GSEB Class 10 Science Solutions Chapter 5 Periodic Classification of Elements

GSEB Solutions

 

Question 1. Did Dobereiner's triads also exist in the columns of Newland's Octaves? Compare and find out.
Answer: Yes, they are present in Newland's octaves. These elements are Lithium, Sodium, and Potassium.
In simple words: Dobereiner's groups of three elements, called triads, can indeed be found within Newland's Octaves, specifically with Lithium, Sodium, and Potassium.

Exam Tip: To compare classification systems, identify common elements or patterns. For Dobereiner's triads and Newland's Octaves, look for groups of three elements with similar properties that also fit the octave pattern.

 

Question 2. What were the limitations of Dobereiner's classification?
Answer:

• Not all the elements known at that time could be categorized.
• Dobereiner managed to find just three triads among the elements.

So, this particular method was not very helpful.
In simple words: Dobereiner's classification had limits because it couldn't group all known elements, and he found only three triads in total, making the system less useful.

Exam Tip: When asked about limitations, focus on what the theory or system failed to explain or include, and how general or specific its applicability was.

 

Question 3. What were the limitations of Newlands' Law of Octaves?
Answer:

• Newlands' law applied only up to calcium; beyond that, every eighth element did not exhibit similar characteristics to the first.
• Newly found elements did not conform to this rule.
• Elements were arranged incorrectly; for instance, Cobalt and Nickel, which are dissimilar to halogens, were placed in the same category, while Iron, similar to Cobalt and Nickel, was kept apart.

In simple words: Newlands' law only worked up to calcium, new elements didn't fit, and some elements were grouped wrongly, making it inconsistent.

Exam Tip: Remember to list specific examples of elements that didn't fit Newlands' Octaves, like Co, Ni, and Fe, to show a deeper understanding.

 

Question 4. Use Mendeleev's Periodic Table to predict the formulae for the oxides of the following elements: K, C, Al, Si, Ba
Answer:

In simple words: By using Mendeleev's table, you can guess how elements like Potassium, Carbon, Aluminium, Silicon, and Barium will combine with oxygen to form their specific oxide formulas.

Exam Tip: Remember that the valency of an element (its combining capacity) determines the formula of its oxide. Mendeleev's table groups elements with similar valencies together, which helps in predicting compound formulas.

 

Question 5. Besides gallium, which other elements have since been discovered that were left by Mendeleev in his Periodic Table? (any two)
Answer: In addition to gallium, Germanium and Scandium have also been found.
In simple words: Besides gallium, Germanium and Scandium are two more elements that Mendeleev predicted but were discovered later.

Exam Tip: When answering questions about Mendeleev's predictions, recall 'Eka-aluminium' (Gallium), 'Eka-silicon' (Germanium), and 'Eka-boron' (Scandium) as key examples of his foresight.

 

Question 6. What were the criteria used by Mendeleev in creating his Periodic Table?
Answer:

1. Elements were arranged by their increasing atomic mass.
2. Elements within a group shared comparable properties.
3. The chemical structure of the oxides and hydrides each element formed.

In simple words: Mendeleev built his Periodic Table by ordering elements based on their atomic mass, grouping those with similar properties, and considering the formulas of their oxides and hydrides.

Exam Tip: Mendeleev's main criteria were atomic mass and chemical properties, especially the formulas of compounds with oxygen and hydrogen. Make sure to mention these two key points.

 

Question 7. Why do you think noble gases are placed in a separate group?
Answer: Noble gases are inert, they do not look like other elements, and all of them exhibit similar traits, so they are placed in a distinct group.
In simple words: Noble gases are in their own group because they are very unreactive, don't chemically resemble other elements, and all share similar properties.

Exam Tip: The key reason for noble gases' separate grouping is their full outermost electron shell, making them very stable and unreactive compared to other elements.

 

Question 8. How could the Modern Periodic Table remove various anomalies of Mendeleev's Periodic Table?
Answer: Modern Periodic Table is based on the atomic number of elements, therefore

1. The issue with isotopes was resolved since isotopes share the same atomic number.
2. The incorrect placement of Argon, Potassium, Cobalt, and Nickel was corrected.

In simple words: The Modern Periodic Table fixed Mendeleev's problems by using atomic number for organization. This sorted out isotope placement and corrected several elements' positions.

Exam Tip: The atomic number (number of protons) is the fundamental basis of the Modern Periodic Table, which resolved the anomalies of Mendeleev's table that was based on atomic mass.

 

Question 9. Name two elements you would expect to show chemical reactions similar to magnesium. What is the basis for your choice?
Answer: Calcium and Barium.

• Both these elements are part of the same group as magnesium.
• Barium and Calcium have the same number of valence electrons as Magnesium, and will display similar chemical characteristics.

In simple words: Calcium and Barium would react like magnesium because they are in the same group, meaning they have the same number of valence electrons and thus similar chemical properties.

Exam Tip: Elements in the same group of the Periodic Table have similar chemical properties because they possess the same number of valence electrons.

 

Question 10. Name:
(a) Three elements that have a single electron in their outermost shells.
(b) Two elements that have two electrons in their outermost shells.
(c) Three elements with filled outermost shells.
Answer:
(a) Li (2, 1), Na (2, 8, 1), K (2, 8, 8, 1)
(b) Be (2, 2), Mg (2, 8, 2)
(c) He (2), Ne (2, 8), Ar (2, 8, 8)
In simple words: For part (a), Lithium, Sodium, and Potassium have one electron in their outer shells. For part (b), Beryllium and Magnesium have two outer electrons. For part (c), Helium, Neon, and Argon have full outer shells.

Exam Tip: The number of valence electrons determines an element's group, and elements with filled outer shells are noble gases, which are stable.

 

Question 11.
(a) Lithium, sodium, potassium are all metals that react with water to liberate hydrogen gas. Is there any similarity in the atoms of these elements?
(b) Helium is an unreactive gas and neon is a gas of extremely low reactivity. What, if anything, do their atoms have in common?
Answer:
(a) All these metals are very reactive; they each possess one valence electron and can easily shed electrons to form positive ions.
(b) Helium and Neon both feature fully occupied outermost electron shells.
In simple words: (a) Lithium, sodium, and potassium all have one valence electron, which makes them very reactive and able to form positive ions. (b) Helium and neon are both unreactive because their outermost electron shells are completely full.

Exam Tip: Similar chemical properties stem from identical numbers of valence electrons. Full outermost shells explain why elements are unreactive (inert).

 

Question 12. In the Modern Periodic Table, which are the metals among the first ten elements.
Answer: Lithium and Beryllium are classified as metals.
In simple words: Among the first ten elements in the Modern Periodic Table, Lithium and Beryllium are considered metals.

Exam Tip: Remember that metals are typically found on the left side of the periodic table. For the first ten elements (H to Ne), hydrogen and helium are non-metals, lithium and beryllium are metals, and boron, carbon, nitrogen, oxygen, fluorine, and neon are non-metals.

 

Question 13. By considering their position in the Periodic Table, which one of the following elements would you expect to have maximum metallic characteristics? Ga, Ge, As, Se, Be.
Answer: Among the elements provided, Beryllium (Be) will display the highest metallic characteristics because it is positioned on the far left side of the Periodic Table.
In simple words: Beryllium (Be) has the most metallic character because it is furthest to the left in the Periodic Table among the given options.

Exam Tip: Metallic character increases as you move down a group and decreases as you move across a period from left to right. Elements on the far left (like Group 1 and 2) are most metallic.

 

In-Text Activities Solved

 

Activity 5.1
Answer:

1. Hydrogen is similar to alkali metals since it joins with halogens, oxygen, and sulfur to create compounds with chemical structures resembling metals.
2. Hydrogen also shows traits like halogens because it exists as two-atom molecules and combines with both metals and non-metals to create covalent compounds.
3. A specific position could not be given to hydrogen within Mendeleev's Periodic Table.

In simple words: Hydrogen's ability to form compounds like alkali metals and halogens, and its diatomic nature, made it hard to place consistently in Mendeleev's table.

Exam Tip: The dual nature of hydrogen (resembling both alkali metals and halogens) caused issues in early periodic classifications due to its unique electron configuration.

 

Activity 5.2
Answer: Chlorine possesses two isotopes, Chlorine-35 and Chlorine-37. Both these isotopes exhibit identical chemical characteristics, therefore both isotopes should occupy the same space.
In simple words: Chlorine-35 and Chlorine-37 are isotopes, and since they have the same chemical properties, they should occupy the same spot in the periodic table.

Exam Tip: Isotopes have the same atomic number but different mass numbers. They belong to the same element and thus share identical chemical behaviors.

 

Activity 5.3
Answer: The placement of Cobalt (Co) and Nickel (Ni) was fixed by arranging them according to their increasing atomic numbers in the Modern Periodic Table.

1. Isotopes were not factored in during the arrangement of elements in the Modern Periodic Table because elements were ordered by their increasing atomic number.
2. An element's atomic number cannot be 1.5, as atomic number represents the quantity of protons in an atom, which must always be a full number.
3. Hydrogen ought to be positioned in Group-I of the Modern Periodic Table.

In simple words: Cobalt and Nickel's positions were sorted by atomic number. Isotopes aren't an issue for classification by atomic number, and atomic numbers must be whole numbers. Hydrogen belongs in Group 1.

Exam Tip: The Modern Periodic Table's organization by atomic number resolved ambiguities that arose from Mendeleev's mass-based classification, particularly for isotopes and specific element pairs.

 

Activity 5.4
Answer:

In simple words: This table shows the properties of Group I elements, including their atomic number, electron configuration, valence electrons, valency, and relative atomic size.

Exam Tip: Understand that elements in Group I (alkali metals) all have one valence electron, leading to a valency of 1 and similar chemical reactivity.

 

Activity 5.5
Answer:

They contain same number of shells.
In simple words: This table illustrates the atomic number, electron configurations, valence electrons, valency, and relative atomic sizes for elements in Period II. All elements in a given period have the same number of electron shells.

Exam Tip: In a period, the number of electron shells remains constant. Trends like increasing valence electrons, varying valency, and decreasing atomic size are observed from left to right across a period.

 

Activity 5.6
Answer: An element's valency, derived from its electron arrangement, is determined by checking the quantity of electrons in its outermost shell. If this count exceeds 4, then subtract 4 from that number to find the valency).

In simple words: Valency can be figured out from the number of electrons in the outermost shell. If this number is more than four, subtract four from it. The table demonstrates this for Magnesium and Sulphur. Valency changes across a period but stays the same within a group.

Exam Tip: Valency is often 8 minus the number of valence electrons for non-metals (if valence electrons > 4) and equals the number of valence electrons for metals (if valence electrons < 4).

 

Activity 5.7
Answer: Atomic radii of the second period elements

Given:
B (88), Be (111), O (66), N (74), Li (152), C (77)

In decreasing order of their atomic radii
Li (152), Be (111), B (88), C (77), N (74), O (66)

Lithium possesses the largest atom size, while Oxygen has the smallest. As you move from left to right within a period, the atomic size consistently gets smaller.
In simple words: For second-period elements, Lithium has the biggest atomic radius, and Oxygen has the smallest. Atomic size shrinks as you go from left to right across a period.

Exam Tip: Remember the trend: atomic radius generally decreases across a period from left to right due to increased nuclear charge pulling electrons closer to the nucleus, even with the same number of shells.

 

Activity 5.8
Answer: Group I elements.

Increases

The size of atoms grows as you move down a group, due to the addition of new electron shells.
In simple words: For Group I elements, atomic size consistently increases from Lithium to Cesium because new electron shells are added as you go down the group.

Exam Tip: Atomic size increases down a group because each successive element adds a new electron shell, which increases the distance between the outermost electrons and the nucleus.

 

Activity 5.9
Answer: Elements of III period

Metals typically appear on the left side of the periodic table, while non-metals are found on the right side.
In simple words: Elements in Period III show a clear distinction: metals like Na, Mg, Al are on the left, while non-metals like P, S, Cl, Ar are on the right, with Si acting as a metalloid.

Exam Tip: Understand that the diagonal line from Boron (B) to Astatine (At) separates metals from non-metals, with elements near this line being metalloids.

 

Activity 5.10
Answer: The likelihood of losing electrons increases within a group.
The likelihood of losing electrons decreases across a period.

L \( \leftarrow \) R
readily lose \( e^- \) \( \rightarrow \) readily gain \( e^- \)
In simple words: The tendency to lose electrons goes up as you move down a group but goes down as you move from left to right across a period.

Exam Tip: Metallic character, which is related to the tendency to lose electrons, follows the same trend: increases down a group and decreases across a period.

 

Activity 5.11
Answer:

• The ability to acquire electrons grows when moving across a period.
• The ability to acquire electrons lessens when moving down a group.

In simple words: The chance of an element gaining electrons goes up across a period, but it goes down as you move deeper into a group.

Exam Tip: Non-metallic character, which is linked to the tendency to gain electrons, increases across a period and decreases down a group.

 

Gujarat Board Class 10 Science Periodic Classification of Elements Textbook Questions and Answers

 

Question 1. Which of the following statements is not a correct statement about the trends when going from, left to right across the periods of the Periodic Table?
(a) The elements become less metallic in nature.
(b) The number of valence electrons increases.
(c) The atoms lose their electrons more easily.
(d) The oxides become more acidic.
Answer: (c) The atoms lose their electrons more easily.
In simple words: Moving from left to right across a period, elements tend to hold onto their electrons more tightly, making it harder for atoms to give them away.

Exam Tip: Remember that metallic character decreases and non-metallic character increases across a period, which means elements are less likely to lose electrons (and more likely to gain them).

 

Question 2. Element X forms a chloride with the formula \(XCl_2\), which is a solid with a high melting point. X would most likely be in the same group of the Periodic Table as -
(a) Na
(b) Mg
(c) Al
(d) Si
Answer: (b) Mg.
In simple words: Element X creates a chloride with two chlorine atoms, similar to how Magnesium (Mg) forms \(MgCl_2\). This indicates X belongs to the same group as Mg, which usually forms \(XCl_2\) compounds.

Exam Tip: The formula \(XCl_2\) indicates that element X has a valency of 2. Elements with a valency of 2 belong to Group 2, like Magnesium (Mg).

 

Question 3. Which element has -
(a) two shells, both of which are completely filled with electrons?
(b) electronic configuration 2, 8, 2?
(c) a total of three shells, with four electrons in its valence shell?
(d) a total of two shells, with three electrons in its valence shell?
(e) twice as many electrons in its second shell as in its first shell?
Answer:
(a) Neon (Ne), with electron configuration (2, 8).
(b) Magnesium (Mg), with electron configuration (2, 8, 2).
(c) Silicon (Si), with electron configuration (2, 8, 4).
(d) Boron (B), with electron configuration (2, 3).
(e) Carbon (C), with electron configuration (2, 4).
In simple words: This question asks you to identify elements based on their electron shells and electron count. For instance, 'two shells, both filled' means all electron slots are taken, like in Neon.

