RBSE Solutions Class 10 Science Chapter 7 Atomic Theory, Periodic Classification, and

I. Multiple Choice Questions

Question 1. Which type of radiations were used in Rutherford's experiment?
(a) a
(b) β
(c) γ
(d) λ
Answer: (a) a
In simple words: Rutherford used alpha (a) particles in his famous gold foil experiment. These particles helped him understand the structure of an atom.

🎯 Exam Tip: Remember Rutherford's experiment proved the existence of a small, dense, positively charged nucleus within an atom.

Question 2. Which is the smallest particle of matter?
(a) Molecule
(b) Atom
(c) Element
(d) Compound
Answer: (b) Atom
In simple words: An atom is considered the basic building block of all matter. It is the smallest unit that retains the chemical identity of an element.

🎯 Exam Tip: While sub-atomic particles exist, an atom is generally considered the smallest particle of matter that can exist independently and participate in chemical reactions.

Question 3. Who was the first to carry out periodic classification of elements?
(a) Dobereiner
(b)
(c)
(d)

🎯 Exam Tip: Dobereiner proposed the law of triads, grouping elements with similar properties into sets of three. However, his classification was limited to only a few elements.

Question 4. The modern periodic table is based on which property of matter?
(a) Atomic structure
(b) Atomic weight
(c) Atomic number
(d) Valency
Answer: (c) Atomic number
In simple words: The modern periodic table arranges elements based on their atomic number. This is different from earlier tables that used atomic weight.

🎯 Exam Tip: Atomic number (number of protons) is a more fundamental property for arranging elements as it determines their chemical behavior.

Question 5. What is the number of periods and groups in the modern periodic table?
(a) 7 and 18
(b) 9 and 18
(c) 7 and 20
(d) 9 and 20
Answer: (c) 7 and 18
In simple words: The modern periodic table has 7 horizontal rows called periods and 18 vertical columns called groups. Each period and group shows specific trends in properties.

🎯 Exam Tip: Periods represent the number of electron shells, while groups relate to the number of valence electrons and similar chemical properties.

Question 6. Which trend is shown by the size of atom when we come down in a group?
(a) Reduces
(b) No change
(c) Is irregular
(d) Increases
Answer: (d) Increases
In simple words: As you go down a group in the periodic table, the size of atoms gets bigger. This is because new electron shells are added with each new period.

🎯 Exam Tip: The increasing atomic size down a group is due to the addition of new principal energy shells, which increases the distance of valence electrons from the nucleus.

Question 7. van der Waal's radius is - than covalent radius?
(a) Smaller than
(b) Larger than
(c)
(d)

🎯 Exam Tip: Van der Waal's radius applies to atoms that are not bonded, representing half the distance between non-bonded atoms in a solid. Covalent radius is for atoms bonded together.

Question 8. What is the number of elements in a small period?
(a) 2
(b) 8
(c) 18
(d) 32
Answer: (b) 8
In simple words: A "small period" in the periodic table, like the second and third periods, contains 8 elements. These periods help show how properties change gradually.

🎯 Exam Tip: The first period has 2 elements, the second and third have 8, the fourth and fifth have 18, and the sixth and seventh have 32 elements.

Question 9. Which energy is given to separate electrons from a neutral atom?
(a) Electron gain enthalpy
(b) Electronegativity
(c) Ionisation enthalpy
(d) Activation energy
Answer: (c) Ionisation enthalpy
In simple words: The energy needed to remove an electron from a neutral atom is called ionization enthalpy. It's like the "cost" to pull an electron away.

🎯 Exam Tip: Ionisation enthalpy is an important property that indicates an atom's metallic character (lower value means more metallic) and its tendency to form positive ions.

Question 10. Which of the following has the maximum electronegativity?
(a) H
(b) Na
(c) Ca
(d) F
Answer: (d) F
In simple words: Fluorine (F) has the strongest pull for electrons in a chemical bond. This makes it the most electronegative element among the options.

🎯 Exam Tip: Electronegativity generally increases across a period and decreases down a group in the periodic table, making fluorine the most electronegative element overall.

Question 11. Elements of which group show the highest metallic properties?
(a) 1
(b) 2
(c) 5
(d) 6
Answer: (a) 1
In simple words: Group 1 elements, also known as alkali metals, are the most metallic. They easily lose their single outermost electron.

