Periodic Table Class 9 ICSE — Chemistry Chapter 5

The periodic table class 9 ICSE chapter shows how chemists arranged all known elements and why the modern table works. You will meet Dobereiner’s triads, Newland’s law of octaves, and Mendeleev’s predictions. The chapter ends with group and period trends you can apply to any element. Keep your periodic table open as you read through each section. The worked patterns here use new numbers, so practise them alongside your textbook examples.

Download The Periodic Table Class 9 ICSE Chemistry Chapter 5 PDF

Download the The Periodic Table Class 9 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 18 pages, 15 sections.

What is inside Chapter 5

18 pages · 15 sections

5.1 REASONS FOR CLASSIFICATION OF ELEMENTS

Classification serves three purposes. It helps you study elements in an organised way. It links element properties to the fundamental properties of matter. It defines how one element relates to another.

Early chemists tried grouping by density, malleability, and ductility. These attempts failed for three reasons. Some groups held far too many elements. Some properties changed under different conditions. Some elements showed both metallic and non-metallic traits.

William Prout found the stable basis: atomic mass never changes for an element. That gave chemists a fixed property to classify by.

5.2 DOBEREINER’S TRIADS

German chemist J. W. Dobereiner grouped elements with similar properties into triads — sets of three. The middle element’s atomic mass equals the arithmetic mean of the other two.

For calcium, strontium, and barium:

\frac{40+137}{2} = 88.5

Strontium’s atomic mass is 88, so the triad holds.

Worked pattern: if A has mass 9 and C has mass 35, then B’s mass is \frac{9+35}{2} = 22.

The law failed for two reasons. Not all known elements could form triads. Even within one family the rule broke: for F (19), Cl (35.5), Br (80), the mean of 19 and 80 is 49.5, not 35.5.

5.3 NEWLAND’S LAW OF OCTAVES

John Newland, who loved music, arranged elements by increasing atomic mass. He noticed the eighth element repeated the first element’s properties — like the eighth note of an octave. So he divided elements into rows of seven.

The law worked well for lighter elements. Lithium, sodium, and potassium grouped together. It related properties to atomic mass. It showed periodicity for the first time. Newland received the Davy Medal in 1887.

The law failed beyond calcium. Cobalt and nickel were squeezed into one slot. Iron sat far from cobalt and nickel despite similar properties.

5.4 MENDELEEV’S PERIODIC TABLE

In 1869, Dmitri Mendeleev arranged the 63 known elements by increasing atomic mass. Similar elements fell in the same vertical column. He stated: properties are a periodic function of atomic mass.

The table had 8 groups (I-VIII). Groups I-VII split into A and B subgroups. Group VIII had none. He used R for a group’s element: Group IV oxides are RO_2, hydrides are RH_4.

Mendeleev left gaps for undiscovered elements. He predicted eka-aluminium (gallium) and eka-silicon (germanium) with accurate properties. He corrected atomic masses of gold and platinum.

Defects remained: anomalous pairs (argon before potassium), isotopes needing separate places, copper and silver grouped with alkali metals, gold separated from platinum, no electron arrangement, and hydrogen’s unsettled position.

5.5 ATOMIC NUMBER AS BASIS FOR MODERN PERIODIC LAW

In 1913, Henry Moseley fired cathode rays at metal anodes in a discharge tube. The X-rays produced had wavelengths that decreased regularly as he moved through the table’s order. He concluded that nuclear positive charge — the atomic number — is the fundamental property.

The modern periodic law states: properties are a periodic function of atomic number.

This fixed two defects. Isotopes share one position because they share the same atomic number. Argon (18) correctly precedes potassium (19).

For this ICSE Class 9 chapter, the board accepts both notations: old groups IA-VIIA, VIII, 0 and IUPAC numbering 1-18. This builds directly on the atomic structure ideas in the atomic structure and chemical bonding chapter.

5.7 TYPES OF ELEMENTS

Elements fall into four types:

• Representative elements (s and p block)
• Transition elements (d block)
• Inner transition elements (f block)
• Inert or noble gases

This block classification rests on electronic configuration, which you study in higher classes. Transition elements are metals with high melting points, conduct heat and electricity well, and some are magnetic. Many act as catalysts. They show variable valency and form coloured ions.

Inner transition elements — lanthanides and actinides — are heavy metals with coloured ions. Actinides are radioactive. Noble gases have 8 outer electrons except helium’s 2, so they barely react.

Bridge elements show diagonal resemblance: Li-Mg, Be-Al, B-Si. Period 3 elements like Na, Mg, Al summarise their groups and are called typical elements.

5.7.1 Representative elements

Representative elements fill groups 1, 2, and 13-17. Group 1 alkali metals form strong alkalis with water. Group 2 alkaline earth metals form weaker alkalis. Group 13 is the boron family, group 14 the carbon family, group 15 the nitrogen family, group 16 the oxygen family, and group 17 halogens — the salt formers.

These elements include both metals and non-metals. Metallic character gives way to non-metallic character across a period. They form electrovalent and covalent compounds with non-metals. Metallic nature increases down each group. Heavier members like tin and lead show variable valency.

5.8 MERITS OF THE MODERN PERIODIC TABLE

The modern table is based on atomic number, a better fundamental property than atomic mass. An element’s position relates directly to its electronic configuration. Properties change regularly across a period and down a group. The table is easier to remember and reproduce.

Periodic properties reappear at regular intervals or vary gradually. The cause is recurrence of similar electronic configuration. Elements in a group share the same outer electrons, so they share valency and similar properties, even as shells increase down the group.

5.9 DEFECTS OF THE MODERN PERIODIC TABLE

Two defects remain. Hydrogen’s position is still unsatisfactory — its properties relate to both Group 1 and Group 17. And the inner transition elements, lanthanides and actinides, cannot fit into the main body of the table. They sit below it instead.

5.10 GENERAL TRENDS OF THE MODERN PERIODIC TABLE

This section gives the rules you apply in questions. Learn both group and period directions.

Trends down a group

Down a group, the number of shells increases. Valence electrons equal the group number. Valency stays constant: up to group 4, valency = group number; for groups 13-17, valency = 18 – group number.

Atomic size increases down a group: \mathrm{Li} \lt \mathrm{Na} \lt \mathrm{K} \lt \mathrm{Rb} \lt \mathrm{Cs}. Metallic character increases down a group. Reactivity rises down Group 1 but falls down Group 17.

Worked pattern: element T has mass number 35 and atomic number 17. Its configuration is 2, 8, 7. Group: 17. Period: 3. Valency: 1. It is a non-metal.

Trends across a period

Across a period, shells stay constant. Valence electrons increase left to right. Valency with hydrogen rises 1 to 4 then falls to 1: LiH to HF gives 1, 2, 3, 4, 3, 2, 1. Valency with oxygen rises 1 to 7: Na_2O to Cl_2O_7.

Atomic size decreases across a period: \mathrm{Li} \gt \mathrm{Be} \gt \mathrm{B} \gt \mathrm{C} \gt \mathrm{N} \gt \mathrm{O} \gt \mathrm{F}. Nuclear pull increases as protons build up. Metallic character decreases. Oxides change from basic to acidic.

7.11 STUDY OF SPECIFIC GROUPS

This section applies group trends to real families. You will see how electronic configuration explains reactivity, physical state, and compound formation for the alkali metals, halogens, and noble gases.

7.11.1 Group I (Alkali Metals)

Alkali metals — Li, Na, K, Rb, Cs, Fr — have one valence electron. They react with water to form strong alkalis:

2M + 2H_2O \rightarrow 2MOH + H_2

They are soft enough to cut with a knife, have low melting points, and form ionic salts. They react with air, water, and dilute acids to release hydrogen. They form unipositive ions M^+ and act as strong reducing agents. Flame colours: lithium crimson red, sodium golden yellow, potassium pale violet.

Alkaline earth metals — Be, Mg, Ca, Sr, Ba, Ra — have two valence electrons. They are less reactive, harder, and greyish white. They form M^{2+} ions. Flame colours: calcium brick red, strontium crimson, barium apple green. Both families are obtained by electrolysis of molten salts.

5.11.3 Group VIIA or Group 17 (The hologens)

Halogens means “sea salt formers” in Greek. These elements — F, Cl, Br, I, At — have 7 valence electrons. They are the most reactive non-metals. Reactivity decreases down the group: fluorine is most reactive, iodine least.

Physical states change down the group: fluorine and chlorine are gases, bromine is a liquid, iodine a solid. Colours deepen from pale yellow (F) to greenish yellow (Cl) to reddish brown (Br) to violet (I). All are poisonous with a pungent odour. They form diatomic molecules F_2, Cl_2, Br_2, I_2, have valency 1, and form halides with metals. Astatine is radioactive and rare.

5.11.4 Group Zero or 18 group (Noble gases)

Mendeleev’s table had no room for this group. Argon was found by analysing air. Helium was detected in sunlight during an eclipse. William Ramsay isolated the noble gases and won the Nobel Prize.

These gases have a full octet — helium has 2. This makes them stable and unreactive. The name changed from inert to noble after xenon and krypton compounds were made in the 1960s.

They are colourless, tasteless, odourless gases. They neither burn nor support combustion. They are monoatomic and liquefy with difficulty. They emit coloured light under electric discharge. Helium fills airships and balloons. Argon fills light bulbs. Noble gases power neon advertising lights.

The alkali metals and halogens are reactive because each can reach a noble gas arrangement by losing or gaining one electron.

5.12 USES OF PERIODIC TABLE

The periodic table predicts new elements. It makes the study of elements and compounds systematic. An element’s position reveals its atomic number, electronic configuration, valence electrons, and properties.

Position also predicts bond nature. Metals in groups 1, 2, 3 lose electrons and combine with non-metals in groups 15, 16, 17 to form electrovalent compounds. Non-metals in groups 14-17 combine with each other to form covalent compounds.

Valency and metal/non-metal character are visible at a glance: metals sit at the extreme left and bottom, non-metals at the extreme right. The chapter ends with a CHAPTER AT A GLANCE summary of triads, octaves, Mendeleev’s table, and the modern periodic law. For the official scope, check the CISCE syllabus. You can also browse more Class 9 Chemistry resources on this site.

What to practise from Periodic Table Class 9 ICSE

• EXCERCISE 5(A) — 15 questions covering classification history, Dobereiner’s triads, Newland’s law of octaves, and Mendeleev’s table.
• EXCERCISE 5(C) — 15 questions covering the modern periodic table, groups, periods, electronic configuration, and specific groups. Questions 1, 2, and 15 ask you to derive group, period, and valency from electronic configuration — this pattern repeats across the exercise.

Periodic Table Class 9 ICSE Frequently Asked Questions

How should I revise the periodic table class 9 ICSE chapter for the exam?

Start with the classification history: triads, octaves, Mendeleev, then Moseley. Then memorise the group and period trends in section 5.10. Practise writing electronic configurations and deriving group, period, and valency. The exercises 5(A) and 5(C) test exactly these skills.

Why was Dobereiner’s law of triads discarded?

Dobereiner could not arrange all known elements as triads. The rule also failed inside the halogen family: the mean of F (19) and Br (80) is 49.5, not 35.5.

What is the difference between Mendeleev’s table and the modern periodic table?

Mendeleev’s table used atomic mass and left gaps for unknown elements. The modern table uses atomic number, which fixes isotope positions and anomalous pairs like argon and potassium.

How does valency change across a period and down a group?

Across a period, valency with hydrogen rises 1 to 4 then falls to 1. Valency with oxygen rises 1 to 7. Down a group, valency stays constant.

How do I find the group and period of an element from its electronic configuration?

Write the electronic configuration. The highest shell number is the period. The number of valence electrons gives the group: group = valence electrons for groups 1-2, and group = 10 + valence electrons for groups 13-18. Common mistakes: writing valency as the group number for groups 13-17, saying atomic size increases across a period, or saying metallic character increases across a period.

Reference: ICSE Class 9 Chemistry textbook material, chapter 5: The Periodic Table — CISCE-aligned.

More from this book

• Chapter 1: The Language of Chemistry
• Chapter 2: Chemical Changes and Reactions
• Chapter 3: Water
• Chapter 4: Atomic Structure and Chemical Bonding
• Chapter 6: Study of the First Element Hydrogen
• Chapter 7: Study of Gas Laws
• Chapter 8: Atmospheric Pollution