Selina Concise Solutions for ICSE Class 10 Chemistry Chapter 1 Periodic Table

Intext - Question - 1

Question 1.
(a) State modern periodic law. Name the scientist who stated the law.
(b) What is a periodic table? How many groups and periods does modern periodic table have?
Answer:
(a) The modern periodic law states that "The properties of elements are the periodic functions of their atomic number." Henry Moseley put forward the modern periodic law.
(b) A tabular arrangement of the elements in groups (vertical columns) and periods (horizontal rows) highlighting the regular trends in properties of elements is called a Periodic Table. Modern Periodic table has 7 periods and 18 groups.
In simple words: The modern periodic law says that when elements are arranged by their atomic number, their properties repeat in a pattern. The periodic table is like a chart that organizes all elements into rows and columns to show these patterns.

📝 Teacher's Note: Use a simple analogy like a calendar - just as dates repeat in patterns, element properties repeat when arranged by atomic number. This helps students understand periodicity before diving into electron configurations.

🎯 Exam Tip: Always mention both "atomic number" and "Henry Moseley" together - these are key marking points that examiners look for in periodic law questions.

Question 2. Why sodium element of group 1 and chlorine element of group 17 both have valency 1?
Answer: Valency is the combining capacity of the atom of an element. It is equal to the number of electrons an atom can donate or accept or share. It is just a number and does not have a positive or negative sign.
Group 1 elements have 1 electron in their outermost orbital, while Group 17 elements have 7 electrons in their outermost orbital.
Valency depends on the number of electrons in the outermost shell (i.e. valence shell).
If the number of electrons present in the outermost shell is 1, then it can donate one electron while combining with other elements to obtain a stable electronic configuration.
If the number of electrons present in the outermost shell is 7, then its valency is again 1 (8 - 7 = 1) as it can accept 1 electron from the combining atom.
In a given period, the number of electrons in the valence (outermost) shell increases from left to right. But the valency increases only up to Group 14, where it becomes 4, and then it decreases, that is, it becomes 1 in Group 17.
In simple words: Sodium has 1 electron to give away, while chlorine needs 1 electron to complete its outer shell. Both need just 1 electron to become stable, so both have valency 1.

📝 Teacher's Note: Draw electron dot diagrams for Na and Cl to show visually how one gives and one takes an electron. This makes the concept of "opposite but equal" valency crystal clear.

🎯 Exam Tip: Remember the formula: for groups 15-18, valency = 8 minus group electrons. This shortcut helps calculate valency quickly in exams.

Question 3. What are horizontal rows and vertical columns in a periodic table known as?
Answer: The horizontal rows are known as periods and vertical columns in the periodic table are known as groups.
In simple words: Think of the periodic table like a building - the floors going across are periods, and the columns going up and down are groups.

📝 Teacher's Note: Use hand gestures - horizontal for periods, vertical for groups. Students often confuse these basic terms, so physical movement helps memory.

🎯 Exam Tip: Write "horizontal = periods" and "vertical = groups" at the top of your answer sheet to avoid mixing them up during the exam.

Question 4. Periodicity is observed due to the similar ………………… (Number of valence electrons / atomic number / electronic configuration).
Answer: Periodicity is observed due to the similar electronic configuration.
In simple words: Elements behave similarly because their electrons are arranged in the same pattern in their outer shells.

📝 Teacher's Note: Emphasize that it's the overall electronic configuration pattern, not just valence electrons, that determines periodic behavior. Use examples from the same group.

🎯 Exam Tip: "Electronic configuration" is the complete answer - don't just write "valence electrons" as it's only part of the reason.

Question 5. How does the electronic configuration in atoms change
(i) in a period from left to right?
(ii) in a group top to bottom?
Answer:
(i) Though the number of shells remain the same, number of valence electrons increases by one, as we move across any given period from left to right.
(ii) While going from top to bottom in a group, the number of shells increases successively i.e. one by one but the number of valence electrons remains the same.
In simple words: Going across a period, atoms get one more electron in the same outer shell. Going down a group, atoms get a new shell but the outer shell always has the same number of electrons.

📝 Teacher's Note: Draw the first three periods showing electron configurations to illustrate these trends. Visual representation makes the pattern obvious to students.

🎯 Exam Tip: Remember "across = more electrons, same shells" and "down = more shells, same outer electrons" - this covers both parts of the question.

Question 6. Correct the statements.
(i) Elements in the same periods have equal valency.
(ii) Valency depends upon the number of shells in an atom.
(iii) Copper and zinc are representative elements.
(iv) Transition elements are placed at extreme right of the periodic table.
Answer:
(i) Elements in the same group have equal valency.
(ii) Valency depends upon the number of valence electrons in an atom.
(iii) Copper and zinc are transition elements.
(iv) Noble gases are placed at the extreme right of the periodic table.
In simple words: These corrections show that groups (not periods) have similar valency, valency depends on outer electrons (not total shells), copper and zinc are transition metals, and noble gases are at the far right.

📝 Teacher's Note: Give counter-examples for each wrong statement to help students understand why the original statements were incorrect.

🎯 Exam Tip: In correction questions, clearly write the corrected version - don't just cross out words, as partial marks may be lost.

Question 7. Name two elements in each case:
(i) Alkali metals
(ii) Alkaline earth metals
(iii) halogens
(iv) Inert gas
(v) Transition element
(vi) Lanthanides
(vii) Actinides
Answer:

In simple words: Each group of elements has specific properties - alkali metals are very reactive, halogens are toxic gases, noble gases don't react, and transition metals are strong and colorful.

📝 Teacher's Note: Connect these classifications to everyday examples - sodium in salt, chlorine in swimming pools, neon in signs, iron in tools. Real-world connections help memory.

🎯 Exam Tip: Learn at least 3 examples for each category - if you forget one name, you'll have backups to score full marks.

Question 8. What do you understand by?
(i) Periodicity:
(ii) Typical elements:
(iii) Orbits:
Answer:
(i) The properties that reappear at regular intervals, or in which there is a gradual variation at regular intervals, are called periodic properties and the phenomenon is known as the periodicity of elements.
(ii) The third period elements, Na, Mg, Al, Si, P and Cl summarize the properties of their respective groups and are called typical elements.
(iii) The elements of the second period show resemblance in properties with the elements of the next group of the third period leading to a diagonal relationship. Such elements are called bridge elements.
In simple words: Periodicity means properties repeat in patterns, typical elements represent their whole group's behavior, and bridge elements show diagonal relationships between periods.

📝 Teacher's Note: The third definition seems to be about bridge elements, not orbits. Clarify this discrepancy and explain that orbits are electron paths around the nucleus.

🎯 Exam Tip: For periodicity questions, always mention "regular intervals" and "repeating patterns" - these are key phrases examiners look for.

Question 9. Why are noble gases placed in a separate group?
Answer: Noble gases are unreactive since they have their outermost orbit complete. Due to stable electronic configuration they hardly react with other elements. So these elements are placed in a separate group i.e.18.
In simple words: Noble gases have full outer electron shells, making them very stable and unreactive, like someone who is completely satisfied and doesn't need anything more.

📝 Teacher's Note: Use the analogy of a person who has everything they need - they don't seek out interactions. This helps students remember why noble gases are inert.

🎯 Exam Tip: Always mention "complete outermost shell" and "stable configuration" - these two points together explain their unreactive nature completely.

Question 10. Name two elements you would expect to show chemical reactions similar to calcium. What is the basis of your choice?
Answer: Beryllium and magnesium will show similar chemical reactions as calcium. Since these elements belong to same group 2 and also have two electrons in their outermost shell like calcium.
In simple words: Elements in the same group behave similarly because they have the same number of electrons in their outer shell, like calcium which has 2 outer electrons.

📝 Teacher's Note: Demonstrate with simple reactions like metal + water to show how Group 2 elements react similarly, with reactivity increasing down the group.

🎯 Exam Tip: Always state the group number and mention "same number of valence electrons" as the basis for similar behavior.

Question 11. Name the metal(s) and non-metals in the first twenty elements.
Metals:
Non-metals:
Answer:
Metals: Lithium, Beryllium, Sodium, Magnesium, Aluminium, Potassium, Calcium.
Non-metals: Hydrogen, Helium, Carbon, Nitrogen, Oxygen, Fluorine, Neon, Phosphorus, Sulphur, Chlorine, Argon.
In simple words: Among the first 20 elements, metals are mostly on the left side of the periodic table, while non-metals are mostly on the right side.

📝 Teacher's Note: Point out the metalloid Silicon (element 14) which has properties of both metals and non-metals, helping students understand the gradual transition.

🎯 Exam Tip: Remember that hydrogen, though in Group 1, is a non-metal. This is a common trick question in exams.