🎯 Exam Tip: Metallic character is defined by the ease with which an element can lose electrons. Elements on the far left and bottom of the periodic table are most metallic.

Question 12. What is the name of Thomson's model?
Answer: The name of Thomson's model is the Plum pudding model. This model proposed that atoms are made up of a positively charged sphere with negatively charged electrons embedded within it, like plums in a pudding.
In simple words: Thomson's model of the atom was called the Plum pudding model. He thought atoms were like a ball of positive charge with tiny negative bits inside.

🎯 Exam Tip: Remember Thomson's model was one of the early atomic models that recognized the existence of electrons, but it was later disproved by Rutherford's experiment.

Question 13. What is another name of Bohr's orbit?
Answer: Another name for Bohr's orbit is energy level. Bohr's model introduced the idea that electrons move in specific, fixed paths or shells around the nucleus, each with a distinct amount of energy.
In simple words: Bohr's orbits are also called energy levels. They are specific paths where electrons move around the atom's center.

🎯 Exam Tip: Bohr's model helped explain the stability of atoms and their emission spectra by quantizing electron energies.

Question 14. Write the modern periodic law.
Answer: The modern periodic law states that the chemical properties of elements are a periodic function of their atomic number. This means that when elements are arranged by increasing atomic number, their properties repeat in a regular pattern.
In simple words: The modern periodic law says that element properties depend on their atomic number and repeat in a pattern.

🎯 Exam Tip: This law is key to understanding the arrangement of the modern periodic table, which corrected the limitations of Mendeleev's periodic law based on atomic mass.

Question 15. Write the periodic law of Mendeleev.
Answer: The periodic law of Mendeleev states that the chemical properties of elements are a periodic function of their atomic mass. He arranged elements based on this principle, noticing repeating patterns in their properties.
In simple words: Mendeleev's periodic law said that element properties repeat based on their atomic mass.

🎯 Exam Tip: Mendeleev's table was a groundbreaking achievement, but it had limitations because atomic mass is not the most fundamental property.

Question 16. Mendeleev arranged the elements on the basis of which property?
Answer: Mendeleev arranged the elements primarily on the basis of atomic mass. He observed that when elements were arranged in increasing order of their atomic masses, elements with similar properties appeared at regular intervals.
In simple words: Mendeleev put elements in order using their atomic mass.

🎯 Exam Tip: While atomic mass was his primary criterion, Mendeleev also considered chemical properties, especially those of oxides and hydrides, to group elements.

Question 17. Which name is given to elements of 18th Group?
Answer: Elements of the 18th group are known as Noble elements (or noble gases). These elements have a stable electronic configuration and are chemically unreactive under normal conditions.
In simple words: The elements in Group 18 are called Noble elements or noble gases.

🎯 Exam Tip: Noble gases like Helium, Neon, and Argon are famous for their inertness due to having a complete outer electron shell, which makes them very stable.

Question 18. What are the other names for elements of d-block and f-block?
Answer: Elements of the d-block are called transition elements. These metals are known for their variable valencies and ability to form colored compounds. Elements of the f-block are called inner transition elements, comprising the lanthanides and actinides.
In simple words: d-block elements are called transition elements, and f-block elements are called inner transition elements.

🎯 Exam Tip: Transition elements (d-block) play crucial roles in catalysis and biochemistry, while inner transition elements (f-block) are often radioactive and important in nuclear technology.

Question 19. Explain the "position of metals, non-metals and metalloids in Modern Periodic Table.
Answer: In the modern periodic table, elements are broadly categorized into metals, non-metals, and metalloids. Metals are generally located on the left side of the periodic table. As we move from left to right across a period, the metallic properties of elements gradually decrease. Non-metals are found on the right side of the periodic table. Between the metals and non-metals, there is a diagonal line of elements called metalloids. These elements exhibit properties that are intermediate between those of metals and non-metals, such as silicon and germanium.
In simple words: Metals are on the left side of the periodic table, non-metals are on the right, and metalloids are found in between. Metallic nature lessens as you go from left to right.

🎯 Exam Tip: The zigzag line starting from Boron (B) helps to separate metals from non-metals, with elements along this line often being metalloids.