Exam Tip: To identify elements from their electronic configuration, sum the electrons to find the atomic number. The number of shells indicates the period, and valence electrons indicate the group (or group 10 + valence electrons for p-block elements).

 

Question 4.
(a) What property do all elements in the same column of the Periodic Table as boron have in common?
(b) What property do all elements in the same column of the Periodic Table as fluorine have in common?
Answer:
(a) All other elements in that column possess the same number of valence electrons, and their valency is 3.
(b) All these elements are non-metals; they each have seven valence electrons and a valency of 1. They all acquire electrons to create negative ions.
In simple words: (a) Elements in Boron's column share the same number of outer electrons and a valency of 3. (b) Elements in Fluorine's column are all non-metals with seven outer electrons, a valency of 1, and form negative ions.

Exam Tip: Elements in the same group (column) have similar chemical properties because they have the same number of valence electrons, leading to the same valency and similar reactivity patterns.

 

Question 5. An atom has electronic configuration 2, 8, 7.
(a) What is the atomic number of this element?
(b) To which of the following elements would it be chemically similar? (Atomic numbers are given in parentheses) N (7) F(9) P(15) Ar (18)
Answer:
(a) The element's atomic number is 17.
(b) Fluorine (F) with an atomic number of 9 and electron configuration (2, 7) will exhibit similar chemical characteristics to the specified element.
In simple words: An atom with electron configuration 2, 8, 7 has an atomic number of 17. It will be chemically similar to Fluorine (F) because both have 7 valence electrons.

Exam Tip: The sum of electrons in the configuration gives the atomic number. Chemical similarity depends on the number of valence electrons, as these determine reactivity.

 

Question 6. The position of three elements A, B and C in the Periodic Table are shown below:

(a) State whether A is a metal or non-metal.
(b) State whether C is more reactive or less reactive than A.
(c) Will C be larger or smaller in size than B?
(d) Which type of ion, cation or anion, will be formed by element A?
Answer:
(a) 'A' represents a non-metal.
(b) 'C' shows lower reactivity compared to 'A'.
(c) 'C' has a smaller physical dimension than 'B'.
(d) 'A' will create a negatively charged ion, known as an Anion.
In simple words: (a) Element A is a non-metal. (b) C is less reactive than A. (c) C is smaller than B. (d) A will form a negatively charged ion (anion).

Exam Tip: Remember periodic trends: Non-metals are on the right. Reactivity of non-metals decreases down a group and increases across a period. Atomic size decreases across a period and increases down a group. Non-metals gain electrons to form anions.

 

Question 7. Nitrogen (atomic number 7) and phosphorous (atomic number 15) belong to group 15 of the Periodic Table. Write the electronic configuration of these two elements. Which of these will be more electronegative? Why?
Answer: Nitrogen, with an atomic number of 7, has an electron configuration of 2, 5.
Phosphorus, with an atomic number of 15, has an electron configuration of 2, 8, 5
Nitrogen, possessing two electron shells, will be more electronegative as it can readily obtain electrons because its smaller atomic size means the nucleus exerts a stronger pull on incoming electrons, forming a negative ion.
In simple words: Nitrogen (2, 5) and Phosphorus (2, 8, 5) are in Group 15. Nitrogen is more electronegative because its smaller atomic size means the nucleus pulls electrons more strongly, making it easier to gain them.

Exam Tip: Electronegativity increases across a period and decreases down a group. Nitrogen is above phosphorus, has fewer shells, and thus its nucleus has a stronger pull on valence electrons, making it more electronegative.

 

Question 8. How does the electronic configuration of an atom relate to its position in the Modern Periodic Table?
Answer: An element's placement relies on its electronic configuration. The quantity of electron shells matches the period number. The number of valence electrons determines the group number; elements with one valence electron belong to Group 1, and those with two valence electrons belong to Group 2.
In simple words: An atom's electron configuration dictates its place in the Periodic Table. The number of electron shells tells you the period, and the valence electrons tell you the group.

Exam Tip: Remember: the number of principal energy shells (occupied shells) corresponds to the period number, and the number of valence electrons typically corresponds to the group number (for representative elements).

 

Question 9. In the Modern Periodic Table, calcium (atomic number 20) is surrounded by elements with atomic numbers 12, 19, 21 and 38. Which of these have physical and chemical properties resembling calcium.
Answer: Calcium's atomic number is 20.
Its electron configuration is 2, 8, 8, 2
Elements with atomic number 12 (configuration 2, 8, 2) and atomic number 38 (configuration 2, 8, 18, 8, 2) will be similar to calcium because they all possess the same valence electrons and show identical chemical characteristics.
In simple words: Calcium (atomic number 20) has electron configuration 2, 8, 8, 2. Elements with atomic numbers 12 (Magnesium) and 38 (Strontium) will share similar properties with calcium because they all have 2 valence electrons.

Exam Tip: Elements with similar chemical properties are found in the same group because they have the same number of valence electrons. Identify elements in Calcium's group (Group 2) by their electron configuration.

 

Question 10. Compare and contrast the arrangement of elements in Mendeleev's Periodic Table and the Modern Periodic Table.
Answer:
Mendeleev's Periodic Table:

• It contains 8 groups and 6 periods.
• Transition elements are not set apart.
• Inert gases were absent from this table.
• Lanthanides and Actinides were not included.
• An element's position, meaning its group and period number, could not be anticipated.
• Elements are ordered based on their atomic mass.

Modem Periodic Table:

• It features 18 groups and 7 periods.
• Transition elements are allocated a distinct position.
• Inert gases are included in a distinct group.
• Lanthanides and Actinides are located at the lower part of the Periodic Table.
• The group number and period number can be determined from an element's electron configuration.
• Elements are ordered based on their atomic number.

In simple words: Mendeleev's table had fewer groups and periods, didn't separate transition elements or include inert gases, and was based on atomic mass. The Modern Table has more groups and periods, separates transition elements and includes inert gases, and is based on atomic number, making element prediction easier.

Exam Tip: Highlight the fundamental difference: Mendeleev used atomic mass, leading to anomalies; the Modern table uses atomic number, which resolves these anomalies and provides a more systematic arrangement.

 

Additional Important Questions and Answers

 

Very Short Answer Type Questions

 

Question 1. Give one example of Dobereiner's Triad.
Answer: Lithium (7), Sodium (23), and Potassium (39) form a Dobereiner's Triad.
In simple words: Lithium, Sodium, and Potassium make up a Dobereiner's Triad.

Exam Tip: A Dobereiner's Triad consists of three elements with similar chemical properties where the atomic mass of the middle element is approximately the average of the other two.

 

Question 2. How many triads could Dobereiner identify from the existing elements then?
Answer: Dobereiner was only able to recognize three triads.
In simple words: Dobereiner could identify only three triads.

Exam Tip: Knowing the exact number of triads identified by Dobereiner demonstrates a precise understanding of his classification system's limitations.

 

Question 3. What is the limitation of Dobereiner triads?
Answer: He was unable to organize all elements into triads that shared identical chemical characteristics.
In simple words: Dobereiner's triads couldn't classify all known elements into groups with similar properties.

Exam Tip: The main limitation was its narrow applicability; only a few elements could be grouped into triads, leaving many unclassified.

 

Question 4. What was the basis of classification of elements made by Newlands?
Answer: Newlands ordered the elements according to their increasing atomic masses.
In simple words: Newlands arranged elements based on their increasing atomic masses.

Exam Tip: Newlands' Law of Octaves was based on atomic mass, similar to musical octaves, where every eighth element had similar properties.

 

Question 5. Give two limitations of Newlands' law of Octaves.
Answer: Two limitations of Newlands' law of octaves are:

• This law was only valid up to calcium.
• To make elements fit into his table, Newlands put some elements in incorrect positions and placed dissimilar elements together in the same groups.

In simple words: Newlands' law was limited because it only worked for lighter elements up to calcium, and he sometimes forced elements into incorrect slots, grouping unlike ones together.

Exam Tip: Always remember the calcium limit and the issue of grouping dissimilar elements (e.g., Co, Ni with halogens) as key drawbacks of Newlands' Octaves.

 

Question 6. On what basis did Mendeleev classified the element?
Answer: Mendeleev organized the elements based on their rising atomic masses and their comparable chemical characteristics.
In simple words: Mendeleev classified elements using their increasing atomic masses and shared chemical properties.

Exam Tip: Mendeleev's classification was primarily based on atomic mass and recurring chemical properties, especially how elements reacted with oxygen and hydrogen.

 

Question 7. Which two chemical properties were considered by Mendeleev for grouping of elements?
Answer: The two chemical properties are:

• The type of compounds elements formed with oxygen.
• The type of compounds elements formed with hydrogen.

In simple words: Mendeleev used how elements combined with oxygen and hydrogen to group them.

Exam Tip: Mendeleev observed the formulas of oxides and hydrides to understand an element's valency and chemical nature, which was crucial for his periodic table.

 

Question 8. State Mendeleev's Periodic Law.
Answer: An element's characteristics repeat regularly, depending on its atomic mass.
In simple words: Mendeleev's Periodic Law states that the properties of elements periodically repeat when arranged by increasing atomic mass.

Exam Tip: This is a definition you must memorize. The key phrase is "periodic function of their atomic masses."

 

Question 9. Define 'groups and periods'.
Answer: The upright columns in a Periodic Table are known as groups, and the horizontal rows are referred to as periods.
In simple words: Groups are the vertical columns in the Periodic Table, and periods are the horizontal rows.

Exam Tip: Groups indicate similar chemical properties (same valence electrons), while periods indicate the number of electron shells (elements in the same period have the same number of shells).

 

Question 10. What is the formula of oxide and hydride of Group I elements?
Answer: The oxide formula is \(R_2O\), and the hydride formula is \(RH\). Here, 'R' stands for an element from Group I.
In simple words: Group I elements form oxides with the formula \(R_2O\) and hydrides with the formula \(RH\), where R is the Group I element.

Exam Tip: Group I elements (alkali metals) have a valency of 1. Oxygen has a valency of 2, and hydrogen has a valency of 1, which explains these simple formulas.

 

Question 11. Name three elements discovered later, which filled gaps left by Mendeleev for them.
Answer: Scandium, Gallium, and Germanium.
In simple words: Scandium, Gallium, and Germanium were later found and fit into the empty spaces Mendeleev had left in his Periodic Table.

Exam Tip: Mendeleev's ability to predict the existence and properties of these elements (originally named Eka-boron, Eka-aluminium, and Eka-silicon) was a major success of his periodic table.

 

Question 12. What are isotopes?
Answer: Isotopes are atoms of the same chemical element that share the same atomic number but possess different mass numbers.
Example:
\(^{12}_{6}C\); \(^{14}_{6}C\); \(^{1}_{1}H\), \(^{2}_{1}H\), \(^{3}_{1}H\)
In simple words: Isotopes are versions of the same element that have the same number of protons but different numbers of neutrons, so their atomic mass is different.

Exam Tip: The key identifiers for isotopes are the same atomic number (bottom left) but different mass numbers (top left). The number of protons defines the element, while the number of neutrons defines the specific isotope.

 

Question 13. How many groups and periods are present in the Modern Periodic Table?
Answer: The Modern Periodic Table comprises 18 groups and 7 periods.
In simple words: The Modern Periodic Table has 18 vertical columns called groups and 7 horizontal rows called periods.

Exam Tip: Remember these numbers for the structure of the Modern Periodic Table. Groups run vertically, periods horizontally.

 

Question 14. What is the location of metals and non-metals in the Modern Periodic Table?
Answer: Metals are situated on the left portion, while non-metals are located on the right portion of the Periodic Table.
In simple words: In the Periodic Table, metals are on the left side, and non-metals are on the right side.

Exam Tip: A zigzag line (metalloids) typically separates metals from non-metals, running from Boron (B) to Astatine (At) on the periodic table.

 

Question 15. State Modern Periodic Law.
Answer: The characteristics of elements show a regular, repeating pattern based on their atomic number.
In simple words: The Modern Periodic Law states that the properties of elements are a periodic function of their atomic number.

Exam Tip: Memorize this definition. The key difference from Mendeleev's Law is the basis on "atomic number" instead of "atomic mass."

 

Question 16. In Modern Periodic Table what is common among all the elements in a group.
Answer: All elements within the same group display identical valence electrons and similar chemical characteristics.
In simple words: All elements in the same group have the same number of valence electrons and thus similar chemical properties.

Exam Tip: The number of valence electrons dictates an element's chemical behavior, which is why elements in the same group behave similarly.

 

Question 17. Fluorine (F) atomic number = 9 and chlorine (Cl) atomic number =17 are placed in group number 17, what are the number of valence electrons present in them.
Answer: Fluorine has an atomic number of 9, with an electron configuration of 2, 7.
Chlorine has an atomic number of 17, with an electron configuration of 2, 8, 7.
Both elements possess 7 valence electrons.
In simple words: Fluorine (atomic number 9) and Chlorine (atomic number 17) are both in group 17, and both have 7 valence electrons.

Exam Tip: Elements in the same group (like Group 17, halogens) always have the same number of valence electrons, which is why they share similar chemical properties.

 

Question 18. What is common among all the elements present in one period?
Answer: All elements within the same period exhibit an identical number of electron shells; for example, all elements in Period 3 have three electron shells.
In simple words: All elements in the same period share the same number of electron shells; for instance, all Period 3 elements have three shells.

Exam Tip: The period number directly corresponds to the number of electron shells an element possesses, which is a key characteristic of its position in the periodic table.

 

Question 19. How many elements are present in first and second period?
Answer:

In simple words: The first period contains 2 elements and the second period contains 8 elements. This is based on the maximum number of electrons that can fit in each shell, which is given by the formula \(2n^2\).

Exam Tip: The number of elements in a period is determined by the maximum number of electrons that can be accommodated in the outermost shell, following the \(2n^2\) rule, where 'n' is the shell number.

 

Question 20. What is atomic size?
Answer: The radius of an atom, which means the space between its nucleus's centre and its outermost electron shell, is known as atomic size. This atomic radius is measured using picometre. For instance, 1 picometre equals \( 10^{-12} \) meters.
In simple words: Atomic size is how far an atom's outer electrons are from its center. We measure it in tiny units called picometres.

Exam Tip: Define both "atomic size" and its unit of measurement, picometre, to earn full credit.

 

Question 21. What happens to the size of atom as we move from left to right in a period.
Answer: The atom's size in a period gets smaller as we go from left to right.
In simple words: When you move from left to right across a row in the periodic table, atoms become smaller.

Exam Tip: Remember that atomic size trends horizontally across periods, typically decreasing from left to right due to increasing nuclear charge pulling electrons closer.

 

Question 22. How does the tendency to lose electrons will change in a period.
Answer: The likeliness to give up electrons will reduce when moving across a period, as the stronger nuclear pull on outer electrons grows.
In simple words: Moving left to right across a period, atoms are less likely to give away electrons because the nucleus pulls them more strongly.

Exam Tip: Relate the trend in electron loss to the effective nuclear charge: a higher charge means stronger attraction and a lower tendency to lose electrons.

 

Question 23. How do you think the tendency to lose electrons will change in a group?
Answer: As you go down a group, the force on outer electrons gets weaker because atomic size grows, so they can easily release electrons.
In simple words: When moving down a group, atoms are more likely to lose electrons because they are larger, and the outer electrons are not held as tightly.

Exam Tip: Contrast group trends with period trends: down a group, atomic size increases, weakening the pull on valence electrons and making them easier to lose.

 

Question 24. How was the anomaly in arrangement of elements in the Mendeleev's Periodic Table removed?
Answer: When elements were ordered by their rising atomic number, the problems with element arrangement in Mendeleev's Periodic Table were resolved.
In simple words: The problems in Mendeleev's table were solved by arranging elements by atomic number, not atomic mass.

Exam Tip: Remember that the Modern Periodic Table uses atomic number as its fundamental organizing principle, which successfully addressed the inconsistencies of Mendeleev's atomic mass-based table.

 

Question 25. What are noble gases/inert gases?
Answer: Elements that are unreactive, do not easily combine with other elements, and have their outermost electron shell full are known as inert gases or noble gases. For instance: He, Ne, Ar, Xe.
In simple words: Noble gases are elements that don't react much because their outer electron shells are already full. Examples include Helium and Neon.

Exam Tip: Focus on two key characteristics: chemical inertness (unreactive) and a full outermost electron shell, giving examples to illustrate your definition.

 

Question 26. Name two alkali metals present in Group I.
Answer: Some alkali metals include Li, Na, and K.
In simple words: Lithium, Sodium, and Potassium are all alkali metals found in Group I.

Exam Tip: Recall that Group I elements are alkali metals, known for their high reactivity, and memorize a few examples from this group.

 

Question 27. An element 'X' belongs to II group and 2nd period. Write the atomic number and name of element.
Answer: Since element 'X' is in Group II and Period 2, its electronic configuration reveals 2 electrons in the K shell and 2 in the L shell. Its atomic number is 4. The element is Beryllium.
In simple words: An element in Group II, Period 2 has an atomic number of 4 and is called Beryllium.

Exam Tip: Determine the atomic number from the period (number of shells) and group (number of valence electrons) to correctly identify the element.

 

Question 28. An element 'A' has atomic number 11, name the period and group number to which it belongs.
Answer: Element 'A' has an atomic number of 11. Its electronic configuration is 2, 8, 1. The period number, based on the number of shells, is 3. The group number, determined by its valence electron, is 1.
In simple words: An element with atomic number 11 is in Period 3, Group 1.

Exam Tip: The number of electron shells indicates the period number, and the number of valence electrons (for main group elements) indicates the group number.

 

Question 29. An element 'P' belongs to group 2 and period = 3, state whether it is a metal or non-metal and nature of its oxides.
Answer: An element 'P' is in Group 2 and Period 3. Elements in Group 2 are metals. Their oxides are basic in nature.
In simple words: An element from Group 2, Period 3 is a metal and forms basic oxides.

Exam Tip: Recall that elements in Group 1 and 2 are always metals, and metals typically form basic oxides.

 

Question 30. The electronic configuration of an atom is 2, 8, 7. Give its atomic number, nature of oxide.
Answer: The electron configuration is 2, 8, 7. Its atomic number is 17. Its oxide is acidic in nature.
In simple words: An atom with electron configuration 2, 8, 7 has an atomic number of 17 and forms an acidic oxide.

Exam Tip: An electron configuration ending in 7 (like 2,8,7 for Chlorine) indicates a halogen, which is a non-metal, and non-metals form acidic oxides.

 

Question 31. An element belongs to group 13 and period 3, name the element and give its valency.
Answer: This element is Aluminium. Its valency is 3.
In simple words: The element in Group 13, Period 3 is Aluminium, and it has a valency of 3.

Exam Tip: For main group elements, the group number often (with some adjustments for groups 13-18) relates directly to the number of valence electrons and thus valency. Period number identifies the element's row.

 

Question 32. What are metalloids? Give 2 examples.
Answer: Elements which exhibit characteristics of both metals and non-metals are known as semi-metals or metalloids. For instance - Boron, Silicon, Germanium, Arsenic.
In simple words: Metalloids are elements that have properties like both metals and non-metals. Boron and Silicon are examples.

Exam Tip: Remember that metalloids typically lie along the staircase line in the periodic table, separating metals from non-metals, and provide at least two common examples.

 

Question 33. An element X belongs to group 17 and element Y belongs to group 1. What type of bond will they form?
Answer:

Both elements will form an ionic bond, resulting in the formula XY.
In simple words: Element X (Group 17) and Element Y (Group 1) will form an ionic bond, creating a compound with the formula XY.

Exam Tip: Remember that elements from Group 1 (alkali metals) readily lose one electron to form a cation, while elements from Group 17 (halogens) readily gain one electron to form an anion, leading to an ionic bond.

 

Question 34. The following elements belong to same period arrange them in order.
Answer:Given atomic radii:

In decreasing order of their atomic radii:The atomic size reduces in a period as you go from left to right.
In simple words: The elements are arranged from largest to smallest atomic radius, showing that atom size decreases when moving right across a period.

Exam Tip: Remember the trend that atomic radius decreases across a period due to increasing nuclear charge pulling the electron shells closer to the nucleus.

 

Question 35. What is the valency of magnesium with atomic number 12 and nitrogen with atomic number 7?
Answer:
Magnesium has atomic number 12. Its electronic configuration is 2, 8, 2. Its valency is 2.
Nitrogen has atomic number 7. Its electronic configuration is 2, 5. Its valency is 3.
In simple words: Magnesium (atomic number 12) has a valency of 2, and Nitrogen (atomic number 7) has a valency of 3.

Exam Tip: Valency is determined by the number of electrons an atom needs to lose or gain to achieve a stable octet (or duet for smaller atoms) in its outermost shell.

 

Question 36. How many shells are present in all the elements that belong to period 3?
Answer: All elements in Period 3 have 3 electron shells in which their electrons are arranged (K, L, M).
In simple words: Every element in Period 3 of the Periodic Table uses three electron shells to hold its electrons.

Exam Tip: The period number in the periodic table directly corresponds to the number of electron shells an element possesses.

 

Question 37. What happens to the electropositive character of elements as we move from left to right of the period in the Periodic Table?
Answer: When you go from left to right across a period, the electropositivity lessens as the likeliness to give up electrons also lessens.
In simple words: Moving from left to right in a period, elements become less electropositive because they are less likely to lose electrons.

Exam Tip: Electropositivity is linked to metallic character; as metallic character decreases across a period, so does the ability to lose electrons easily.

 

Question 38. Fluorine, chlorine, bromine belong to same group. What is common between them?
Answer: All three elements - fluorine, chlorine, and bromine - are in the same group. They share the same number of outer electrons and valency. Because of this, they display similar chemical traits.
In simple words: Fluorine, chlorine, and bromine are in the same group and share similar chemical behaviors because they have the same number of valence electrons.

Exam Tip: Elements in the same group exhibit similar chemical properties because they have the same number of valence electrons, which dictate their reactivity.

