
C 6 |
N 7 |
O 8 |
F 9 |
Ne 10 |
| Period 3 | Na 11 |
Mg 12 |
Al 13 |
Si 14 |
P 15 |
S 16 |
Cl 17 |
Ar 18 |
|
| Period 4 | K 19 |
Ca 20 |
Ga 31 |
Ge 32 |
As 33 |
Se 34 |
Br 35 |
Kr 36 |
|
| Period 5 | Rb 37 |
Sr 38 |
In 49 |
Sn 50 |
Sb 51 |
Te 52 |
I 53 |
Xe 54 |
|
| Period 6 | Cs 55 |
Ba 56 |
Tl 81 |
Pb 82 |
Bi 83 |
Po 84 |
At 85 |
Rn 86 |
|
| Period 7 | Fr 87 |
Ra 88 |
Nh 113 |
Fl 114 |
Mc 115 |
Lv 116 |
Ts 117 |
Og 118 |
|
| Lanthanoids and Actinoids (placed separately below) | |||||||||
Position Of Elements In The Modern Periodic Table
The Modern Periodic Table is organised into 18 vertical columns called groups and 7 horizontal rows called periods.
The position of an element in a particular group and period is determined by its electronic configuration:
- Groups: Elements in the same group have the same number of valence electrons (electrons in the outermost shell). The number of valence electrons often equals the group number (for groups 1 and 2) or the group number minus 10 (for groups 13-18, excluding Helium). As you move down a group, the number of electron shells increases. Elements in the same group have similar chemical properties due to having the same number of valence electrons.
- Periods: Elements in the same period have the same number of occupied electron shells. As you move from left to right across a period, the atomic number increases by one, and the number of valence electrons increases by one, while the number of shells remains the same. Each new period corresponds to the filling of a new electron shell.
The number of elements in each period is determined by the maximum number of electrons that can be accommodated in the corresponding shell(s):
- 1st Period (n=1, K shell): $2 \times 1^2 = 2$ elements.
- 2nd Period (n=2, L shell): $2 \times 2^2 = 8$ elements.
- 3rd Period (n=3, M shell): $2 \times 3^2 = 18$ elements theoretically, but only 8 are placed here due to electron filling patterns (first 8 elements fill up to 8 electrons in M shell, next electrons go to N shell before M continues filling).
- 4th Period (n=4): 18 elements.
- 5th Period (n=5): 18 elements.
- 6th Period (n=6): 32 elements (includes Lanthanoids).
- 7th Period (n=7): 32 elements (includes Actinoids).
The arrangement of elements by atomic number and electronic configuration explains the periodic recurrence of properties and justifies Mendeléev's idea of using properties (like formulae of compounds) for grouping elements, as these are directly related to valence electrons.
Trends In The Modern Periodic Table
Several properties of elements show predictable trends across periods (left to right) and down groups (top to bottom) in the Modern Periodic Table.
- Valency: The valency of an element is determined by the number of valence electrons. It represents an element's combining capacity.
- Across a period (left to right): Valency generally increases from 1 to 4, then decreases from 4 to 0. For example, in the 2nd period: Li (1), Be (2), B (3), C (4), N (3), O (2), F (1), Ne (0).
- Down a group: Valency generally remains the same for all elements in a group because they have the same number of valence electrons.
- Atomic Size (Atomic Radius): The atomic size is the distance from the nucleus centre to the outermost electron shell.
- Across a period (left to right): Atomic size decreases. As the atomic number increases, the nuclear charge (number of protons) increases, pulling the electron shells closer to the nucleus.
- Down a group: Atomic size increases. As you move down a group, new electron shells are added with each period, increasing the distance between the outermost electrons and the nucleus, despite the increasing nuclear charge.
- Metallic and Non-metallic Properties: Metallic character is the tendency of an element to lose electrons (electropositivity). Non-metallic character is the tendency to gain electrons (electronegativity).
- Across a period (left to right): Metallic character decreases, and non-metallic character increases. As the effective nuclear charge increases across a period, atoms hold onto their valence electrons more strongly, making it harder to lose them and easier to gain them. Metals are on the left, non-metals on the right.
- Down a group: Metallic character increases, and non-metallic character decreases. As atomic size increases down a group, the outermost electrons are farther from the nucleus and less strongly held, making it easier for metals to lose them. Non-metals become less electronegative down a group.