Question 20. Explain the periodicity of electron gain enthalpy in a group.
Answer: Electron gain enthalpy refers to the energy change when an electron is added to a neutral gaseous atom. When moving from left to right across a period, atomic size generally decreases, and the effective nuclear charge increases. This means electrons are held more tightly, so electron gain enthalpy becomes more negative (more energy is released). However, when moving down a group, atomic size increases, and the added electron is further from the nucleus, leading to a decrease in the effective nuclear attraction. Consequently, electron gain enthalpy becomes less negative, meaning less energy is released upon gaining an electron.
In simple words: As you go down a group, atoms get bigger, so it's harder for them to attract new electrons. This makes their electron gain enthalpy less negative.

🎯 Exam Tip: Remember that negative electron gain enthalpy means energy is released (exothermic), indicating a strong attraction for the incoming electron.

Question 21. What do you understand by van der Waal radius and covalent radius?
Answer: The van der Waal's radius is defined as half the distance between the nuclei of two non-bonded atoms of the same element that are closest to each other in a solid state. This radius represents the overall size of the atom in a non-bonded state. In contrast, the covalent radius is half the distance between the nuclei of two atoms of the same element when they are joined together by a single covalent bond. It indicates the size of an atom when it is chemically bonded to another.
In simple words: Covalent radius is half the distance between two atoms that are chemically bonded. Van der Waal's radius is half the distance between two non-bonded atoms that are just touching.

🎯 Exam Tip: Van der Waal's radius is usually larger than covalent radius for the same element because it represents the non-bonded distance between atoms.

Question 22. A cation is smaller than neutral atom and an anion is larger than a neutral atom. Why?
Answer: A cation is formed when a neutral atom loses one or more electrons. The removal of electrons from the outermost shell leads to a decrease in the effective nuclear charge per remaining electron, pulling the remaining electrons closer to the nucleus. This results in a smaller atomic size for the cation compared to its neutral atom. Conversely, an anion is formed when a neutral atom gains one or more electrons. The addition of extra electrons increases electron-electron repulsion, causing the electron cloud to expand. This expansion leads to a larger atomic size for the anion compared to its neutral atom, as the effective nuclear charge per electron decreases.
In simple words: Cations are smaller because they lose electrons, making the nucleus pull the remaining electrons closer. Anions are larger because they gain electrons, causing more repulsion and expanding the electron cloud.

🎯 Exam Tip: Remember that the number of protons (nuclear charge) remains constant when an atom forms an ion; only the number of electrons changes, affecting the electron cloud's size.

Question 23. What is nuclear charge? How does it behave in a period?

🎯 Exam Tip: Understanding nuclear charge helps explain trends like atomic radius, ionization energy, and electronegativity across the periodic table.

Question 24. How does valency vary when we move from left to right through a period?
Answer: As we move from left to right across a period in the periodic table, the valency of elements generally follows a specific trend. It first increases from 1 to 4, and then it decreases from 4 back to zero. This is because the number of valence electrons increases from 1 to 8 across a period, influencing how many bonds an atom can form.
In simple words: Valency first goes up from 1 to 4 when you move right across a period, and then it goes down to zero.

🎯 Exam Tip: Valency is determined by the number of electrons an atom needs to gain, lose, or share to achieve a stable outer electron shell, typically 8 electrons (octet rule).

Question 25. Write Dalton's atomic theory.
Answer: Dalton's Atomic Theory provided the foundational understanding of atoms. Its main points are:

• Each substance is made up of small particles which are called atoms.
• Atoms are indivisible and cannot be created or destroyed during chemical reactions.
• All atoms of an element have the same properties, including mass, size, and chemical behavior.
• Atoms of different elements have different properties, such as different masses, sizes, and chemical behaviors.
• Atoms of different elements combine in simple whole number ratios to form molecules or compounds.
• Chemical reactions involve the rearrangement of atoms. An atom can neither be created nor destroyed in a chemical reaction.

In simple words: Dalton's theory states that all matter is made of tiny, indivisible atoms. Atoms of the same element are identical, and atoms of different elements are different. Atoms combine in simple ways to make new substances.

🎯 Exam Tip: While some aspects of Dalton's theory have been modified (e.g., atoms are divisible into subatomic particles, isotopes exist), it remains a cornerstone of modern chemistry.

Atomic Theory, Periodic Classification, and Properties of Elements Long Answer Type Questions

Question 26. Enlist merits and demerits of Mendeleev's Periodic Table.
Answer: Merits of Mendeleev's Periodic Table:

• Mendeleev was the first to properly arrange and classify all known elements, explaining the periodicity in their chemical properties.
• He left some gaps for undiscovered elements, accurately predicting their existence and properties. For example, he predicted Eka-aluminium (Gallium) and Eka-silicon (Germanium), which were later discovered.

Demerits of Mendeleev's Periodic Table:

• The position of hydrogen was not fixed, as it showed properties similar to both alkali metals and halogens.
• Isotopes, which have different atomic masses but the same chemical properties, could not be given separate places in the table.
• The increasing order of atomic mass was not always maintained; some elements with higher atomic mass were placed before those with lower atomic mass to maintain chemical similarities.

In simple words: Mendeleev's table was good because he grouped elements well and even predicted new ones. But it had problems like not finding a clear spot for hydrogen or isotopes.

🎯 Exam Tip: Focus on the predictive power as the biggest merit and the lack of explanation for isotopes and hydrogen's position as key demerits for Mendeleev's table.

Question 27. What trend is shown by following properties in the periodic table?
(a) Atomic radius
(b) Enthalpy of ionisation
(c) Electronegativity
Answer:
(a) Atomic radii decrease when we move from left to right across a period. Atomic radii increase when we move down a group.
(b) Enthalpy of ionisation increases across a period and decreases when we move down a group.
(c) Electronegativity increases when we move across a period. It decreases when we move down a group. Elements on the left side of the periodic table have lower electronegativity.
In simple words: Atomic size gets smaller across a row and bigger down a column. Ionisation energy gets higher across a row and lower down a column. Electronegativity gets higher across a row and lower down a column.

🎯 Exam Tip: Understanding these periodic trends (atomic size, ionization enthalpy, and electronegativity) is fundamental to predicting an element's chemical behavior.

Question 28. Explain the classification of elements as per the Modern Periodic Table.
Answer: The Modern Periodic Table arranges elements based on their atomic number, leading to a clear classification into periods and groups. It features 7 horizontal rows, known as periods, and 18 vertical columns, known as groups. Each period is associated with a specific energy level or electron shell number. For example, the first period is the smallest, containing only two elements. The second and third periods each contain 8 elements, known as small periods. The fourth and fifth periods each have 18 elements, referred to as long periods, and the sixth and seventh periods contain 32 elements, called very long periods. The group number of an element typically corresponds to the number of valence electrons it possesses, which dictates its chemical properties. Metals are generally located on the left side of the table, non-metals on the right, and noble elements (or noble gases) are found in the last group, Group 18.
In simple words: The Modern Periodic Table organizes elements by atomic number into 7 periods (rows) and 18 groups (columns). Periods show energy levels, and groups show similar properties. Metals are on the left, non-metals on the right, and noble gases are in the last group.

🎯 Exam Tip: Recall that the modern periodic table successfully addressed the limitations of earlier classifications, particularly regarding the placement of isotopes and anomalous pairs, by using atomic number as the fundamental classifying property.

Question 29. Explain the gold leaf experiment by Rutherford. What were the observations and conclusions made by Rutherford?
Answer: Rutherford's Gold Foil Experiment:
Rutherford conducted an experiment where he took a very thin gold foil, about 100 nm thick, and bombarded it with high-speed alpha (α) rays. Alpha particles are positively charged helium nuclei. The setup included a fluorescent screen around the gold foil to detect the alpha particles after they passed through.
Observations:

• Most of the fast-moving alpha particles passed straight through the gold foil without any deflection.
• Some alpha particles were deflected by small angles.
• A very small fraction (about 1 in 12,000) of the alpha particles bounced back almost completely.

Conclusions:

• Since most alpha particles passed straight through, a major portion of the atom must be empty or hollow and bear no charge.
• The slight deflection of some alpha particles suggested that there is a small, positively charged region in the atom. This positive charge must be dense enough to repel the positively charged alpha particles.
• The very few alpha particles that bounced back indicated that the positively charged portion of the atom, which Rutherford named the nucleus, is extremely small and very dense compared to the overall size of the atom.

In simple words: Rutherford shot tiny positive particles at a thin gold sheet. Most went straight through, some bent a little, and a few bounced back. He concluded that atoms are mostly empty space with a tiny, heavy, positive center called the nucleus.

🎯 Exam Tip: Rutherford's experiment revolutionized the atomic model, replacing Thomson's "plum pudding" model with the nuclear model, which described a dense, positively charged nucleus surrounded by electrons.

Atomic Theory, Periodic Classification, and Properties of Elements Additional Questions Solved

I. Multiple Choice Questions:

Question 1. 14 elements after actinium are called
(a) Lanthanides
(b) Actinides
(c) d-block elements
(d) p-blockelements
Answer: (b) Actinides
In simple words: The 14 elements that come after actinium in the periodic table are known as Actinides. They are part of the f-block elements.

🎯 Exam Tip: Remember that both Lanthanides and Actinides are inner transition elements, typically placed in a separate block below the main periodic table.

Question 2. An element has an atomic number of 15 with which of the following elements will it show similar chemical properties?
(a) Be(4)
(b) Ne(10)
(c) N(7)
(d) O(8)
Answer: (c) N(7)
In simple words: An element with atomic number 15 is Phosphorus (P). It is in Group 15, just like Nitrogen (N) which has atomic number 7. Elements in the same group have similar chemical properties.

🎯 Exam Tip: Elements in the same group have the same number of valence electrons, which largely determines their chemical reactivity and properties.

Question 3. 8 is
(a) 6, 2
(b) 16, 2
(c) 6, 8
(d) 16,4

🎯 Exam Tip: Always double-check that the question is fully stated and makes sense before attempting to answer, as incomplete questions can lead to confusion.

Question 4. An element belongs to period 2 and group 2 the number of valence electrons in the atoms of this element is
(a) 2
(b) 4
(c) 3
(d) 1
Answer: (a) 2
In simple words: For an element in Group 2, it will always have 2 valence electrons. The period number tells you which shell these electrons are in, but the group number directly tells you the number of valence electrons.

🎯 Exam Tip: The group number (for main group elements, excluding transition metals) directly indicates the number of valence electrons in an atom, which determines its chemical bonding behavior.

Question 5. In the third period of the Periodic Table the element having smallest size is
(a) Na
(b) Ar
(c) Cl
(d) Si
Answer: (b) Ar
In simple words: As you move across a period from left to right, the atomic size gets smaller. Argon (Ar) is the last element in the third period (a noble gas), so it has the smallest size due to increased nuclear charge pulling electrons inward.

🎯 Exam Tip: Remember that noble gases generally have the smallest atomic radii within their respective periods due to high effective nuclear charge and a complete outer shell.

Question 6. Electronic configuration of \( \text{Al}^{+3} \) is
(a) 2,8, 3
(b) 2, 8, 8
(c) 2, 8
(d) 2, 8, 8, 3
Answer: (c) 2, 8
In simple words: Aluminium (Al) has 13 electrons (2,8,3). When it loses 3 electrons to become \( \text{Al}^{+3} \), it now has 10 electrons, which means its electron configuration is 2, 8. This is the same as Neon, making it very stable.

🎯 Exam Tip: When forming positive ions (cations), metals lose electrons from their outermost shell to achieve a stable noble gas configuration, reducing their atomic size.

Question 8. Which of the following sets does not belong to a group?
(a) Li, Na, K
(b) B, C, N
(c) B, Al, Ga
(d) O, S, Se
Answer: (b) B, C, N
In simple words: Li, Na, K are all in Group 1. B, Al, Ga are all in Group 13. O, S, Se are all in Group 16. But B, C, N are in different groups (13, 14, 15), so they don't form a group together.

🎯 Exam Tip: Elements in the same group share similar chemical properties because they have the same number of valence electrons.

Question 9. An element with atomic number 11 will form a basic oxide.
(a) 7
(b) 17
(c) 14
(d) 11
Answer: (d) 11
In simple words: An element with atomic number 11 is Sodium (Na), which is an alkali metal. Metals typically form basic oxides. Non-metals (like atomic number 7 for Nitrogen, 17 for Chlorine) form acidic oxides, and metalloids (like atomic number 14 for Silicon) form amphoteric oxides.

🎯 Exam Tip: Generally, metallic character increases down a group and decreases across a period, meaning elements on the left side of the periodic table tend to form basic oxides.

Question 10. Which of the following elements with atomic numbers 11,19,14,18,23 belong to the same period?
(a) 11,14, 23
(b) 11,18, 23
(c) 11,14,18
(d) 14,19, 23
Answer: (c) 11,14,18
In simple words: Elements in the same period have the same number of electron shells. Atomic number 11 (Na), 14 (Si), and 18 (Ar) all belong to the third period, meaning they have 3 electron shells. Atomic number 19 (K) is in Period 4, and 23 (V) is also in Period 4.