Question 12. Name the type of elements, which have their:
(i) Outermost shell complete – ....................
(ii) Outermost shell incomplete – ……………………
(iii) two outermost shell incomplete – ……………………
(iv) one electron short of octet – ………………………
(v) two electrons in the outermost orbit - ……………
Answer:
(i) Outermost shell complete – Noble gases
(ii) Outermost shell incomplete – Representative elements
(iii) two outermost shell incomplete – Transition elements
(iv) one electron short of octet – Halogens
(v) two electrons in the outermost orbit - Alkaline Earth metals
In simple words: Each type of element is defined by how their outer electrons are arranged - complete shells make noble gases, incomplete makes representative elements, and specific numbers give specific groups.

📝 Teacher's Note: Draw electron configurations for examples of each type to show students the pattern behind these classifications.

🎯 Exam Tip: "Octet" means 8 electrons - so "one short of octet" means 7 electrons, which identifies halogens in Group 17.

Question 13. An element has 2 electrons in its N shell.
(i) What is its atomic number?
(ii) State its position in periodic table
(iii) Is is metal or non-metal?
(iv) State the name assigned to this group?
Answer:
(i) 30
(ii) It belongs to group 12 and fourth period.
(iii) It is a metal.
(iv) The name assigned to this group is IIB
In simple words: An element with 2 electrons in the N shell (4th shell) has filled the first three shells completely, giving it atomic number 30, making it zinc - a transition metal.

📝 Teacher's Note: Work through the electron filling order (K=2, L=8, M=18, N=2) to show how we arrive at atomic number 30. This reinforces electron configuration concepts.

🎯 Exam Tip: Remember the electron capacity: K=2, L=8, M=18, N=32. Add these up step by step to find atomic numbers accurately.

Question 14. State the valency of the elements of periods 3 and write the formula of their oxides.
Answer:

In simple words: Valency increases from 1 to 4, then decreases back to 1 across period 3. The oxide formulas show how many atoms of each element combine with oxygen.

📝 Teacher's Note: Explain why valency peaks at 4 (silicon) and then decreases - this demonstrates the concept of gaining vs losing electrons for stability.

🎯 Exam Tip: For sulfur, mention that both SO₂ and SO₃ are possible, but SO₂ is more common and corresponds to valency 2.

Question 15. An element A has atomic number 14. To which period does this element belong and how many elements are there in this period.
Answer: An element A with atomic number 14 belongs to period three and there are eight elements in this period.
In simple words: Element with atomic number 14 is silicon, which is in the third period that contains 8 elements from sodium to argon.

📝 Teacher's Note: Show the electron configuration (2,8,4) to demonstrate how the number of electron shells determines the period number.

🎯 Exam Tip: Period 1 has 2 elements, period 2 has 8, period 3 has 8. Memorize these numbers for quick period identification.

Question 16. Answer the following in respect of element \( ^{31}_{15} P \)
(i) Give its electronic configuration
(ii) To which group and period does it belong?
(iii) What is its valency?
(iv) Is it a metal or non - metal
(v) Is it a reducing agent or oxidizing agent?
(vi) Give its formula with chlorine.
Answer:
(i) Electronic configuration of P: 2, 8, 5
(ii) 15th Group and 3rd Period.
(iii) Valency of P = 8 - 5 = 3
(iv) It is a non-metal.
(v) It can act as both reducing and oxidizing agent depending on the reaction.
(vi) Formula with chlorine: PCl₃ or PCl₅
In simple words: Phosphorus has 5 outer electrons, is in group 15 and period 3, is a non-metal that can either gain or lose electrons, and forms compounds like PCl₃ with chlorine.

📝 Teacher's Note: Explain that phosphorus can show variable valency (3 or 5) depending on the compound formed, which is why both PCl₃ and PCl₅ are possible.

🎯 Exam Tip: For elements that can show multiple valencies, mention both possibilities unless the question specifically asks for one. This shows complete understanding.

Question 1. Name any five periods properties.
Answer: (i) Electron affinity
(ii) Atomic size
(iii) Metallic character
(iv) Non-metallic character
(v) Ionization energy
In simple words: These are properties of elements that change in a predictable pattern as you move across the periodic table from left to right or top to bottom.

📝 Teacher's Note: Help students visualize these trends by drawing simple arrows on a blank periodic table showing increase/decrease directions. This visual aid makes memorization easier.

🎯 Exam Tip: Always list five different properties - don't write similar ones like "metallic" and "non-metallic" as separate points unless specifically asked.

Question 2. What do you understand by atomic size? State its unit.
Answer: Atomic size is the distance between the centre of the nucleus of an atom and its outermost shell. It's measured in Angstrom and Picometre.
In simple words: Think of atomic size as the radius of an atom - how far the outermost electrons are from the center. It's like measuring the size of a ball from its center to its edge.

📝 Teacher's Note: Use the analogy of measuring a cricket ball's radius to help students understand atomic radius. Show that atoms aren't solid spheres but have fuzzy boundaries.

🎯 Exam Tip: Always mention both units (Angstrom and Picometre) and define it as distance from nucleus center to outermost shell for full marks.

Question 3. Give the trends in atomic size on moving:
(i) down the group
(ii) across the period right to left.
Answer: (i) The atomic size of an atom increases when we go down a group from top to bottom
(ii) It increases as we move from right to left in a period
In simple words: Atoms get bigger when you go down a column (more electron shells added) and also get bigger when you go from right to left in a row (less nuclear pull on outer electrons).

📝 Teacher's Note: Use the analogy of building floors (electron shells) going down a group, and nuclear charge effect going across periods. Draw arrows on periodic table.

🎯 Exam Tip: Remember the directions clearly - down a group = increase, right to left across period = increase. Many students confuse the period direction.

Question 4. Arrange the elements of second and third periods in increasing order of their atomic size.
(i) Second Period
(ii) Third Period
Answer: (i) Second Period: Fluorine < Neon < Oxygen < Nitrogen < Carbon < Boron < Beryllium < Lithium.
(ii) Third Period: Chlorine < Argon < Sulphur < Phosphorus < Silicon < Aluminum < Magnesium < Sodium.
In simple words: In both periods, atomic size increases from right to left, so the smallest atoms are on the right side and largest on the left side.

📝 Teacher's Note: Point out the pattern that noble gases (Neon, Argon) appear second smallest due to complete electron shells, which is an important exception students often miss.

🎯 Exam Tip: Write the order with clear inequality symbols (<) and double-check that you're going from smallest to largest as asked in "increasing order".

Question 5. Why is the size of (i) neon greater than fluorine? (ii) sodium is greater than magnesium?
Answer: (i) The size of Neon is bigger compared to fluorine because the outer shell of neon is complete (octet). As a result, the effect of nuclear pull over the valence shell electrons cannot be seen. Hence the size of Neon is greater than fluorine.
(ii) Since atomic number of magnesium is more than sodium but the numbers of shells are same, the nuclear pull is more in case of Mg atom. Hence its size is smaller than sodium.
In simple words: Neon is larger because its outer shell is full and relaxed, while magnesium is smaller than sodium because it has more protons pulling its electrons closer.

📝 Teacher's Note: Emphasize that noble gases break the usual trend due to electron-electron repulsion in complete shells. Use the analogy of a crowded bus where people spread out.

🎯 Exam Tip: For noble gas exceptions, always mention "complete octet" and "reduced nuclear pull effect" as key phrases examiners look for.

Question 6. Which is greater in size?
(i) an atom or a cation
(ii) an atom or an anion
(iii) \( Fe^{2+} \) or \( Fe^{3+} \)
Answer: (i) An atom is always bigger than cation since cation is formed by the loss of electrons; hence protons are more than electrons in a cation. So the electrons are strongly attracted by the nucleus and are pulled inward.
(ii) An anion is bigger than an atom since it is formed by gain of electrons and so the number of electrons are more than protons. The effective positive charge in the nucleus is less, so less inward pull is experienced. Hence the size expands.
(iii) \( Fe^{2+} \) is bigger than \( Fe^{3+} \) since \( Fe^{3+} \) has lost more electrons, making the nuclear pull on remaining electrons stronger.
In simple words: Losing electrons makes ions smaller (cations), gaining electrons makes them larger (anions), and losing more electrons makes even smaller ions.

📝 Teacher's Note: Use the analogy of people leaving or joining a room - fewer people means more space per person (anion), more people means crowding (cation).

🎯 Exam Tip: Always explain in terms of proton-electron ratio and effective nuclear charge for complete answers that score full marks.

Question 7. Metallic character and non-metallic character are periodic properties discuss.
Answer: The periodic variation in electronic configuration as one move sequentially in increasing order of atomic number produces a periodic variation in properties. As the elements are arranged in increasing order of atomic number, the metals with tendency to lose electrons are placed on the left and the metallic character decreases from left to right and increases down a group and non-metals with tendency to gain electrons are placed automatically on the right and the non-metallic character increase across a period and decreases down a group.
In simple words: Metallic behavior (giving away electrons easily) decreases as you go right and increases as you go down, while non-metallic behavior (taking electrons) shows the opposite pattern.

📝 Teacher's Note: Draw a periodic table showing the metal-metalloid-nonmetal boundaries. Explain that this relates to how easily atoms give or take electrons.

🎯 Exam Tip: Always mention both trends - across periods and down groups - for both metallic and non-metallic character to get complete marks.

Question 8. Give the trend in metallic character:
(i) across the period left to right,
(ii) down the group top to bottom.
Answer: (i) The metallic character decreases as we go from left to right in a period.
(ii) It increases as we go down a group
In simple words: Metallic behavior weakens as you move right across a row, but gets stronger as you move down a column in the periodic table.

📝 Teacher's Note: Connect this to ionization energy trends - easier to lose electrons means more metallic. Use examples like sodium vs chlorine in period 3.

🎯 Exam Tip: Remember the opposite trends: left to right = decrease, top to bottom = increase. Draw arrows to avoid confusion.

Question 9. State the trends in chemical reactivity:
(i) across the periods left to right
(ii) Down the group
Answer: (i) Across a period, the chemical reactivity of elements first decreases and then increases.
(ii) Down the group, chemical reactivity increases as the tendency to lose electrons increases down the group.
In simple words: Across a period, reactivity is high at both ends (metals and halogens) but low in the middle. Down a group, reactivity generally increases.

📝 Teacher's Note: Draw a U-shaped curve for period reactivity trends. Explain that noble gases in the middle are least reactive, while alkali metals and halogens at ends are most reactive.

🎯 Exam Tip: For periods, mention the U-shaped pattern - high reactivity at both ends, low in middle. This shows understanding beyond simple increase/decrease.

Question 10. State the trends in physical properties on moving down the group. Give an example to illustrate.
Answer: The melting and boiling points of metals decrease on going down the group.
Example: Observe the trend in group 1 elements given in the following table:

From the above table, it is clear that m.p. and b.p. decrease from Li to K.
The melting and boiling points of non-metals increase on going down the group.
Example: Observe the trend in Group 17 elements given in the following table:

From the above table, it is clear that m.p. and b.p. increase from F to I.
In simple words: For metals, melting and boiling points decrease down a group, but for non-metals, they increase down a group - showing opposite trends.

📝 Teacher's Note: Emphasize the contrasting trends between metals and non-metals. This helps students understand that periodic trends aren't always uniform across all elements.

🎯 Exam Tip: Always provide data tables or specific examples when asked to illustrate trends - this demonstrates thorough understanding and earns extra marks.

Question 11. An element X belong to 4th period and 17th group, state.
(i) no of valence electrons in it
(ii) name of the element.
(iii) name the family to which it belong.
(iv) Write the formula of the compound formed when it reacts with \( \frac{27}{13}Y \)
Answer: (i) The element from the 17th group has 7 electrons in its outermost shell.
(ii) The name of the element is bromine.
(iii) Bromine belongs to the halogen family.
(iv) The element \( \frac{27}{13}Y \) (aluminum) has three electrons in its outermost shell which it can donate; hence, its valency is three. While the valency of bromine is 1. Thus, aluminum can donate three electrons, and bromine can accept 1 electron to get the stable electronic configuration. Therefore, the formula of the compound is \( AlBr_3 \)
In simple words: Bromine has 7 outer electrons and belongs to the halogen family. When it reacts with aluminum (which gives away 3 electrons), we get aluminum bromide with formula AlBr₃.

📝 Teacher's Note: Show students how to decode element notation like ²⁷₁₃Y to identify aluminum. Practice identifying elements from their position in periodic table.

🎯 Exam Tip: When writing formulas, use the criss-cross method with valencies to get the correct subscripts. Always check that charges balance out.

Question 12. The given table shows elements with the same number of electrons in its valence shell.

State:
(i) Whether these elements belong to same group or period.
(ii) Arrange them in order of increasing metallic character.
Answer: (i) Yes, these elements belong to the same group but are not from the same period.
(ii) We know that m.p. decreases on going down the group. Hence, from the above table, the elements can be ordered according to their period as follows:

The metallic character increases as one moves down the group. Hence, the order of the given elements with increasing metallic character is as follows: B < C < A
In simple words: Since these elements have same outer electrons, they're in the same group. The one with highest melting point is at the top, and metallic behavior increases going down.

📝 Teacher's Note: Help students connect melting point trends with position in periodic table. Lower melting point indicates the element is further down the group.

🎯 Exam Tip: Always arrange data in order first, then apply the trend rules. This systematic approach prevents errors in ordering elements.

Question 1. (a) Define the term 'ionisation potential'
(b) Represent in the form of an equation. In which unit it is measured?
Answer: (a) The energy required to remove an electron from a neutral isolated gaseous atom and convert it into a positively charged gaseous ion is called Ionization energy or ionization potential.
(b) M (g) + I.E \( \rightarrow \) M⁺(g) + e⁻
M can be any element
It is measured in electron volts per atom. Its S.I unit kJmol⁻¹.
In simple words: Ionization energy is the energy needed to pull an electron completely away from an atom, like the energy needed to remove a magnet from a metal surface.

📝 Teacher's Note: Emphasize that the atom must be gaseous and isolated to avoid interference from other atoms. Use the magnet analogy to explain why energy is needed.

🎯 Exam Tip: Always write the equation with proper state symbols (g) and include both the common unit (eV) and SI unit (kJ/mol) for full marks.

Question 2. What do you understand by successive ionization energies?
Answer: The energy required to remove the residual electrons one by one is called successive ionization energy.
In simple words: After removing the first electron, it takes even more energy to remove the second electron, then more for the third, and so on - like removing layers of protection one by one.

📝 Teacher's Note: Explain that successive ionization energies always increase because you're removing electrons from an increasingly positive ion. Draw a simple energy diagram.

🎯 Exam Tip: Mention that successive ionization energies show large jumps when electrons are removed from inner, more stable electron shells.

Question 3. State the trends in ionization energy:
(a) across the period:
(b) down the group.
Answer: (a) Ionization energy increases as we move from left to right across a period as the atomic size decreases.
(b) Ionization energy decreases down a group as the atomic size increases.
In simple words: It gets harder to remove electrons as you move right across a period (smaller atoms, stronger pull), but easier as you go down a group (bigger atoms, weaker pull).

📝 Teacher's Note: Connect this directly to atomic size trends - smaller atoms hold electrons more tightly. Use the analogy of holding objects close vs far from your body.

🎯 Exam Tip: Always link ionization energy trends to atomic size changes - this shows understanding of the underlying cause, not just memorized patterns.

Question 4. Name the elements with highest and lowest ionization energies.
Answer: Helium has the highest ionization energy of all the elements while cesium has the lowest ionization energy
In simple words: Helium (smallest noble gas) holds its electrons most tightly, while cesium (largest alkali metal) holds its electrons most loosely.

📝 Teacher's Note: Point out the logical extremes - helium is small and stable, cesium is large and wants to lose electrons easily. This reinforces the size-ionization energy relationship.

🎯 Exam Tip: Remember the extremes: helium (highest) and cesium (lowest) - these are commonly asked and easy to remember by their positions in periodic table.

Question 5. Arrange the elements of second and third period in increasing order of ionization energy.
Answer: Second period: Lithium < Beryllium < Boron < Carbon < Nitrogen < Oxygen < Fluorine < Neon
Third period: Sodium < Magnesium < Aluminum < Silicon < Phosphorus < Sulphur < Chlorine < Argon
In simple words: Both periods show the same pattern - ionization energy increases from left to right, with noble gases having the highest values.

📝 Teacher's Note: Point out that this order is essentially the reverse of atomic size order. Help students see the inverse relationship between size and ionization energy.

🎯 Exam Tip: Write the order carefully from lowest to highest as asked. Double-check by remembering that alkali metals have lowest and noble gases have highest ionization energies.

Third Period: Argon > Chlorine > Sulphur > Phosphorus > Silicon > Aluminum > Magnesium > Sodium

Question 6:
(a) Define the term 'electron affinity'.
(b) Arrange the elements of second period in increasing order of their electron affinity. Name the elements which do not follow the trend in this period.
Answer:
(a) Electron affinity is the energy released when a neutral gaseous atom acquires an electron to form an anion.
(b) Second period: Lithium < Boron < Carbon < Oxygen < Fluorine
Neon, Nitrogen and Beryllium do not follow the trend.
In simple words: Electron affinity is like how much an atom "wants" to grab an extra electron. In the second period, most atoms want electrons more as we go from left to right, but some elements break this pattern.

📝 Teacher's Note: Use the analogy of "hunger for electrons" to explain electron affinity. Show students that some elements are exceptions because of their stable electron arrangements.

🎯 Exam Tip: Always mention the exceptions (N, Be, Ne) when discussing electron affinity trends in period 2 to score full marks.

Question 7:
State the factors on which electron affinity depends.
Answer: Electron affinity depends on:
(a) Atomic size
(b) Nuclear charge
In simple words: Electron affinity depends on how big the atom is and how strong its nucleus pulls electrons. Smaller atoms with stronger nuclear pull have higher electron affinity.

📝 Teacher's Note: Relate atomic size to distance and nuclear charge to attraction force. Use the magnet analogy - closer and stronger magnet pulls harder.

🎯 Exam Tip: Write both factors clearly and briefly explain how each affects electron affinity for complete marks.

Question 8:
Electron affinity values generally _______ across the periods left to right and ______ down the group top to bottom.
Answer: Electron affinity values generally increases across the periods left to right and decreases down the group top to bottom.
In simple words: As you move right across a period, atoms want electrons more. As you move down a group, atoms want electrons less.

📝 Teacher's Note: Use periodic table visuals to show the trend arrows. Emphasize that this is the general trend with some exceptions.

🎯 Exam Tip: Remember "increases across, decreases down" as a simple rule for electron affinity trends.

Question 9:
Give reason:
(a) Electron affinity of halogens is comparatively high,
(b) Electronegativity of chlorine is higher than Sulphur.
Answer:
(a) As we move from left to right the increase in atomic number and decrease in size results in a greater nuclear pull. As a result, the ability to attract the electrons increases, and so does the electron affinity. But noble gases have complete stable octet configuration, hence their electron affinity is lower than halogens. Hence halogens on extreme right have highest electron affinity in a period.
(b) Chlorine is smaller than sulphur with a bigger atomic number. Since its nuclear pull is more, hence its electron affinity is higher than sulphur.
In simple words: Halogens are very "greedy" for electrons because they need just one more to complete their outer shell. Chlorine pulls electrons stronger than sulphur because it's smaller and has more protons.

📝 Teacher's Note: Explain that halogens need only one electron to achieve stability, making them highly electronegative. Use visual models to show atomic sizes.

🎯 Exam Tip: Mention both atomic size and nuclear charge when explaining electronegativity differences between elements.

Question 10:
Why fluorine has higher E.N. than chlorine?
Answer: Since size of chlorine is bigger than fluorine hence the electrons being farther away from the nucleus experience a lesser force of attraction, hence electron negativity of chlorine is less than fluorine
In simple words: Fluorine is smaller than chlorine, so its nucleus can pull electrons more strongly. It's like being closer to a magnet - the pull is stronger.

📝 Teacher's Note: Use atomic radius data to show the size difference. Emphasize that smaller size means stronger pull on valence electrons.

🎯 Exam Tip: Always relate electronegativity differences to atomic size when comparing elements in the same group.

Question 11:
Define the term 'Electronegativity' state its unit.
Answer: Electronegativity measures an atom's tendency to attract shared pair of electrons towards itself. Its S.I unit is Pauling unit.
In simple words: Electronegativity is how much an atom pulls on shared electrons in a bond. It's measured in Pauling units, named after a famous scientist.

📝 Teacher's Note: Distinguish between electron affinity (isolated atoms) and electronegativity (bonded atoms). Mention Linus Pauling's contribution to chemistry.

🎯 Exam Tip: Include both the definition and unit in your answer. Remember electronegativity deals with bonded atoms, not isolated ones.

Question 12:
(a) Name the elements with highest and lowest electronegativity,
(b) State the character of the oxide of period 3.
Answer:
(a) The element fluorine has the highest electronegativity and Caesium has the lowest electronegativity.
(b) The nature of oxides changes from basic to acidic as we move from left to right in third period. Hence sodium forms most basic oxide while oxide of Aluminum is amphoteric and oxides of phosphorus, sulphur and chlorine are progressively acidic.
In simple words: Fluorine pulls electrons the strongest, while caesium pulls them the weakest. In period 3, oxides start basic on the left and become acidic on the right.

📝 Teacher's Note: Show the periodic table positions of F and Cs to explain the trend. Use examples of common oxides like \( \text{Na}_2\text{O} \) (basic) and \( \text{SO}_2 \) (acidic).

🎯 Exam Tip: Remember F (highest) and Cs (lowest) for electronegativity. Mention amphoteric nature of aluminum oxide for complete answer.

Question 13:
Name the periodic property which relates to the:
(a) Amount of energy required to remove an electron from an isolated gaseous atom,
(b) character of element which loses one or more electrons when supplied with energy.
(c) tendency of an atom to attract the shared pair of electron.
Answer:
(a) Ionization energy
(b) Metallic character
(c) Electronegativity
In simple words: These are three important properties: how hard it is to remove electrons (ionization energy), how easily elements give up electrons (metallic character), and how strongly atoms pull shared electrons (electronegativity).

📝 Teacher's Note: Help students connect the definitions to the property names. Use examples of metals losing electrons easily and non-metals attracting electrons.

🎯 Exam Tip: Learn the exact definitions of each property. These are frequently asked in one-word answer questions.

Question 14:
Explain the following:
(a) Group 17 elements are strong non-metals, while group 1 elements are strong metals
(b) Metallic character of elements decreases from left to right in a period while it increases in moving down a group.
(c) Halogens have a high electron affinity.
(d) The reducing power of element increases down in the group while decreases in a period.
(e) Size of atom progressively becomes smaller when we move from sodium (Na) to chlorine (Cl) in the third period of the periodic table.
Answer:
(a) On moving across a period, nuclear pull increases because of the increase in atomic number, and thus, the atomic size decreases. Hence, elements cannot lose electrons easily. Hence, Group 17 elements are strong non-metals, while Group 1 elements are strong metals.
(b) On moving across a period, nuclear pull increases because of the increase in atomic number, and thus, the atomic size decreases. Hence, elements cannot lose electrons easily. Hence, Group 17 elements are strong non-metals, while Group 1 elements are strong metals. Down a group, the atomic size increases and the nuclear charge also increases. The effect of an increased atomic size is greater as compared to the increased nuclear charge. Therefore, metallic nature increases as one moves down a group, i.e. they can lose electrons easily.
(c) The atomic size of halogens is very small. The smaller the atomic size, the greater the electron affinity, because the effective attractive force between the nucleus and the valence electrons is greater in smaller atoms, and so the electrons are held firmly.
(d) The reducing property depends on the ionisation potential and electron affinity of the elements. In a period, from left to right in a horizontal row of the periodic table, the atomic size decreases and the nuclear charge increases, so the electron affinity and ionisation energy both increase. Hence, the tendency to lose electrons decreases across the period from left to right and thus the reducing property also decreases across the period from left to right. The electron affinity and ionisation potential decreases along the group from top to bottom. Hence, the tendency to lose electrons increases, and thus, the reducing property also increases along the group from top to bottom.
(e) In a period, the size of an atom decreases from left to right. This is because the nuclear charge, i.e. the atomic number increases from left to right in the same period, thereby bringing the outermost shell closer to the nucleus. Therefore, considering the third period given above, it has been found that sodium is the largest in size, while chlorine is the smallest.
In simple words: These trends happen because of atomic size and nuclear charge changes. Group 1 elements easily give up electrons (metals), while Group 17 elements strongly attract electrons (non-metals). As we move across periods or down groups, these properties change in predictable patterns.

📝 Teacher's Note: Use atomic models to show how nuclear charge and atomic size affect electron behavior. Connect all trends to the fundamental concepts of attraction and repulsion.

🎯 Exam Tip: For multi-part explanations, always mention both atomic size and nuclear charge effects. Use specific examples like Na to Cl for clarity.

Exercise: 1

Question 1:
(a) How does the electronic configuration of an atom relate to its position in the modern periodic table?
(b) Write the number of protons, neutrons and electronic configuration of \( ^{39}_{19}K \), \( ^{31}_{15}P \). Also state their position in periodic table.
Answer:
(a) The total number of electron shells in an atom determines the period to which the element belongs, and the valence electrons determine the group to which it will belong. So with the help of electronic configuration we can figure out the period and group number of an element. Elements with one and two valence electrons belong to group 1 and 2 respectively, while to determine the group number of elements with 3 to 8 valence electrons, we add 10 to their valence electrons. For example an element X has atomic number 15. Its configuration will be: K shell has 2 electrons, L will have 8, and the remaining 5 will be placed in M shell. Since it has three shells it belongs to period 3 and with 5 valence electrons the element will be placed in five plus ten that is the 15th group. So with the help of electronic configuration we can figure out the period and group number of an element.
(b) Atomic number = Number of protons. Hence, number of protons in K atom = 19. Number of neutrons = Mass number − Atomic number. Hence, number of neutrons in K atom = 39−19 = 20. Number of electrons = Number of protons. Hence, number of electrons = 19. And electronic configuration of K atom = 2, 8, 8, 1. Since K atom has 4 shells, hence it belongs to fourth period. With one valence electron, it belongs to group 1. Number of protons in P atom = 15. Number of neutrons in P atom = 31−15 = 16. Number of electrons in P atom = 15. And electronic configuration of P atom = 2, 8, 5. Since it has three shells, it belongs to period 3 and with 5 valence electrons Phosphorus is found in five plus ten that is 15th group.
In simple words: Electronic configuration is like an address system for elements. The number of electron shells tells you the period (row), and valence electrons tell you the group (column) in the periodic table.

📝 Teacher's Note: Use the "address system" analogy to help students understand. Practice with multiple examples to reinforce the shells→period and valence→group concept.

🎯 Exam Tip: For group number calculation, remember: Groups 1-2 use valence electrons directly, Groups 13-18 use valence electrons + 10.

Question 2:
Fluorine, chlorine and Bromine are put in one group on basis of their similar properties.
(a) what are those similar properties?
(b) What is the common name of this group or family?
Answer:
(a) Fluorine, chlorine and bromine are non-metals with seven valence electrons. They are highly electronegative elements with valency of one. They exist as diatomic molecules. They form ionic compounds with alkali metals
(b) They are known as halogens. The term means salt forming and therefore compounds containing these elements are called salts.
In simple words: These three elements all have 7 valence electrons, are very reactive non-metals, and love to form salts. That's why they're called halogens, which means "salt makers."

📝 Teacher's Note: Show examples of common salts like NaCl, NaBr to illustrate the "salt forming" nature. Emphasize the diatomic nature (F₂, Cl₂, Br₂).

🎯 Exam Tip: Mention all key properties: 7 valence electrons, diatomic molecules, high electronegativity, and salt formation with metals.

Question 3:
What is the main characteristic of the last element in each period of the periodic table? What is the general name of such elements?
Answer: The last element in each period of the periodic table is a gaseous element with its valence shell completely filled. Except for helium with complete duplet configuration, rest all the 5 gases have complete octet configuration. These group 18 elements are commonly referred to as noble gases.
In simple words: The last elements in each period have complete outer electron shells, making them very stable and unreactive. They're called noble gases because they rarely react with other elements.

📝 Teacher's Note: Explain why complete electron shells make these gases unreactive. Use the analogy of being "satisfied" and not needing more electrons.

🎯 Exam Tip: Mention both the complete electron configuration and the term "noble gases" for full marks.

Question 4:
According to atomic structure, what determines which element will be the first and which will be the last in a period?
Answer: The electronic configuration of an element determines its position in Modern Periodic table. The element with one valence electron is the first while the element with 8 valence electrons is placed in the 18th group of a period.
In simple words: The number of valence electrons decides the order. Elements start with 1 valence electron and end with 8 valence electrons (complete outer shell) in each period.

📝 Teacher's Note: Show students the pattern across different periods. Emphasize that this creates the repeating pattern that gives the table its name "periodic."

🎯 Exam Tip: Clearly state that 1 valence electron = first element and 8 valence electrons = last element in a period.

Question 5:
How does the number of:
(i) valence electrons and
(ii) valency vary on moving from left to right in the second period of the periodic table?
Answer:
(i) The number of valence electrons increases by one as we move from left to right in a period. The group number 1 and 2 have 1 and 2 valence electrons respectively while group 13 to 18 have group number minus 10 = valence electrons. So, group 13 to 18 have 3, 4, 5, 6, 7 and 8 valence electrons respectively.
(ii) Valency is determined by the number of valence electrons. For elements belonging to group 1, 2 and 13, the valency is equal to the number of valence electrons, so their valency is 1, 2 and 3 respectively.
In simple words: As you move right across a period, valence electrons increase by 1 each time. Valency follows a pattern - it increases up to 4, then decreases back to 1.

📝 Teacher's Note: Explain the valency pattern: increases from 1 to 4, then decreases from 4 to 1. Use examples from period 2 to illustrate this concept clearly.

🎯 Exam Tip: Remember the valency pattern across a period: 1,2,3,4,3,2,1,0. This helps in predicting chemical formulas.

Since the elements in group 14 to 17 needs to gain electrons to complete their octet configuration. Their valency is 8 minus the number of valence electrons. So their valencies are 4, 3, 2 and 1 respectively.

Question. Fill in the blanks:
(a) The horizontal rows in the periodic table are called …………
(b) On moving across a period from right to left in periodic table, the atomic size of the atom ……….
(c) on moving from right to left in the second period, the number of valence electrons……….
Answer:
(a) The horizontal rows in the periodic table are called Periods.
(b) On moving across a period from right to left in periodic table, the atomic size of the atom increases.
(c) on moving from right to left in the second period, the number of valence electrons decreases
In simple words: Periods are the horizontal lines in the periodic table. When you go from right to left in a period, atoms get bigger because there are fewer protons pulling the electrons closer. Also, valence electrons decrease as you move right to left.

📝 Teacher's Note: Use visual aids showing atomic structure to demonstrate how atomic size changes. Have students trace their finger across a period to reinforce the direction of trends.

🎯 Exam Tip: Remember the mnemonic "Right to Left = Bigger and Less" for atomic size increase and valence electron decrease.

Question. An element barium has atomic number 56. Look up its position in the periodic table and answer the following questions.
(a) Is it a metal or a non – metal?
(b) Is it more or less reactive than calcium?
(c) What is its valency?
(d) What will be the formula of its phosphate?
(e) Is it larger or smaller than caesium (Cs)?
Answer:
(a) Since it belongs to group II, it has 2 valence electrons and hence it is a metal.
(b) Barium is placed below calcium in the group. Since, it has more number of shells; it is easier for it to lose its valence electrons to complete its octet configuration. Hence it is more reactive than calcium.
(c) It needs to lose its 2 valence electrons to complete its octet configuration; therefore its valency is also 2.
(d) The formula of its phosphate will be \( (Ba)_3 (PO_4)_2 \)
(e) As we move from left to right in a period, the size decreases, therefore, it will be smaller than Cesium
In simple words: Barium is a metal in group 2 that easily gives away its 2 outer electrons. It's more reactive than calcium because it's bigger and loses electrons more easily. Since barium has +2 charge and phosphate has -3 charge, you need 3 barium and 2 phosphate to balance.

📝 Teacher's Note: Show students how to balance ionic compounds by crisscross method. Emphasize that reactivity increases down a group for metals because electrons are farther from nucleus.

🎯 Exam Tip: For formula writing, always check that total positive and negative charges balance to zero. Group number equals valency for groups 1-2.

Question. How do the following change on moving from left to right in a period of the periodic table? Give examples in support of your answer.
(a) atomic structure (electron arrangements)
(b) Chemical reactivity of elements.
(c) Nature of oxides of the elements.
Answer:
(a) The number of valence electrons increases by one as we move across any given period. Therefore as we move from Lithium to Neon in period 2, the valence electrons will increase from 1 to 7.
(b) The metallic character decreases as we move from left to right while the non metallic character increases. Ongoing from left to right in a period, the chemical reactivity of elements first decreases and then increases. For example in period 3, Sodium is the most reactive metal and Chlorine is the most reactive non-metal and Silicon is least reactive
(c) The oxides of metals are basic and that of non-metals are acidic in general. Therefore since metallic strength decreases and non-metallic strength increases on moving from left to right across a period, the strength of basic oxides decreases, while the strength of acidic oxides increases. For example, sodium forms a basic oxide, while sulphur and phosphorus form acidic oxides.
In simple words: As you move left to right across a period, each element has one more valence electron. Reactivity first drops (metals become less reactive) then rises (non-metals become more reactive). Oxides change from basic (like soap) to acidic (like lemon juice).

📝 Teacher's Note: Create a periodic table chart showing the trends visually. Use household examples like baking soda (basic) and vinegar (acidic) to explain oxide nature.

🎯 Exam Tip: Remember that noble gases (column 18) are exceptions to reactivity trends. Always give specific examples from the same period when explaining trends.

Question. This question refers to the elements of the periodic table with atomic number from 3 to 18. Some of the elements are shown by letters, but the letters are not the usual symbols of the elements.

Which of these:
(a) have most electronegative element.
(b) is a halogen?
(c) is an alkali metal?
(d) is an element with valency 4?
(e) have least ionization energy?
(f) have least atomic size in period 3.
Answer:
(a) Noble gases- H and P
(b) Halogens- G and O
(c) Alkali metals - A and I
(d) D and L have valency of 4
(e) I with atomic number 11.
(f) Cl has the least atomic size in period 3 with atomic number 17.
In simple words: Look at the position in the periodic table to identify element families. Noble gases are most electronegative, halogens need 1 electron, alkali metals have 1 valence electron, group 14 has valency 4, and bigger atoms lose electrons easier.

📝 Teacher's Note: Create flashcards with group numbers and their properties. Practice identifying families by their position rather than memorizing individual elements.

🎯 Exam Tip: Count the group number from the given atomic numbers to identify families quickly. Group 1=alkali metals, Group 17=halogens, Group 18=noble gases.

Question. In group I of the periodic table, three elements X,Y and Z have ionic radii 1.33 Å, 0.95 Å and 0.60 Å respectively. Giving a reason, arrange them in the order of increasing atomic number in the group.
Answer: As we move down a group, the numbers of shells increases and hence the atomic size increases. Therefore, Z will have the smallest atomic number followed by Y, while X will have the largest atomic number. So the elements in order of increasing atomic number will be ZIn simple words: In a group, atoms get bigger as you go down because they have more electron shells. So the smallest atom (Z) is at the top and largest atom (X) is at the bottom.

📝 Teacher's Note: Use concentric circles to show how additional electron shells make atoms larger down a group. Relate this to Russian nesting dolls getting bigger.

🎯 Exam Tip: Smaller ionic radius = higher position in group = smaller atomic number. This relationship is always inverse in the same group.

Question. How does the chemical reactivity of:
(a) alkali metals vary?
(b) halogens vary?
Answer:
(a) Since, the distance of the valence electrons from the nucleus keeps on increasing down the group, therefore, the ionization energy keeps on decreasing. Hence the reactivity of alkali metals increases from lithium to francium.
(b) As we move down a group, the size keeps on increasing, so it becomes more difficult for atoms to attract electrons. Thus reactivity of halogens decreases from Fluorine to Astatine.
In simple words: Alkali metals get more reactive going down because it's easier to lose electrons when they're far from the nucleus. Halogens get less reactive going down because it's harder to grab electrons when the atom is bigger.

📝 Teacher's Note: Demonstrate with videos of alkali metals reacting with water - show how cesium reacts more violently than lithium. Emphasize opposite trends for metals vs non-metals.

🎯 Exam Tip: Remember: metals lose electrons (easier down the group), non-metals gain electrons (harder down the group). Reactivity trends are opposite for metals and non-metals.

Question. An element X belong to 3rd periods and group II of the periodic table state:
(a) the number of valence electrons,
(b) the valency,
(c) name of the element,
(d) whether it is a metal or a non-metal.
Answer:
(a) Since it belongs to period 3 it has 3 shells, K, L and M. The outermost M shell will have 2 valence electrons as it is placed in group II
(b) With 2 valence electrons, its valency will be 2.
(c) Since it has electronic configuration of 2, 8, 2, its atomic number is 12 and hence X is Magnesium
(d) It is a metal.
In simple words: Period number tells you how many electron shells, group number tells you valence electrons. Adding up all electrons gives atomic number, which identifies the element as magnesium.

📝 Teacher's Note: Practice electron configuration writing with students. Show how period and group numbers directly give shell and valence electron information.

🎯 Exam Tip: For groups 1-2, valency equals group number. For groups 13-18, valency equals (18 - group number). Always verify by electron configuration.

Question. The electronic configuration of an element T IS 2, 8, 8, 1.
(a) What is the group number of T?
(b) What is the period number of T?
(c) How many valence electrons are there in an atom of T?
(d) What is the valency of T?
(e) Is it a metal or a non-metal?
Answer:
(a) Group 1since the valence electrons is 1
(b) With 4 shells T belong to period 4.
(c) Number of electrons = 2+8+8+1=19
(d) T needs to lose one electron to complete its octet hence its valency is 1
(e) Since it has one valence electron, it is a metal.
In simple words: Count the electron shells to get period number (4 shells = period 4). The outermost shell electrons give group number (1 electron = group 1). Elements with 1-3 valence electrons are metals.

📝 Teacher's Note: Emphasize that metals have few valence electrons and prefer to lose them. Use the electron configuration to deduce all properties systematically.

🎯 Exam Tip: Always count total electrons to verify atomic number. Number of shells = period number, outermost shell electrons = group number (for main groups).

Question. Arrange the elements of group 17 and group 1 according to the given conditions.
(a) Increasing order of atomic size,
(b) Increasing non – metallic character
(c) Increasing ionization potential
(d) Increasing electron affinity
(e) Decreasing electro negativity.
Answer:
(a) Group 1: Lithium< Sodium< Potassium< Rubidium < Caesium< Francium
Group 17: Fluorine < Chlorine < Bromine< Iodine < Astatine
(b) Group 1: Francium
Group 17: Astatine< Iodine< Bromine< Chlorine< Fluorine
(c) Group 1: Francium< Cesium< Rubidium< Potassium< Sodium< Lithium
Group 17: Astatine< Iodine< Bromine< Chlorine< Fluorine
(d) Group 1: Francium
Group 17: Astatine
(e) Group 1: Lithium>Sodium> Potassium> Rubidium> Cesium> Francium
Group 17: Fluorine > Chlorine> Bromine > Iodine > Astatine
In simple words: Down a group, atoms get bigger, so atomic size increases. Non-metallic character and ionization potential decrease down both groups. Electronegativity decreases down groups (top elements hold electrons more tightly).

📝 Teacher's Note: Create a trends chart showing arrows for increasing/decreasing properties. Students should memorize that most properties decrease down a group except atomic size.

🎯 Exam Tip: Remember "Size up, everything else down" for group trends. Fluorine is always highest for electronegativity, ionization energy, and electron affinity.

Question. Complete the following sentences choosing the correct word or words from those given in brackets at the end of each sentence:
(a) The properties of the elements are a periodic function of their …………… (atomic number, mass number, reative atomic mass).
(b) Moving across a ………….. of the periodic table the elements show increasing ………………..character (group, period, metallic, non-metallic).
(c) The elements at the bottom of a group would be expected to show …….. metallic character than the element at the top. (less, more).
(d) The similarities in the properties of a group of elements are because they have the same ………………… (electronic configuration, number of outer electrons, atomic numbers).
Answer:
(a) The properties of the elements are a periodic function of their atomic number (atomic number, mass number, reative atomic mass).
(b) Moving across a periods of the periodic table the elements show increasing non-metallic character (group, period, metallic, non-metallic).
(c) The elements at the bottom of a group would be expected to show more metallic character than the element at the top. (less, more).
(d) The similarities in the properties of a group of elements are because they have the same number of outer electrons (electronic configuration, number of outer electrons, atomic numbers).
In simple words: Properties repeat based on atomic number (Mendeleev's law). Across periods, elements become more non-metallic. Down groups, elements become more metallic. Same valence electrons give similar properties.

📝 Teacher's Note: Emphasize that periodic law is based on atomic number, not mass. Show examples of how valence electrons determine chemical behavior across families.

🎯 Exam Tip: Atomic number determines properties, not atomic mass. Period trends go towards non-metallic, group trends go towards metallic character.

Question. Give reasons for the following:
(a) The size of the anion is greater than the size of the parent atom.
(b) argon atom is bigger than chlorine atom.
(c) Ionisation potential of the element increases across a period.
Answer:
(a) When an atom gains electrons to form an anion, the electron-electron repulsion increases while the nuclear charge remains the same. This causes the electron cloud to expand, making the anion larger than the neutral atom.
(b) Argon atom is bigger than chlorine atom because argon has one more electron shell (3 shells) compared to chlorine which has 3 shells but argon has a complete outer shell with 8 electrons while chlorine has 7.
(c) Ionisation potential increases across a period because as we move from left to right, the nuclear charge increases while the number of electron shells remains the same. The increased nuclear attraction makes it harder to remove electrons.
In simple words: Anions are bigger because extra electrons push each other apart. Argon is bigger because it has more electrons in the same shell. Ionization energy increases across periods because more protons pull electrons tighter.

📝 Teacher's Note: Use balloon analogies - more air (electrons) makes balloon bigger (anion). Show how increasing nuclear charge across period pulls electrons closer.

🎯 Exam Tip: Remember electron-electron repulsion makes anions larger. Across periods, increasing nuclear charge overcomes constant shielding, so ionization energy increases.

Question. (a) Anion is formed by the gain of electrons. Thus the numbers of electrons are more than protons. The effective positive charge in the nucleus is less, so less inward pull is experienced. Hence the size expands. So the size of an atom is greater than the size of parent atom.
(b) Since Argon has stable octet configuration, so due to the inter- electronic repulsions the effect of nuclear pull over the valence shell electrons cannot be seen which results in the bigger size.
(c) Since size of Bromine is bigger than chlorine, so it becomes more difficult for Br atoms to attract electrons. Thus, Cl is more reactive than Br.
Answer: (a) Anions are larger than their parent atoms because gaining electrons increases electron-electron repulsion while nuclear charge remains the same, causing the electron cloud to expand. (b) Argon's complete octet creates strong electron repulsions that overcome nuclear attraction, making it larger than expected. (c) Chlorine is more reactive than bromine because smaller atoms can attract electrons more effectively due to stronger nuclear pull.

📝 Teacher's Note: Use visual models or diagrams to show how adding electrons to an atom makes it "puff up" like a balloon. Compare it to a crowded room where people push each other away.

🎯 Exam Tip: Always mention "nuclear charge remains constant" when explaining ionic size changes — this is a key concept examiners look for.

Question 17. Which element has:
(a) two shells, both of which are completely filled with electrons?
(b) the electronic configuration 2, 8, 3?
(c) a total of three shells with five electrons in its valence shell?
(d) a total of four shells with two electrons in its valence shell?
(e) twice as many electrons in its second shell as in its first shell?
Answer: (a) Neon (b) Aluminum (c) Phosphorus (d) Calcium (e) Carbon
In simple words: These questions test your ability to count electrons in different shells and match them to specific elements on the periodic table.

📝 Teacher's Note: Have students practice drawing electron configurations for the first 20 elements until they can do it automatically. This builds foundation for all periodic trends.

🎯 Exam Tip: When asked about electron configurations, always count total electrons to verify your answer matches the atomic number.

Question 18. Name
(a) An alkali metal in period 3 and halogen in period 2.
(b) The noble gas with 3 shells.
(c) The non-metals present in period 2 and metals in period 3.
(d) The element of period 3 with valency 4
(e) The element in period 3 which does not form oxide
(f) The element of lower nuclear charge out of Be and Mg.
(g) Which has higher E.A. fluorine or Neon.
(h) Which has maximum metallic character Na, Li or K.
Answer: (a) Na and F (b) Argon (c) C, N, O and F are non-metals present in period 2 while Na, Mg and Al are metals in period 3. (d) Silicon (e) Argon (f) Mg (g) Fluorine (h) K
In simple words: This tests your knowledge of element families and their positions in the periodic table.

📝 Teacher's Note: Create a colorful periodic table chart highlighting different element families — visual memory works better than just reading names.

🎯 Exam Tip: Remember that noble gases don't form compounds easily, so they won't have electron affinity values like other elements.

Question 19. Chorine in the periodic table is surrounded by the elements with atomic number 9, 16, 18 and 35.
(a) Which of these have physical and chemical properties resembling chlorine.
(b) Which is more electronegative than chlorine
Answer: (a) Element with atomic number 9 and 35 (b) Element with atomic number 9.
In simple words: Elements in the same group (family) have similar properties, and electronegativity increases as you go up a group.

📝 Teacher's Note: Explain that atomic numbers 9 and 35 are fluorine and bromine — both halogens like chlorine. This reinforces the concept of element families.

🎯 Exam Tip: When comparing electronegativity, remember it increases going up and to the right in the periodic table.

Question 20. (a) State the number of elements in periods 1, Periods 2, and Period 3, of the periodic table.
(b) name the elements in period 1.
(c) What is the common feature of the electronic configuration of the elements at the end of period 2, and period 3?
(d) if an element is in group 17, it is likely to be ………….. (Metallic / non-metallic) in character while with one electron in its outermost energy level (shell), then it is likely to be ……………(Metallic / Non-metallic)
Answer: (a) Period 1 has 2 elements while period 2 and period 3 have 8 elements each. (b) Hydrogen and helium (c) The elements at the end of period 2 and Period 3 have 8 electrons in its outermost shell. (d) if an element is in group 17, it is likely to be Non metallic (Metallic / non-metallic) in character while with one electron in its outermost energy level (shell), then it is likely to be metallic (Metallic / Non-metallic)
In simple words: The first period is special with only 2 elements, while periods 2 and 3 each have 8 elements ending with noble gases that have complete outer shells.

📝 Teacher's Note: Use the phrase "2-8-8" to help students remember the pattern of elements in the first three periods.

🎯 Exam Tip: Group 17 elements (halogens) are always non-metals, while Group 1 elements (with one outer electron) are always metals.

Question 21. First ionization enthalpy of two elements X and Y are 500KJ mol⁻¹ and 375KJ mol⁻¹ respectively. Comment about their relative position in a group as well as in a period.
Answer: Position in a group: X and Y Position in a period: Y and X
In simple words: Element X has higher ionization energy, so it's either above Y in a group or to the right of Y in a period.

📝 Teacher's Note: Explain that higher ionization energy means it's harder to remove electrons — this happens when atoms are smaller or have more nuclear pull.

🎯 Exam Tip: Remember the trends: ionization energy decreases down a group and increases across a period from left to right.

Question 22. A metal M forms as oxide having the formula M₂O₃. It belongs to third period. Write the atomic number and valency of the metal.
Answer: Period no. = no. of shells, so n = 3 From the formula M₂O₃ its valency is 3. Since it is a metal, its valence shell has 3 electrons. So its electronic configuration is 2, 8, 3 Atomic number =13 Hence the metal is Aluminum with valency 3.
In simple words: From the formula M₂O₃, we can figure out the metal has valency 3, and being in period 3 means it has 3 shells, leading us to aluminum.

📝 Teacher's Note: Show students how to work backwards from chemical formulas to determine valency — this is a crucial skill for compound formation.

🎯 Exam Tip: Always state the electronic configuration when asked for atomic number — it shows your complete understanding.

Question 23. Explain why are the following statements not correct:
(a) All groups contain metals and non metals.
(b) Atoms of elements in the same group have the same number of electron(s)
(c) Non- metallic character decreases across a period with increase in atomic number
(d) Reactivity increases with atomic number in a group as well as in a period.
Answer: (a) Since the elements in a group have same number of valence electrons, they can either contain metals or non-metals like alkali and alkaline metals have only metals whereas halogens are non-metals. (b) No two elements have the same number of electrons instead atoms of the same elements in the same group have the same number of valence electrons. (c) Non-metals have the tendency to gain electrons to attain stable configuration and therefore are said to be electronegative. As we move from left to right the increase in atomic number and decrease in size results in a greater nuclear pull. As a result the non-metallic character increases across a period. (d) On moving from left to right in a period, the reactivity first decreases and then increases since the tendency to lose electrons first decreases on going from left to right and then from P to Cl, tendency to gain electrons increases, so reactivity increases then. In case of a group, reactivity increases on going down since the tendency to lose electrons increases but for non-metals, reactivity decreases on going down the group as the tendency to gain electrons decreases down the group.
In simple words: These statements contain common misconceptions about periodic trends that students often make.

📝 Teacher's Note: Use these incorrect statements as discussion starters — ask students to identify what's wrong before giving the explanation.

🎯 Exam Tip: When explaining why statements are incorrect, always provide the correct trend and reasoning — partial answers lose marks.

Question 24. Arrange the following in order of increasing radii:
(a) CI⁻, CI
(b) Mg²⁺, Mg, Mg⁺
(c) N, O, P
Answer: (a) Cl < Cl⁻ (b) Mg²⁺ < Mg⁺ < Mg (c) O < N < P
In simple words: Anions are larger than atoms, cations are smaller, and size increases down a group while decreasing across a period.

📝 Teacher's Note: Draw circles of different sizes to show how atoms change size when they gain or lose electrons — visual representation helps retention.

🎯 Exam Tip: Remember that losing electrons makes atoms smaller (more positive charge pulling fewer electrons) while gaining electrons makes them larger.

Question 25. Which element from the following has the highest ionization energy? Explain your choice.
(a) P, Na, CI
(b) F, O, Ne
(c) Ne, He, Ar
Answer: (a) Cl Metals have low ionisation energy and non-metals have high ionisation energy. Also, across the period, ionisation energy tends to increase. The elements P, Na and Cl belong to the third period. Na - Group 1, P - Group 15 and Cl - Group 17. (b) Ne Inert gases have zero electron affinity because of their stable electronic configuration. (c) He Ionisation energy decreases with an increase in the atomic size, i.e. it decreases as one moves down a group. Ne, He and Ar are inert gases. He - Period 1, Ne - Period 2 and Ar - Period 3.
In simple words: Ionization energy is highest for small atoms with strong nuclear pull, and noble gases have very high values due to their stable configurations.

📝 Teacher's Note: Emphasize that noble gases have the highest ionization energies in their periods because of their complete electron shells.

🎯 Exam Tip: Always justify your choice by mentioning periodic trends — this shows you understand the underlying principles.

Question 26. Choose the correct answer.
(a) An element in period 3 whose electron affinity is zero
(i) Sulphur
(ii)Sodium
(iii) Neon
(iv) Argon
(b) An alkaline earth metal
(i) Lead
(ii) potassium
(iii) calcium
(iv) Copper
(c) An element with highest ionization potential
(i) Calesium
(ii) Fluorine
(iii) Helium
(iv) Neon
Answer: (a) (iv) Argon (b) (iii) Calcium (c) (iii) Helium
In simple words: Noble gases don't accept electrons easily, alkaline earth metals are in Group 2, and helium has the highest ionization energy of all elements.