 

Question 39. What are halogens? Where are they located in the Periodic Table?
Answer: Halogens are elements possessing 7 valence electrons and a valency of 1. They combine with metals to produce salts. They are found in Group 17 of the Periodic Table.
In simple words: Halogens are elements with 7 outer electrons and a valency of 1, found in Group 17, and they react with metals to form salts.

Exam Tip: Key characteristics of halogens include 7 valence electrons, high reactivity, formation of salts with metals, and their location in Group 17.

 

Question 40. Atomic number of 4 elements is given below which element will belong to the group of Helium.

Answer: Element Y, which has atomic number 36, will be in the same group as Helium. Its electronic configuration is 2, 8, 18, 8, and Helium's is 2. Both are inert gases, meaning their outermost shells are completely filled.
In simple words: Element Y (atomic number 36) belongs to Helium's group because both are inert gases, meaning they don't react much and have full outer electron shells.

Exam Tip: Elements in the same group have similar chemical properties. Noble gases (Group 18) like Helium and Xenon (atomic number 36) are characterized by their complete outermost electron shells and inertness.

Short Answer Type Question

 

Question 1. Why Mendeleev could not assign fixed position to hydrogen in the table?
Answer:
(a) Hydrogen is similar to alkali metals; for example, it joins with halogens, oxygen, and sulfur to create compounds with formulas like those of alkalis.
(b) Like halogens, hydrogen also can be a diatomic molecule and can join with both metals and non-metals, forming covalent compounds.
In simple words: Mendeleev couldn't place hydrogen definitely because it acts like both alkali metals and halogens, showing properties of both.

Exam Tip: Explain hydrogen's dual nature by citing its similarities to both Group 1 (alkali metals) and Group 17 (halogens) to demonstrate the positioning problem.

 

Question 2. Name the group number of the following elements, halogens, alkali metals, inert gases, hydrogen, in the Modern Periodic Table.
Answer:
Halogens - Group No. 17
Alkali metals - Group No. 1
Inert gases - Group No. 18
Hydrogen - Group No. 1
In simple words: Halogens are in Group 17, alkali metals in Group 1, inert gases in Group 18, and hydrogen is also in Group 1.

Exam Tip: Memorize the group numbers for these common element categories as they are fundamental to understanding periodic trends.

 

Question 3. State two characteristics of groups.
Answer: All the elements within a group possess these characteristics:
(a) Every element in a group displays the same number of outer electrons, and consequently exhibits similar characteristics.
(b) When you go from top to bottom in a group, the atomic radius gradually increases, and there is a small, gradual change in properties.
In simple words: Elements in a group have the same number of outer electrons, leading to similar properties. Also, atomic size grows as you move down a group.

Exam Tip: Clearly state that constant valence electrons lead to similar properties, and increasing atomic radius down a group impacts other trends like metallic character.

 

Question 4. What happens to the valency of elements as we move from left to right in a Periodic Table?
Answer: When going from left to right across a Periodic Table, valency initially rises up to 4, and then declines once more.
L-R, Valency in a period \( \rightarrow \) 1 2 3 4 3 2 1 0
In simple words: Moving across a period, valency increases to 4, then decreases to 0.

Exam Tip: Remember the specific pattern of valency across a period: it increases up to 4 (Group 14) and then decreases to 0 (Group 18, noble gases).

 

Question 5. The number of electrons goes on increasing in the outer shell as we move from left to right in a period, why does the atomic size goes on decreasing?
Answer: In a period, all elements possess the same number of electron shells. As we move from left to right across a period, the number of electrons rises, and simultaneously, the number of protons also grows. This means the nuclear attractive force gets stronger, which then draws the outer valence electrons closer to the nucleus. As a result, the atom's size reduces.
In simple words: Atomic size decreases across a period because more protons pull the increasing number of outer electrons closer to the nucleus, even though the number of shells stays the same.

Exam Tip: Explain that despite adding electrons to the same shell, the increased nuclear charge causes a stronger pull, resulting in a smaller atomic radius across a period.

 

Question 6. What happens to the metallic character as we move from top to bottom in a group?
Answer: Metallic character grows when going from top to bottom in a group, as the likeliness to give up electrons rises.
In simple words: Down a group, elements become more metallic because it's easier for them to lose electrons.

Exam Tip: Metallic character is directly related to the ease of losing electrons; down a group, increased atomic size and shielding make electron loss easier, thus increasing metallic character.

 

Question 7. What happens to the non-metallic character as we move from top to bottom in a group?
Answer: The non-metallic character lessens when going from top to bottom in a group, as the likeliness to acquire electrons drops down the group.
In simple words: Moving down a group, elements become less non-metallic because they are less likely to gain electrons.

Exam Tip: Non-metallic character relates to the ease of gaining electrons; down a group, the increasing atomic size and shielding make it harder to attract additional electrons, thus decreasing non-metallic character.

 

Question 8. The atomic number of 'X' is 17. Predict its (a) valency, (b) formula of halide, (c) type of ion formed, (d) reactivity with respect to the other members of same group.
Answer: Element 'X' has atomic number is 17. Its electronic configuration is 2, 8, 7.
(a) Valency = 1
(b) Formula of halide = HX
(c) It forms a negative ion, an anion, denoted as \( X^- \).
(d) It is the most reactive among the elements situated below X in the same group.
In simple words: Element X, with atomic number 17, has an electronic configuration of 2, 8, 7, and its valency is 1. It forms HX halides, creates \( X^- \) anions, and is the most reactive halogen in its group compared to elements below it.

Exam Tip: For halogens (Group 17), reactivity decreases down the group. Therefore, the element with atomic number 17 (Chlorine) is more reactive than elements below it.

 

Question 9. Given below are 4 elements W, X, Y and Z the atomic numbers are 9, 10, 16, 17. Predict the following:
(a) Two elements lying in same group.
(b) Elements in second period.
Answer:

(a) The two elements in the same group are W and Z.
(b) The elements in the second period are W and X.
In simple words: W (Fluorine) and Z (Chlorine) are in the same group (17), while W and X (Neon) are both in the second period.

Exam Tip: Elements in the same group have the same number of valence electrons, and elements in the same period have the same number of electron shells. Use this to identify their positions.

 

Question 10. Write all the elements present in third period of the Periodic Table and give their electronic configuration.
Answer: The elements found in the third period are:
Na, Mg, Al, Si, P, S, Cl, Ar.
Their electronic configurations are:

In simple words: The third period contains Sodium, Magnesium, Aluminium, Silicon, Phosphorus, Sulfur, Chlorine, and Argon, each with its electrons arranged in three shells.

Exam Tip: Remember that all elements in a given period (e.g., Period 3) will have the same number of electron shells, which is equal to the period number itself.

 

Question 11. How does electronic configuration helps us to locate the position of element in the Periodic Table?
Answer: An atom's electronic configuration shows the number of valence electrons and electron shells. The valence electrons aids in finding the group number, while the number of shells in an atom indicates its period.
In simple words: Electronic configuration helps us find an element's spot because valence electrons tell us the group, and electron shells tell us the period.

Exam Tip: Clearly state that valence electrons determine the group number (for main group elements) and the number of shells determines the period number.

 

Question 12. What are the merits of Mendeleev's Periodic Table?
Answer: The advantages of Mendeleev's Periodic Table include:
1. Mendeleev included empty spaces in his table. He foretold the chemical traits of three elements, which were later found and possessed the characteristics he had anticipated, namely gallium, germanium, and scandium.
2. He organized the elements highly systematically into periods and groups.
In simple words: Mendeleev's table was good because he predicted new elements and arranged them in a very organized way.

Exam Tip: Highlight Mendeleev's successful prediction of new elements and their properties, as well as his systematic arrangement of known elements, as key strengths.

 

Question 13. Why does the reactivity of metals increases and that of non-metals decreases as we move down the group?
Answer: The reactivity of metals is linked to how easily they give up electrons. As an atom's size grows down a group, the outer electrons are readily detached, thus forming a positive ion, which increases metallic reactivity.
Non-metals' reactivity relies on their ability to acquire electrons. As we move down the group, the likeliness to gain electrons reduces because the atomic size expands, and the nuclear pull lessens.
In simple words: Metal reactivity increases down a group because larger atoms lose electrons more easily. Non-metal reactivity decreases down a group because larger atoms attract new electrons less strongly.

Exam Tip: For metals, greater atomic size down a group means lower ionization energy and higher reactivity. For non-metals, greater atomic size means lower electron affinity and lower reactivity.

 

Question 14. List the elements present in 2nd period. Write their atomic number and electronic configuration.
Answer: The elements present in Period 2, along with their atomic numbers and electronic configurations, are:

In simple words: The elements in the second period are Lithium to Neon, each having two electron shells, with their atomic numbers increasing from 3 to 10.

Exam Tip: For any period, list elements in increasing order of atomic number, and ensure their electronic configuration correctly reflects the period number (number of shells) and group number (valence electrons).

 

Question 15. For the following given elements predict the;
(a) Valency
(b) Period number
(c) Group number Na (11), Al(13), Cl(17), K(19)
Answer:

In simple words: The table shows the atomic number, electron arrangement, valency, period, and group for Sodium, Aluminium, Chlorine, and Potassium.

Exam Tip: Remember that valency is determined by the number of valence electrons, the period number by the number of shells, and the group number by the valence electrons (or valence electrons + 10 for groups 13-18).

 

Question 16. Elements of group 1 are given below with their atomic number: Li (3) Na (11) K (19)
Give the order of their (a) atomic size. (b) reactivity.
Answer:

(a) The atomic size increases as a new shell is included in each element as we go from top to bottom. So, the order is Li < Na < K.
(b) Reactivity increases because it is easier to lose electrons as the atom's size grows, due to a reduction in nuclear pull. So, the order is K > Na > Li.
In simple words: Down Group 1, atomic size increases (Li < Na < K) and reactivity increases (K > Na > Li) because it's easier for larger atoms to lose electrons.

Exam Tip: Remember that down a group, atomic size increases, making it easier for metals to lose electrons and thus increasing their reactivity. For non-metals, reactivity generally decreases down a group.

 

Question 17. Lithium, sodium and potassium belong to same group called alkali metals. Why?
Answer: Lithium, sodium, and potassium all have one valence electron, so they are in the same group. This group is named the alkali metals group as all these elements create oxides that dissolve in water to make alkalis.
In simple words: Lithium, sodium, and potassium are called alkali metals and are in the same group because they all have one valence electron and form basic oxides.

Exam Tip: Define alkali metals by their valence electron count (1) and their property of forming strong bases when their oxides react with water.

 

Question 18. Carbon with atomic number 6 and silicon with atomic number 14 belong to same group although carbon is non-metal and silicon is semi-metal. Comment.
Answer: Carbon, having atomic number 6, and silicon, having atomic number 14, are in the same group. Silicon possesses three electron shells with an electronic configuration of 2, 8, 4, while carbon possesses two electron shells with an electronic configuration of 2, 4. Thus, silicon is larger in size and has a weaker nuclear attraction compared to carbon, which results in electrons being easier to remove, giving silicon more metallic character than carbon.
In simple words: Carbon and silicon are in the same group but have different natures because silicon is larger and has more electron shells, affecting its metallic character, making it more metallic than carbon.

Exam Tip: When comparing elements in the same group, remember that increasing atomic size (down the group) generally leads to a decrease in nuclear attraction on outer electrons, making it easier to lose them and increasing metallic character.

 

Question 19. The position of 3 elements A, B and C in the Periodic Table is shown below:

Giving reasons, explain the following:
(a) Element A is a non-metal.
(b) Atom of element C has larger size than A.
(c) Element B has a valency of 1.
Answer:
(a) Element A is in Group VII and has 7 valence electrons. It can acquire one electron to become stable. So, it is a non-metal as it forms a negative ion.
(b) Element C has three electron shells, while element A has two electron shells. Therefore, C is larger than A.
(c) Element B has one valence electron. It can give up this electron to become stable, so its valency is 1.
In simple words: A is a non-metal (Group VII, gains electrons). C is larger than A because it has more electron shells (Period 3 vs Period 2). B has a valency of 1 because it's in Group I and loses one electron.

Exam Tip: Use the group number to infer valence electrons and chemical nature, and the period number to compare atomic size. Valency is determined by the number of electrons lost or gained.

 

Question 20. The elements of the third period of the Periodic Table are given below:

(a) Which atom is bigger – Na or Mg? Why?
(b) Identify the most – (i) Metallic and (ii) Non-metallic element in period 3.
Answer:
(a) The Sodium atom is larger than Magnesium. This happens because as we progress from Na to Cl across the period, the atomic number rises, and the nuclear charge also grows. This pulls the outer electrons closer to the nucleus, hence the atomic size shrinks.
(b) (i) The most metallic element is Na.
(ii) The most non-metallic element is Cl.
In simple words: Sodium is bigger than Magnesium because atomic size decreases across a period due to stronger nuclear pull. Sodium is the most metallic, and Chlorine is the most non-metallic element in Period 3.

Exam Tip: Remember the general trends: atomic size decreases from left to right across a period, and metallic character decreases while non-metallic character increases across a period.

Long Answer Type Questions

 

Question 1. The atomic number of element X is 17 predict its
(a) Physical state.
(b) Name of element.
(c) Formula of its compound with hydrogen
(d) Metal or Non-metal.
(e) Formula of its molecule.
Answer: Element X has an atomic number of 17. Its electronic configuration is 2, 8, 7.
(a) Its physical state is Gas.
(b) The name of the element is Chlorine.
(c) Its compound with hydrogen is HCl.
(d) It is a Non-metal.
(e) Its molecular formula is Cl\( _2 \).
In simple words: Element X, with atomic number 17 and configuration 2, 8, 7, is Chlorine, a non-metal gas. It forms HCl with hydrogen and exists as Cl\( _2 \).

Exam Tip: Recognize that an atomic number of 17 corresponds to Chlorine, a Group 17 element (halogen). From this, deduce its non-metallic nature, gaseous state, common compounds, and diatomic molecular form.

 

Question 2. Two elements A and B belong to group 1 and 2 respectively in the same period. Compare them with respect to:
(a) Valency
(b) Size of atom
(c) Formula of oxide
(d) Nature of oxide
(e) Metallic character
Answer:
(a) Valency of A is 1, and Valency of B is 2.
(b) Atom A is larger in size compared to Atom B.
(c) The formula of A's oxide is \( A_2O \), and B's oxide is BO.
(d) Both oxides are Basic in nature.
(e) A has a more metallic character than B.
In simple words: In the same period, Group 1 element A has a valency of 1, is larger, forms \( A_2O \) oxide, and is more metallic than Group 2 element B, which has a valency of 2 and forms BO oxide. Both oxides are basic.

Exam Tip: Remember trends across a period: atomic size decreases from left to right, metallic character decreases, and valency generally increases then decreases. Group 1 elements are more metallic and have a lower valency than Group 2 elements in the same period.