🎯 Exam Tip: To find the period an element belongs to, write its electronic configuration. The highest principal quantum number (shell number) corresponds to its period.

Question 12. Identify the group which is not a Dobereiner triad
(a) Li, Na, K
(b) Be, Mg, Cr
(c) Ca, Sr, Ba
(d) Cl, Br, I
Answer: (b) Be, Mg, Cr
In simple words: Dobereiner's triads grouped three elements where the middle one had properties that were an average of the other two. Beryllium, Magnesium, and Chromium do not form such a triad. Li, Na, K; Ca, Sr, Ba; and Cl, Br, I are all known Dobereiner triads.

🎯 Exam Tip: Dobereiner's triads were an early attempt at classifying elements, where the atomic mass of the middle element was approximately the average of the other two.

Question 13. Which is not true about noble gases?
(a) They are non-metallic in nature
(b) They exist in atomic forth
(c) They are radioactive in nature
(d) Xenon is the most reactive among these
Answer: (c) They are radioactive in nature
In simple words: It is not true that all noble gases are radioactive. While some noble gases like Radon are radioactive, it's not a general characteristic of the entire group. They are non-metals, exist as single atoms, and Xenon is the most reactive among them.

🎯 Exam Tip: Noble gases are characterized by their extreme stability and low reactivity, primarily due to their full valence electron shells. They generally exist as monatomic gases.

Question 14. Which of the given elements A, B, C, D, and E with atomic number 3,11,15,18,19 respectively belong to the same group
(a) A, B, C
(b) B, C, D
(c) A, D, E
(d) A, B, E
Answer: (d) A, B, E
In simple words: Element A (atomic number 3, Lithium), Element B (atomic number 11, Sodium), and Element E (atomic number 19, Potassium) all belong to Group 1 of the periodic table. Elements in the same group share similar chemical properties.

🎯 Exam Tip: To determine if elements belong to the same group, write their electronic configurations. Elements with the same number of valence electrons will be in the same group.

Question 15. Id Loading [MathJax]/extensions/MathMenu.js ents in a group

🎯 Exam Tip: When faced with an incomplete question, try to identify any relevant keywords or concepts that might hint at the intended topic, although a full answer may not be possible.

Question 16. Two elements X and Y have (i) X has 17 protons and 18 neutrons (ii) Y has 17 protons and 20 neutrons Both X and Y are
(a) Isobars
(b) Isotopes
(c) Isotones
(d) None of the options
Answer: (b) Isotopes
In simple words: Elements with the same number of protons (atomic number) but different numbers of neutrons (and thus different mass numbers) are called isotopes. Here, both X and Y have 17 protons but different neutrons.

🎯 Exam Tip: Remember the definitions: isotopes have same Z, different A; isobars have same A, different Z; isotones have same neutrons, different Z and A.

Question 17. Mendeleev predicted the existence of two elements and named them as eka-silicon and eka-aluminium. Identify the elements which took their position at later stage
(a) Si and Ge
(b) Si and Ga
(c) Ge and Ga
(d) Si and Al
Answer: (c) Ge and Ga
In simple words: Mendeleev left gaps for elements he knew would be discovered. Eka-silicon was later identified as Germanium (Ge), and Eka-aluminium was identified as Gallium (Ga).

🎯 Exam Tip: Mendeleev's predictions for these elements were remarkably accurate, demonstrating the power of his periodic table structure based on properties.

Question 18. An element 'X' is placed in group 13 and third period of the Periodic Table. It burns in oxygen to form an oxide which is amphoteric in nature. Identify the chemical formula of its chloride
(a) CCl4
(b) BCl2
(c) GaCl3
(d) AlCl3
Answer: (d) AlCl3
In simple words: An element in Group 13, Period 3, that forms an amphoteric oxide is Aluminium (Al). Aluminium forms a chloride with the formula \( \text{AlCl}_3 \), as Aluminium has a valency of 3.

🎯 Exam Tip: Amphoteric oxides can react with both acids and bases. Aluminium oxide (\( \text{Al}_2\text{O}_3 \)) is a common example of an amphoteric oxide.