📝 Teacher's Note: Create mnemonics for element families — "Be Mg Ca Sr Ba Ra" for alkaline earth metals helps students remember Group 2.

🎯 Exam Tip: Noble gases have zero electron affinity because their electron shells are already complete and stable.

Question 2003. The table given below represents the first three periods. Study the table and answer the question as given below
(a) Write the formula of the sulphate of the element with atomic number 13
(b) what type of bonding will be present in the oxide of the element with atomic number 17?
(c) Which feature of the atomic structure accounts for the similarities in the chemical properties of the elements in group 7A of the periodic table?
(d) Name the element which has the highest ionization potential.
(e) How many electrons are present in the valency shell of the element with atomic number 18?
(f) What is the name given to the energy released when an atom in its isolated gaseous state accepts an electron to form an anion?
Answer: Based on the periodic table shown: (a) Al₂(SO₄)₃ - aluminum forms Al³⁺ ions, so its sulfate formula requires three sulfate ions to balance two aluminum ions. (b) Covalent bonding - element 17 is chlorine, a non-metal, so Cl₂O₇ will have covalent bonds between non-metal atoms. (c) Same number of valence electrons - all Group 7A elements have 7 valence electrons, giving them similar chemical behavior. (d) Helium (atomic number 2) - it has the highest ionization potential as the smallest noble gas. (e) 8 electrons - atomic number 18 is argon, which has a complete octet in its valence shell. (f) Electron affinity - this is the energy change when an isolated gaseous atom gains an electron.
In simple words: This question tests your ability to use the periodic table to predict chemical formulas, bonding types, and atomic properties.

📝 Teacher's Note: Use the provided periodic table as a reference tool — teach students to extract atomic numbers, periods, and groups from such diagrams.

🎯 Exam Tip: When writing chemical formulas, always check that positive and negative charges balance to give a neutral compound.

Question. Fill in the blanks: The atomic size ………. as we move from left to right across the periods, because the ………. increases but the …………… remains the same
Answer: The atomic size decreases as we move from left to right across the periods, because the atomic number increases but the number of shells remains the same
In simple words: As we go across a period, atoms get smaller because more protons pull the electrons closer, even though the number of electron shells stays the same.

📝 Teacher's Note: Use the analogy of a magnet getting stronger - more protons mean stronger pull on electrons. Draw atomic diagrams to show same shell number but increasing nuclear charge.

🎯 Exam Tip: Remember the key phrase "atomic number increases but number of shells remains same" - this exact wording often appears in mark schemes.

Question. The electro negativities (according to pauling) of the elements in periodic table are as follows with the elements arranged in alphabetical order: Al Cl Mg Na P S Si 1.5 3.0 1.2 0.9 2.1 2.5 1.8
(a) Arrange the elements in the order in which they occur in the periodic table from left to right. (The group 1 element first, followed by the group 2 element and so on, up to group 7)
(b) Choose the word or phrase from the brackets which correctly completes each of the following statements:
(i) The element below sodium in the same group would be expected to have a………. (lower/higher) electro-negativity than sodium and the element above chlorine would be expected to have a ……. (lower/ higher) ionization potential than chlorine.
(ii) On moving from left to right in a given period, the number of shells (remains the same/ increases/ decreases).
(iii) On moving down a group, the number of valence electrons (remains the same/ increases/ decreases).
Answer:
(a) Na, Mg, Al, Si, P, S, Cl
(b) (i) The element below sodium in the same group would be expected to have a Lower (lower/higher) electro-negativity than sodium and the element above chlorine would be expected to have a higher (lower/ higher) ionization potential than chlorine.
(ii) remains the same
(iii) remains the same
In simple words: Elements are arranged by groups and periods. As you go down a group, electronegativity decreases but valence electrons stay the same. As you go up a group, ionization potential increases.

📝 Teacher's Note: Create a blank periodic table outline and have students place these elements in correct positions. Emphasize that group number = valence electrons for main group elements.

🎯 Exam Tip: Always arrange elements by group number first (1, 2, 13, 14, 15, 16, 17) when asked to show periodic table order. Remember trends: electronegativity decreases down, increases across.

Question. Parts (a) to (e) refer to changes In the properties of elements on moving from left to right across a period of the periodic table. For each property, choose the correct answer.
(a) The non-metallic character of the elements:
(i) decreases
(ii) increases
(iii) remains the same
(iv) depends on the period
(b) The electronegativity:
(i) depends on the number of valence electrons
(ii) remains the same
(iii) decreases
(iv) increases
(c) The ionization potential:
(i) goes up and down
(ii) decreases
(iii) increases
(iv) remains the same
(d) The atomic size:
(i) decreases
(ii) increases
(iii) remains the same
(iv) sometimes increases and sometimes decreases
(e) The electron affinity of the elements in groups 1 to 7:
(i) goes up and then down
(ii) decreases and then increases
(iii) increases
(iv) decreases
Answer:
(a) Increases
(b) Increases
(c) Increases
(d) Decreases
(e) Increases
In simple words: As you move left to right across a period, atoms become more non-metallic, hold electrons more tightly, get smaller, and want electrons more strongly.

📝 Teacher's Note: Use the trend Na → Mg → Al → Si → P → S → Cl to show how properties change from metallic to non-metallic. Visual aids showing atomic size decrease work well.

🎯 Exam Tip: For across-period trends, remember "increases" for most properties except atomic size which "decreases". The pattern is consistent for periods 2 and 3.

Question. The elements of one short period of the periodic table are given below in order from left to right: Li Be B C O F Ne
(a) To which period do these elements belong?
(b) One element of this period is missing. Which is the missing element and where should it be placed?
(c) Which one of the elements in this period shows the property of catenation?
(d) Place the three elements fluorine, beryllium and nitrogen in the order of increasing electronegativity.
(e) Which one of the above elements belongs to the halogen series?
Answer:
(a) Period 2
(b) Nitrogen (N), between carbon and oxygen
(c) Carbon
(d) Be < N < F
(e) Fluorine
In simple words: These are Period 2 elements. Nitrogen is missing between carbon and oxygen. Carbon can form chains with itself. Electronegativity increases across the period. Fluorine is a halogen.

📝 Teacher's Note: Point out that catenation is carbon's unique ability to bond with itself repeatedly, forming chains and rings. This property makes organic chemistry possible.

🎯 Exam Tip: Remember Period 2 has 8 elements (Li to Ne). Catenation questions always refer to carbon. Halogens are Group 17 elements (F, Cl, Br, I).

Question. A group of elements in the periodic table are given below (boron is the first member of the group and thallium is the last). Boron, Aluminium, Gallium, Indium, Thallium. Answer the following questions in relation to the above group of elements:
(a) Which element has the most metallic character?
(b) Which element would be expected to have the highest electronegativity?
(c) If the electronic configuration of aluminium is 2, 8, 3, how many electrons are there in the outer shell of thallium?
(d) The atomic number of boron is 5. Write chemical formula of the compound formed when boron reacts with chlorine.
(e) Will the elements in the group to the right of this boron group be more metallic or less metallic in character? Justify your answer.
Answer:
(a) Thallium has the most metallic character since metallic character increases down the group
(b) Boron has the highest electronegativity since it has the smallest size in the group
(c) 3. Since all the elements in a group have same number of valence electrons
(d) \( BCl_3 \)
(e) The elements in the group to the right of boron group would be less metallic as with the decrease in size and increase in atomic number, it will be more difficult for them to lose electrons
In simple words: Going down a group, elements become more metallic. Boron is most electronegative as it's smallest. All group members have 3 valence electrons. Boron forms \( BCl_3 \) with chlorine. Elements to the right are less metallic.

📝 Teacher's Note: Emphasize that Group 13 elements all have 3 valence electrons. Show how metallic character relates to ease of losing electrons - larger atoms lose electrons more easily.

🎯 Exam Tip: For group trends, remember metallic character increases down the group. For valence electrons, all elements in the same group have identical numbers in outer shell.

Question. Select the correct answer from the choice A, B, C, D which are given. Write down only the letter corresponding to the correct answer. With reference to the variation of properties in the periodic table, which of the following is generally true?
A. Atomic size increases from left to right across a period.
B. Ionization potential increases from left to right across a period.
C. Electron affinity increases going down a group.
D. electro-negativity increases going down a group.
Answer: (B) Ionization potential increases from left to right across a period.
In simple words: As you move across a period, it becomes harder to remove electrons because the nuclear charge increases while the distance remains similar.

📝 Teacher's Note: Explain that ionization potential is the energy needed to remove an electron. Demonstrate with examples from Period 2 or 3 showing increasing values.

🎯 Exam Tip: Know the basic periodic trends - atomic size decreases across, ionization potential increases across, electronegativity increases across and decreases down.