 

Question 3. Give all characteristics of group.
Answer: Characteristics of a group:
(a) Valence electrons: All elements in a group exhibit the same number of valence electrons.
(b) Valency: The valency of all elements in a group stays constant.
(c) Atomic size: The atomic size increases as you move down the group.
(d) Metallic character: For metals, the metallic character grows as you descend the group.
(e) Non-metallic character: For non-metals, the non-metallic character lessens as you descend the group.
In simple words: Elements in a group share similar valence electrons and valency. Going down a group, atomic size and metallic character increase, while non-metallic character decreases.

Exam Tip: For groups, remember the key trends: constant valence electrons/valency, increasing atomic size, increasing metallic character (for metals), and decreasing non-metallic character (for non-metals) as you move down.

 

Question 4. Give the characteristics of a period.
Answer: In a period, as we go from left to right:
(a) Valence electrons: Increases (1, 2, 3, 4, 5, 6, 7, 8.)
(b) Valency: Valency initially increases to 4 and then reduces to zero (1, 2, 3, 4, 3, 2, 1, 0).
(c) Size of atom: The size of the atom gets smaller.
(d) Metallic character: Decreases.
(e) Non-metallic character: Increases.
In simple words: Across a period, valence electrons increase, valency goes up then down, atomic size shrinks, metallic character decreases, and non-metallic character increases.

Exam Tip: For periods, focus on how the number of valence electrons changes, which directly impacts valency, atomic size, and metallic/non-metallic properties across the row.

 

Question 5. You are given some descriptions about five different elements. Identify these elements to place them in the Modern Periodic Table.
(a) Essential for breathing and burning.
(b) Inactive, two electrons in the outermost shell.
(c) Atom has same number of protons, electrons and neutrons, used in fertilizer industry.
(d) Number of neutrons, protons are same used in building our bones.
(e) This element form the hardest naturally occurring substance as allotrope.
Answer:

In simple words: The elements described are Oxygen (for breathing/burning), Helium (inactive, 2 outer electrons), Nitrogen (used in fertilizer), Calcium (for bones), and Carbon (forms hardest naturally occurring substance).

Exam Tip: Connect the given characteristic properties and uses directly to specific elements and their positions in the periodic table.

 

Question 6. Name the following elements
(a) Two shells, both of which are completely filled.
(b) Three shells with 2 valence electrons.
(c) Group 1, two shells.
(d) Group 17, period 3.
(e) Metal, with valency 3 group number 13 period 3.
Answer:
(a) Having two shells (K and L), with full electron shells (2 and 8). This gives an atomic number of 10, so the element is Neon.
(b) Having three shells (K, L, M) and 2 valence electrons (2, 8, 2). This results in an atomic number of 12, identifying the element as Magnesium.
(c) Being in Group 1 and having two shells (K and L), its electronic configuration is 2, 1. This makes its atomic number 3, so the element is Lithium.
(d) Being in Group 17 and Period 3 (K, L, M shells), its electronic configuration is 2, 8, 7. This yields an atomic number of 17, so the element is Chlorine.
(e) Being in Group 13 with a valency of 3, and in Period 3 (K, L, M shells), its electronic configuration is 2, 8, 3. This indicates an atomic number of 13, meaning the element is Aluminium.
In simple words: (a) Neon (10), (b) Magnesium (12), (c) Lithium (3), (d) Chlorine (17), and (e) Aluminium (13) are the elements based on the given descriptions.

Exam Tip: Use the number of shells to determine the period and the number of valence electrons (or (Group Number - 10) for groups 13-18) to determine the group, then identify the element.

Practical Based Questions Solved

 

Question 1.

(a) Which atom is bigger, Na or Mg? Why?
(b) Identify the most metallic and most non-metallic element in period 3.
Answer:
(a) Sodium is larger in size than magnesium because the effective nuclear charge on sodium is weaker than on magnesium, and the outer shell electrons are drawn more strongly to the center of the magnesium atom.
(b) The most metallic element is Na, as it can easily lose electrons to create a positive ion. Chlorine is the most non-metallic element as it can easily gain electrons to create a negative ion.
In simple words: Sodium is larger than Magnesium because its nuclear pull is weaker. Sodium is the most metallic element since it easily loses electrons, while Chlorine is the most non-metallic as it readily gains electrons.

Exam Tip: Remember that across a period, atomic size decreases (due to increased effective nuclear charge) and metallic character decreases, while non-metallic character increases.

 

Question 2. For each of the following, give the name of an element from Period 3 (sodium to argon), which matches the description.
(a) an element which is gaseous at room temperature and pressure
(b) an element that is added to water to kill bacteria
(c) an element that forms a basic oxide of the type XO
(d) an element that forms an amphoteric oxide
(e) an element that reacts vigorously with cold water to produce hydrogen.
Answer:
(a) Chlorine
(b) Chlorine
(c) Magnesium
(d) Aluminium
(e) Sodium
In simple words: We identify Period 3 elements based on their properties: Chlorine is a gas and kills bacteria, Magnesium forms basic oxides, Aluminium forms amphoteric oxides, and Sodium reacts strongly with cold water.

Exam Tip: Recall the specific properties of Period 3 elements: Chlorine (gas, disinfectant), Magnesium (Group 2 metal, basic oxide), Aluminium (amphoteric oxide), and Sodium (Group 1 metal, vigorous reaction with water).

 

Question 3. The table given below shows four elements, named P, Q, R and S. Complete the table:
Answer:

Nucleon number is the total number of protons and neutrons in an atom's nucleus.
In simple words: The table is completed by using the relationships between proton number (atomic number), electron number (for ions), neutron number, and nucleon number (mass number) to identify the elements and their symbols.

Exam Tip: Remember that proton number equals atomic number, electron number differs from proton number in ions, nucleon number is protons + neutrons, and the symbol represents the element and its charge.

 

Question 4. Rubidium, Rb, is a Group I element. It has similar physical and chemical properties to the other elements in Group I
(a) When rubidium is added to cold water a reaction occurs. Suggest two observations that would be made when rubidium is added to cold water.
(b) Write a chemical equation for the reaction between rubidium and water.
(c) Suggest one safety measure that should be used when rubidium is added to cold water.
Answer:
(a) The reaction is extremely energetic, fires might appear, and a fizzing noise can be heard.
(b) \( 2Rb + 2H_2O \rightarrow 2RbOH + H_2 \)
(c) The reaction produces much heat, and hot water or rubidium might splash onto the body, so keep far away, use safety eyewear, a laboratory coat, and hand protection.
In simple words: (a) Rubidium reacts strongly with cold water, causing flames and a sizzling sound. (b) Rubidium and water create rubidium hydroxide and hydrogen gas. (c) Because the reaction is hot and splashy, stand far back and wear safety gear like glasses and gloves.

 

Question 5. Aluminium metal is widely used in the making of cooking utensils like pans, cookers etc. Suggest the property of aluminium that makes it suitable for this use.
Answer: Aluminium is a metal with excellent heat conduction, and it resists rust because it creates a defensive coating of aluminium oxide on its surface, which then does not interact with any other element or compound.
In simple words: Aluminium conducts heat well and does not rust easily. It forms a protective oxide layer that stops it from reacting, making it good for cooking pots.

 

Question 6. Zinc metal is widely used in the construction of dry cells and in making of alloys. Name the alloy of zinc metal and state the other metal used in this alloy.
Answer: Zinc is employed in creating brass. The additional metal used in this alloy is copper.
In simple words: Zinc is used in brass. The other metal in brass is copper.