In the Modern Periodic Table, metals are found on the left side, non-metals on the right side, and a zig-zag line separates them. Elements along this line (boron, silicon, germanium, arsenic, antimony, tellurium, polonium) are called metalloids or semi-metals as they exhibit properties intermediate between metals and non-metals.
These trends help predict the properties of elements, including the nature of oxides they form (metal oxides are generally basic, non-metal oxides are generally acidic).
Intext Questions
Page No. 81
Question 1. Did Döbereiner’s triads also exist in the columns of Newlands’ Octaves? Compare and find out.
Answer:
Question 2. What were the limitations of Döbereiner’s classification?
Answer:
Question 3. What were the limitations of Newlands’ Law of Octaves?
Answer:
Page No. 85
Question 1. Use Mendeléev’s Periodic Table to predict the formulae for the oxides of the following elements:
K, C, AI, Si, Ba.
Answer:
Question 2. Besides gallium, which other elements have since been discovered that were left by Mendeléev in his Periodic Table? (any two)
Answer:
Question 3. What were the criteria used by Mendeléev in creating his Periodic Table?
Answer:
Question 4. Why do you think the noble gases are placed in a separate group?
Answer:
Page No. 90
Question 1. How could the Modern Periodic Table remove various anomalies of Mendeléev’s Periodic Table?
Answer:
Question 2. Name two elements you would expect to show chemical reactions similar to magnesium. What is the basis for your choice?
Answer:
Question 3. Name
(a) three elements that have a single electron in their outermost shells.
(b) two elements that have two electrons in their outermost shells.
(c) three elements with filled outermost shells.
Answer:
Question 4. (a) Lithium, sodium, potassium are all metals that react with water to liberate hydrogen gas. Is there any similarity in the atoms of these elements?
(b) Helium is an unreactive gas and neon is a gas of extremely low reactivity. What, if anything, do their atoms have in common?
Answer:
Question 5. In the Modern Periodic Table, which are the metals among the first ten elements?
Answer:
Question 6. By considering their position in the Periodic Table, which one of the following elements would you expect to have maximum metallic characteristic?
Ga Ge As Se Be
Answer:
Exercises
Question 1. Which of the following statements is not a correct statement about the trends when going from left to right across the periods of periodic Table.
(a) The elements become less metallic in nature.
(b) The number of valence electrons increases.
(c) The atoms lose their electrons more easily.
(d) The oxides become more acidic.
Answer:
Question 2. Element X forms a chloride with the formula $XCl_2$, which is a solid with a high melting point. X would most likely be in the same group of the Periodic Table as
(a) Na
(b) Mg
(c) Al
(d) Si
Answer:
Question 3. Which element has
(a) two shells, both of which are completely filled with electrons?
(b) the electronic configuration 2, 8, 2?
(c) a total of three shells, with four electrons in its valence shell?
(d) a total of two shells, with three electrons in its valence shell?
(e) twice as many electrons in its second shell as in its first shell?
Answer:
Question 4. (a) What property do all elements in the same column of the Periodic Table as boron have in common?
(b) What property do all elements in the same column of the Periodic Table as fluorine have in common?
Answer:
Question 5. An atom has electronic configuration 2, 8, 7.
(a) What is the atomic number of this element?
(b) To which of the following elements would it be chemically similar? (Atomic numbers are given in parentheses.)
N(7) F(9) P(15) Ar(18)
Answer:
Question 6. The position of three elements A, B and C in the Periodic Table are shown below –
| Group 16 | Group 17 |
|---|---|
| - | - |
| - | A |
| - | - |
| B | C |
(a) State whether A is a metal or non-metal.
(b) State whether C is more reactive or less reactive than A.
(c) Will C be larger or smaller in size than B?
(d) Which type of ion, cation or anion, will be formed by element A?
Answer:
Question 7. Nitrogen (atomic number 7) and phosphorus (atomic number 15) belong to group 15 of the Periodic Table. Write the electronic configuration of these two elements. Which of these will be more electronegative? Why?
Answer:
Question 8. How does the electronic configuration of an atom relate to its position in the Modern Periodic Table?
Answer:
Question 9. In the Modern Periodic Table, calcium (atomic number 20) is surrounded by elements with atomic numbers 12, 19, 21 and 38. Which of these have physical and chemical properties resembling calcium?
Answer:
Question 10. Compare and contrast the arrangement of elements in Mendeléev’s Periodic Table and the Modern Periodic Table.
Answer: