This chapter covers Periodic Classification of Elements and Periodicity from the 2nd Year (FSc Part-II) Chemistry syllabus of the Punjab Curriculum and Textbook Board (PTB/PCTB). It explains how the periodic table developed, how the modern table is arranged into periods, groups and blocks, and how the properties of elements repeat in a regular way (periodicity) across a period and down a group.
Once you understand the pattern behind the table, you can predict atomic size, ionization energy, electron affinity, electronegativity and the acidic or basic nature of oxides simply from an element position. This makes Chapter 1 the foundation for the whole of Part-II Chemistry.
Learning Objectives
- Describe the historical development of the periodic table from Dobereiner to Moseley.
- State the modern periodic law and explain why atomic number is the true basis of classification.
- Identify periods, groups and the s, p, d and f blocks.
- Explain periodicity and its cause.
- Explain the trends in atomic radius, ionization energy, electron affinity and electronegativity.
- Relate shielding effect and effective nuclear charge to these trends.
- Justify the anomalous position of hydrogen.
Key Concepts
Historical Background
Dobereiner (1829) grouped similar elements into triads of three, where the atomic mass of the middle element was about the average of the other two. Newlands (1864) gave the Law of Octaves, noticing that every eighth element repeated the properties of the first. Mendeleev (1869) arranged elements by increasing atomic mass into periods and groups, left gaps for undiscovered elements and predicted their properties. Finally Moseley (1913) proved that atomic number, not atomic mass, is the fundamental property, giving the Modern Periodic Law.
The Modern Periodic Table
Elements are arranged in order of increasing atomic number. The horizontal rows are called periods (7 of them) and the vertical columns are called groups (18 of them). Elements in the same group have the same number of valence electrons and similar properties. Important families include the alkali metals (Group IA), alkaline earth metals (Group IIA), chalcogens (Group VIA), halogens (Group VIIA) and noble gases (Group 0).
Classification into Blocks
Depending on which sub-shell receives the last electron, elements are divided into four blocks: the s-block (Groups IA and IIA), the p-block (Groups IIIA to VIIA and Group 0), the d-block (the transition elements) and the f-block (lanthanides and actinides).
Periodicity and Its Cause
Periodicity is the regular repetition of physical and chemical properties after certain regular intervals when elements are arranged by increasing atomic number. It is caused by the periodic repetition of the same outer (valence) electronic configuration; for example, all alkali metals have one electron in their outermost shell (ns1), so they behave alike.
Periodic Properties and Their Trends
Two ideas explain almost every trend. Effective nuclear charge (Zeff) is the net positive charge felt by the valence electrons, while the shielding (screening) effect is the reduction of that pull by inner-shell electrons. Across a period the atomic radius decreases while ionization energy, electron affinity and electronegativity increase. Down a group the atomic radius increases while ionization energy, electron affinity and electronegativity decrease. A cation is always smaller than its parent atom and an anion is always larger. Across a period the nature of the oxides changes from basic through amphoteric to acidic.
Position of Hydrogen
Hydrogen has one electron (1s1). Like Group IA metals it can lose one electron to form H+, and like Group VIIA halogens it can gain one electron to form H-. Because it resembles both families but is identical to neither, its position is described as anomalous and it is usually placed separately at the top of the table.
Important Definitions
What is the modern periodic law?
The properties of elements are a periodic function of their atomic numbers.
What is periodicity?
Periodicity is the regular repetition of the properties of elements after certain regular intervals when they are arranged by increasing atomic number.
Define atomic radius.
Atomic radius is the distance from the nucleus to the outermost electron shell of an atom.
Define ionization energy.
Ionization energy is the minimum energy required to remove the outermost electron from an isolated gaseous atom.
Define electron affinity.
Electron affinity is the energy released when an electron is added to an isolated gaseous atom.
Define electronegativity.
Electronegativity is the tendency of an atom to attract the shared pair of electrons in a covalent bond towards itself.
What is the shielding effect?
The shielding effect is the reduction in nuclear attraction on the valence electrons caused by the repulsion of inner-shell electrons.
What is effective nuclear charge?
