CBSE Class 10 Science Periodic Classification of Elements Notes
Introduction to Periodic Classification of Elements
Periodic Classification of Elements is Chapter 5 of Class 10 Science. It traces the fascinating history of how scientists attempted to organise the known elements based on their properties, culminating in the Modern Periodic Table which is one of the greatest achievements in the history of chemistry. Understanding this chapter helps you appreciate why elements behave the way they do and how their position in the periodic table predicts their chemical and physical properties.
By the mid-1800s, scientists had discovered over 60 elements and were struggling to organise them in a meaningful way. The search for patterns in elemental properties led to several classification attempts before Mendeleev's landmark periodic law and eventually the modern periodic table based on atomic number.
This chapter covers every major classification attempt in chronological order — from Dobereiner's Triads to Newlands' Law of Octaves to Mendeleev's Periodic Table to the Modern Periodic Table based on Moseley's discovery of atomic number.
Key Topics Covered
• Dobereiner's Triads — the first attempt at classification by atomic mass
• Newlands' Law of Octaves — musical analogy for elemental properties
• Mendeleev's Periodic Table — periodic law, achievements, and limitations
• Modern Periodic Table — based on atomic number (Moseley)
• Periods and Groups — their significance and trends
• Trends in the Modern Periodic Table: valency, atomic size, metallic/non-metallic character, electronegativity
1. Dobereiner's Triads (1817)
Johann Wolfgang Dobereiner, a German chemist, was one of the first scientists to identify a pattern among elements. He arranged certain elements in groups of three called triads, where the elements in each group had similar properties and the atomic mass of the middle element was approximately the arithmetic mean of the atomic masses of the first and third elements.
Atomic mass of middle element = (Atomic mass of 1st + Atomic mass of 3rd) / 2
Dobereiner's Triads Table
Triad | Elements | Atomic Masses | Verification |
Triad 1 | Li, Na, K | 7, 23, 39 | Mean of Li and K = (7+39)/2 = 23 = Na (exact) |
Triad 2 | Ca, Sr, Ba | 40, 88, 137 | Mean = (40+137)/2 = 88.5 approx = Sr |
Triad 3 | Cl, Br, I | 35.5, 80, 127 | Mean = (35.5+127)/2 = 81.25 approx = Br |
Triad 4 | S, Se, Te | 32, 79, 128 | Mean = (32+128)/2 = 80 approx = Se |
Limitations of Dobereiner's Triads
• All known elements could not be arranged into triads.
• Only a few sets of elements fit the triad pattern, making it an incomplete classification.
• Many elements discovered later did not fit into any triad.
Dobereiner identified only 3 triads from the elements known at his time. His work was a starting point, not a complete classification system.
2. Newlands' Law of Octaves (1866)
John Newlands, an English chemist, arranged the 56 known elements in increasing order of their atomic masses. He observed that every eighth element had properties similar to the first element, just like the eighth note in a musical scale repeats the first. He called this the Law of Octaves.
Every 8th element (arranged by atomic mass) has properties similar to the 1st element
Newlands' arrangement starting from hydrogen:
H Li Be B C N O
F Na Mg Al Si P S
Cl K ...
Li (2nd) was similar to Na (9th), Be (3rd) similar to Mg (10th), and so on — exactly like a musical octave.
Limitations of Newlands' Law of Octaves
• The law worked only up to calcium (Ca). Beyond Ca, the pattern broke down.
• Newlands placed two elements in the same position (e.g., Co and Ni in one slot) to maintain the pattern.
• He placed unlike elements (e.g., Fe, Co, Ni) together with unlike properties.
• When noble gases were discovered later, they could not be accommodated in his arrangement.
• The law did not account for elements yet to be discovered.
Newlands' Octaves was not accepted by the scientific community in his time. The British Chemical Society refused to publish his work. It was recognised only after Mendeleev's success.
3. Mendeleev's Periodic Table (1869)
Dmitri Mendeleev, a Russian chemist, arranged all 63 known elements in order of increasing atomic mass and noticed that elements with similar properties recurred at regular intervals. He stated the Mendeleev's Periodic Law:
Mendeleev's Periodic Law: Properties of elements are a periodic function of their atomic masses
Mendeleev arranged elements in horizontal rows called periods and vertical columns called groups. Elements in the same group had similar chemical and physical properties.