Effective nuclear charge is the net positive charge actually experienced by the valence electrons of an atom.
Formulas & Rules
| Rule | Description |
|---|---|
| Zeff = Z – S | Effective nuclear charge = atomic number minus shielding constant |
| Across a period | Radius decreases; ionization energy, electron affinity and electronegativity increase |
| Down a group | Radius increases; ionization energy, electron affinity and electronegativity decrease |
| Ionic size | Cation is smaller than the parent atom; anion is larger |
| Oxides across a period | Change from basic to amphoteric to acidic |
Diagrams & Illustrations
Blocks of the periodic table: a schematic showing the s-block on the far left (Groups IA and IIA), the p-block on the right (Groups IIIA to 0), the d-block (transition elements) in the middle, and the f-block (lanthanides and actinides) placed at the bottom. Each block is named after the sub-shell being filled by the last electron.
Periodic trends: an arrow diagram showing that across a period (left to right) ionization energy, electron affinity and electronegativity increase while atomic radius decreases, and that down a group (top to bottom) atomic radius increases while those three properties decrease.
Variation of atomic radius: a graph showing atomic radius decreasing across Period 3 (Na to Cl) and increasing down Group 1 (Li to Cs), because a new shell is added down the group while nuclear charge pulls the shell inward across a period.
First ionization energy across Period 2: a graph rising from lithium to neon, with small dips at boron (its 2p electron is easier to remove than a 2s electron) and at oxygen (paired 2p electron repulsion).
Shielding effect: a Bohr-style diagram of a sodium atom showing the inner K and L shells shielding the single valence electron in the M shell from the full pull of the nucleus.
Solved Examples & Numericals
Example 1: Comparing atomic size
Which atom is larger, Na or Cl? Both are in Period 3. Moving from Na to Cl the nuclear charge increases while electrons enter the same third shell, so the shell is pulled in more tightly. Therefore Na is larger than Cl.
Example 2: Order of ionization energy
Arrange Li, Na and K in increasing order of first ionization energy. All three are Group IA metals; down the group size increases and shielding increases, so the outer electron is removed more easily. Hence K < Na < Li.
Example 3: Nature of an oxide
Predict whether MgO is acidic or basic. Magnesium is a metal (Group IIA), and metals form basic oxides, so MgO is basic and reacts with acids to form a salt and water.
Short Questions & Answers
What is a triad? Give one example.
A triad is a group of three elements with similar properties in which the atomic mass of the middle element is about the average of the other two, for example Li, Na and K.
Why does atomic radius decrease across a period?
Because the nuclear charge increases while electrons enter the same shell, pulling it inward.
Why does ionization energy increase across a period?
Smaller size and higher effective nuclear charge hold the outer electron more tightly, so more energy is needed to remove it.
Why is a cation smaller than its parent atom?
It has lost electrons, often a whole shell, so the remaining electrons are held more tightly by the same nucleus.
Why do noble gases have almost zero electron affinity?
Their outermost shells are already complete, so they have no tendency to accept an extra electron.
Why is the position of hydrogen anomalous?
Because it resembles both Group IA (it can lose one electron) and Group VIIA (it can gain one electron) but is identical to neither.
Long Questions & Answers
Q1: Trace the development of the periodic table from Dobereiner to Moseley.
Dobereiner (1829) grouped elements into triads of three similar elements, where the middle element atomic mass was about the average of the other two. Newlands (1864) proposed the Law of Octaves, noticing that every eighth element repeated the properties of the first when arranged by increasing atomic mass, but this failed for heavier elements. Mendeleev (1869) arranged elements by increasing atomic mass into periods and groups, left gaps for undiscovered elements and even predicted their properties, which was a major success. However his table had defects: an uncertain position for hydrogen, no place for isotopes, and a few anomalous pairs where a heavier element came before a lighter one. Finally Moseley (1913) showed by X-ray work that the atomic number, not the atomic mass, is the true basis of classification, which removed the anomalies and gave us the modern periodic law.
Q2: Explain the variation of atomic radius, ionization energy and electronegativity across a period and down a group.
Across a period the nuclear charge and the effective nuclear charge increase while electrons are added to the same shell. As a result the atomic radius decreases, and because the outer electrons are held more tightly, both ionization energy and electronegativity increase. Down a group a new shell is added at each step and the shielding effect increases, so the atomic radius increases while ionization energy and electronegativity decrease, since the outer electrons are farther from the nucleus and more shielded. These opposite directions explain why fluorine (top right) is small and highly electronegative while caesium (bottom left) is large and loses its electron easily.
Q3: What are effective nuclear charge and shielding effect, and how do they control periodic properties?
Effective nuclear charge is the net positive charge felt by the valence electrons after the inner electrons partly cancel the pull of the nucleus (Zeff = Z – S). The shielding or screening effect is the repulsion of the outer electrons by the inner-shell electrons, which reduces the nuclear attraction they feel. When shielding is large, Zeff is small, the outer electrons are loosely held, the atom is larger and ionization energy is low. When shielding is small, as across a period where no new inner shell is added, Zeff is large, electrons are pulled in strongly, the atom is smaller and ionization energy is high. Together these two factors explain almost every trend in the chapter.
MCQs with Answers
The modern periodic law is based on: (a) Atomic mass (b) Atomic number (c) Atomic radius (d) Density
Correct Answer: (b) Atomic number. Moseley showed atomic number is the fundamental property of an element.
Which scientist arranged elements in triads? (a) Newlands (b) Mendeleev (c) Dobereiner (d) Moseley
Correct Answer: (c) Dobereiner. He grouped similar elements in sets of three called triads.
The most electronegative element is: (a) Oxygen (b) Chlorine (c) Fluorine (d) Nitrogen
Correct Answer: (c) Fluorine. It is the smallest, most electron-attracting element.
Down a group, atomic radius: (a) Increases (b) Decreases (c) Remains same (d) First increases then decreases
Correct Answer: (a) Increases, because a new shell is added at each step.
Which of the following is an amphoteric oxide? (a) Na2O (b) MgO (c) Al2O3 (d) SO3
Correct Answer: (c) Al2O3. It reacts with both acids and bases.
Elements in which block are called transition elements? (a) s-block (b) p-block (c) d-block (d) f-block
Correct Answer: (c) d-block. Their last electron enters a d-orbital.
Ionization energy across a period generally: (a) Increases (b) Decreases (c) Stays constant (d) Becomes zero
Correct Answer: (a) Increases, due to smaller size and higher effective nuclear charge.
A cation is always ___ than its parent atom: (a) Larger (b) Smaller (c) Equal (d) Unpredictable
Correct Answer: (b) Smaller, because it loses electrons and often a whole shell.
The number of periods in the modern periodic table is: (a) 6 (b) 7 (c) 8 (d) 18
Correct Answer: (b) 7 periods (rows).
Noble gases have electron affinity that is nearly: (a) Very high (b) Zero (c) Negative infinite (d) Same as halogens
Correct Answer: (b) Zero, because their shells are already complete.
Quick Revision Summary
- Elements are arranged by increasing atomic number (Modern Periodic Law, Moseley).
- 7 periods and 18 groups; four blocks: s, p, d and f.
- Periodicity is caused by the repeating valence-shell configuration.
- Across a period: radius decreases; ionization energy, electron affinity and electronegativity increase.
- Down a group: radius increases; ionization energy, electron affinity and electronegativity decrease.
- Cation is smaller than the parent atom; anion is larger.
- Oxides across a period change from basic to amphoteric to acidic.
- Fluorine is the most electronegative element; hydrogen has an anomalous position.
Exam Tips
- Learn the two master ideas, effective nuclear charge and shielding, and you can reason out every trend instead of memorising them.
- In trend questions, always give both the direction (across or down) and the reason (size plus Zeff or shielding).
- Remember the small ionization-energy dips: B is lower than Be, and O is lower than N.
- Practise drawing and labelling the blocks diagram and the shielding diagram; labelled diagrams earn full marks.
- For nature-of-oxide questions, first decide metal or non-metal, then say basic or acidic (amphoteric for Al, Zn, Be).
- Do not confuse ionization energy (energy absorbed) with electron affinity (energy released).