Achievements of Mendeleev's Periodic Table
• Systematic arrangement: Organised 63 elements in a logical pattern based on atomic mass and properties.
• Prediction of undiscovered elements: Left gaps for elements not yet discovered and predicted their properties (e.g., Eka-Boron = Scandium, Eka-Aluminium = Gallium, Eka-Silicon = Germanium).
• Prediction of properties: Properties of Gallium and Germanium (discovered later) matched Mendeleev's predictions almost exactly.
• Anomalies corrected: Some elements were placed in appropriate groups based on properties even if it meant ignoring atomic mass order.
Mendeleev predicted properties of Eka-Silicon (Germanium) in 1871. When Germanium was discovered in 1886, its properties matched his predictions almost exactly. This was a landmark validation of his periodic table.
Limitations of Mendeleev's Periodic Table
• Position of hydrogen: Hydrogen was placed with Group I (alkali metals) but it also shares properties with Group VII (halogens). Its position was ambiguous.
• Position of isotopes: Isotopes of the same element have different atomic masses. If arranged by atomic mass, isotopes should occupy different positions, which was not possible.
• Anomalous pairs: Certain pairs of elements (e.g., Ar and K; Co and Ni; Te and I) were placed out of atomic mass order to keep them in the correct groups based on properties.
• Lanthanides and Actinides: Rare earth elements could not be properly placed in the table.
• No explanation for periodicity: Mendeleev could not explain why properties repeated at regular intervals.
Eka Element (Predicted) | Predicted Properties | Actual Element Discovered |
Eka-Boron | Atomic mass ~44, valency 3, similar to Al | Scandium (Sc) — discovered 1879 |
Eka-Aluminium | Atomic mass ~68, low melting point, dense metal | Gallium (Ga) — discovered 1875 |
Eka-Silicon | Atomic mass ~72, forms dioxide and chloride | Germanium (Ge) — discovered 1886 |
4. Modern Periodic Table
Henry Moseley, a British physicist, discovered in 1913 that the atomic number (number of protons in the nucleus) was a more fundamental property than atomic mass. This led to the Modern Periodic Law:
Modern Periodic Law: Properties of elements are a periodic function of their atomic numbers
The Modern Periodic Table arranges all 118 known elements in order of increasing atomic number. It has 18 vertical columns called Groups and 7 horizontal rows called Periods.
Switching from atomic mass to atomic number resolved all anomalies in Mendeleev's table. Ar (18) before K (19) and Co (27) before Ni (28) are now correct since they are arranged by atomic number, not mass.
Structure of the Modern Periodic Table
• Periods (horizontal rows): There are 7 periods. Period number = number of electron shells in an atom of that element.
• Groups (vertical columns): There are 18 groups. Group number indicates valence electrons (for Groups 1 and 2) and properties.
• Period 1: 2 elements (H, He) — 1 shell
• Period 2: 8 elements (Li to Ne) — 2 shells
• Period 3: 8 elements (Na to Ar) — 3 shells
• Periods 4 and 5: 18 elements each
• Periods 6 and 7: 32 elements each (include lanthanides and actinides)
5. Trends in the Modern Periodic Table
One of the most important aspects of the Modern Periodic Table is that element properties show regular, predictable trends as you move across a period or down a group. These trends are essential for board exams.
Trend 1: Valency
• Across a period (left to right): Valency first increases from 1 to 4, then decreases from 4 to 0.
• Down a group: Valency remains the same (same number of valence electrons).
Period 3: Na(1) Mg(2) Al(3) Si(4) P(3) S(2) Cl(1) Ar(0)
Group 1: Li(1) Na(1) K(1) Rb(1) Cs(1) -- same valency down the group
Trend 2: Atomic Size (Atomic Radius)
• Across a period (left to right): Atomic size decreases. Nuclear charge increases, pulling electrons closer to the nucleus. Number of shells stays the same.
• Down a group: Atomic size increases. New electron shells are added as atomic number increases, so the outermost electrons are farther from the nucleus.
Period: Atomic size DECREASES (left to right)
Group: Atomic size INCREASES (top to bottom)
Trend 3: Metallic and Non-metallic Character
• Across a period (left to right): Metallic character decreases and non-metallic character increases. Metals lose electrons easily; non-metals gain electrons. As nuclear charge increases across a period, losing electrons becomes harder.
• Down a group: Metallic character increases and non-metallic character decreases. Larger atoms lose electrons more easily because the outermost electrons are farther from the nucleus (less nuclear attraction).
Period: Metallic character DECREASES (left to right)
Group: Metallic character INCREASES (top to bottom)
The most metallic element is Caesium (Cs) in the bottom-left of the table. The most non-metallic element (excluding noble gases) is Fluorine (F) in the top-right.
Trend 4: Electronegativity
• Across a period (left to right): Electronegativity increases. Atoms have a stronger tendency to attract bonding electrons as nuclear charge increases.
• Down a group: Electronegativity decreases. Larger atomic size means the nucleus is farther from bonding electrons, so attraction is weaker.
Period: Electronegativity INCREASES (left to right)
Group: Electronegativity DECREASES (top to bottom)
Periodic Trends Summary Table
Property | Across a Period (left to right) | Down a Group (top to bottom) |
Valency | Increases 1 to 4, then decreases to 0 | Stays the same |
Atomic size | Decreases (more nuclear charge, same shells) | Increases (more electron shells added) |
Metallic character | Decreases | Increases |
Non-metallic character | Increases | Decreases |
Electronegativity | Increases | Decreases |
Electron affinity | Increases (generally) | Decreases (generally) |
6. Comparison of Classification Attempts
Feature | Dobereiner | Newlands | Mendeleev | Modern |
Year | 1817 | 1866 | 1869 | 1913 (Moseley) |
Basis | Atomic mass (triads) | Atomic mass (octaves) | Atomic mass | Atomic number |
Law | Arithmetic mean of triad | Every 8th element similar | Periodic function of mass | Periodic function of at. no. |
Elements covered | Few (triads only) | 56 elements | 63 elements | 118 elements |
Gaps left | No | No | Yes (predicted new ones) | No |
Isotopes | Not considered | Not considered | Problem (anomalous) | No problem (same position) |
Key limitation | Incomplete | Failed beyond Ca | Anomalous pairs; isotopes | Accepted as standard |
7. Important Facts and Memory Points
• Dobereiner's Triads (1817): Groups of 3 elements; middle element's atomic mass = average of outer two.
• Newlands' Law of Octaves (1866): Every 8th element similar to 1st. Worked only up to Calcium.
• Mendeleev's Periodic Law (1869): Properties = periodic function of atomic mass. Arranged 63 elements; predicted 3 missing elements.
• Modern Periodic Law (Moseley, 1913): Properties = periodic function of atomic number. Resolved all anomalies.
• Periods: 7 horizontal rows. Period number = number of electron shells.
• Groups: 18 vertical columns. Elements in same group have same valency and similar properties.
• Atomic size across period: Decreases (left to right).
• Atomic size down group: Increases (top to bottom).
• Metallic character across period: Decreases. Down group: increases.
• Most electronegative element: Fluorine (F).
Period number tells you how many electron shells an atom has. Group number (for Groups 1 and 2) tells you the number of valence electrons. This connects electronic configuration directly to the periodic table position.
8. Board Exam Practice Questions
These questions cover all types from CBSE Class 10 Science (Chapter 5) board examinations.
1 Mark Questions
1. State Mendeleev's Periodic Law.
2. How many periods and groups are there in the Modern Periodic Table?
3. What is the basis of classification in the Modern Periodic Table?
4. State the Law of Octaves given by Newlands.
5. What happens to atomic size as we move across a period from left to right?
3 Mark Questions
1. What are Dobereiner's Triads? Give two examples and show that the middle element's atomic mass is the mean of the other two.
2. State two achievements and two limitations of Mendeleev's Periodic Table.
3. How did Moseley's work lead to the Modern Periodic Law? How does it resolve the anomalous pairs in Mendeleev's table (give example of Ar and K)?
4. State the trends in (a) atomic size and (b) metallic character across a period and down a group with reason.
5 Mark Questions
1. Compare the classification attempts by Dobereiner, Newlands, and Mendeleev under the following headings: (a) year (b) basis of classification (c) law stated (d) number of elements covered (e) main limitation.
2. Explain the following trends in the Modern Periodic Table with reasons: (a) valency across a period (b) atomic size down a group (c) metallic character across a period (d) electronegativity down a group (e) position of isotopes.
3. What were the achievements of Mendeleev's Periodic Table? Why was it necessary to modify it? Explain with the example of isotopes and anomalous pairs.
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