Question 20. In the Modern Periodic Table, calcium (Z = 20) is surrounded by the elements with atomic numbers 12, 19, 21, and 38 which of the following will have physical and chemical properties resembling calcium.
(a) 12, 20, 38
(b) 12,19, 20, 38
(c) 19,20,38
(d) 12,19, 20
Answer: (a) 12, 20, 38
In simple words: Calcium (Z=20) is in Group 2. Elements in the same group have similar properties. Magnesium (Z=12) is above Calcium, and Strontium (Z=38) is below Calcium in Group 2. Therefore, elements 12, 20, and 38 will have similar properties.

🎯 Exam Tip: Elements in the same group (vertical column) of the periodic table share similar chemical properties due to having the same number of valence electrons.

Question 21. An element X (2,8,2) combines separately with \( (\text{SO}_4)^{2-} \) and \( (\text{PO}_4)^{3-} \) radicals. The chemical formulae of the compounds are
(a) \( \text{X}_2\text{SO}_4: \text{X}_3(\text{PO}_4)_2 \)
(b) \( \text{XSO}_4: \text{X}_3(\text{PO}_4) \)
(c) \( \text{X}(\text{SO}_4)_2 : \text{X}_2(\text{PO}_4)_3 \)
(d) \( \text{XSO}_4: \text{X}_3(\text{PO}_4)_3 \)
Answer: (b) \( \text{XSO}_4: \text{X}_3(\text{PO}_4) \)
In simple words: Element X has an electronic configuration of 2,8,2, meaning it has a valency of 2. So, with sulfate \( (\text{SO}_4)^{2-} \) (also valency 2), it forms \( \text{XSO}_4 \). With phosphate \( (\text{PO}_4)^{3-} \) (valency 3), it forms \( \text{X}_3(\text{PO}_4)_2 \). The option presented in the answer choice (b) had a slight mathematical error in the formula for phosphate.

🎯 Exam Tip: Remember to use the criss-cross method to correctly determine the chemical formula by balancing the valencies of the element and the polyatomic ion.

Question 22. Two elements X and Y belong to group 1 and 2 respectively in the same period. The formulae of this oxides are
(a) XO, YO
(b) X2O, YO
(c) X2O, Y2O
(d) XO,Y2O

🎯 Exam Tip: Elements in Group 1 (alkali metals) have a valency of 1, and elements in Group 2 (alkaline earth metals) have a valency of 2. This difference in valency directly affects the stoichiometry of their oxides.

Question 24. Which of the following has maximum non-metallic character?
(a) F
(b) Br
(c) Cl
(d) I
Answer: (a) F
In simple words: Fluorine is the most non-metallic element here.

🎯 Exam Tip: Remember that non-metallic character generally increases across a period and decreases down a group.

Question 25. Arrange elements into the increasing order of their metallic character along a period.
(a) S < Si < P < Al
(b) S < P < Si < Al
(c) Si < P < S < Al
(d) Si < S < P < Al
Answer: (a) S < Si < P < Al
In simple words: Metallic character gets stronger as you move towards the left side of the periodic table in a period.

🎯 Exam Tip: Metallic character tends to increase when going down a group and decrease when moving from left to right across a period.

Question 26. Which of the following is not the characteristics of isotopes of an element? Isotopes of an element
(a) show same atomic mass
(b) show same atomic number
(c) occupy same position in the periodic table
(d)
Answer: (a) show same atomic mass
In simple words: Isotopes always have different atomic masses.

🎯 Exam Tip: Isotopes are atoms of the same element, meaning they have the same number of protons (atomic number) but different numbers of neutrons (and thus different atomic masses).

Question 1. Give one example of Dobereiner's Triad.
Answer: One example of a Dobereiner's Triad is Lithium (Li), Sodium (Na), and Potassium (K). In a triad, the atomic mass of the middle element is approximately the average of the other two.
In simple words: Lithium, Sodium, and Potassium form a Dobereiner's Triad.

🎯 Exam Tip: Remember Dobereiner's Triads were one of the earliest attempts to classify elements, grouping three elements with similar properties.

Question 2. How many triads could Dobereiner identify from the existing elements then?
Answer: Dobereiner was only able to identify three triads from the elements known at his time. This limited number of triads was a major reason why his classification system was not widely accepted.
In simple words: Dobereiner found only three triads of elements.

🎯 Exam Tip: Knowing the exact number of triads identified by Dobereiner can be a quick recall fact.

Question 3. What is the limitation of Dobereiner triads?
Answer: The main limitation of Dobereiner's triads was that he could not arrange all the known elements into triads that shared similar chemical properties. This meant his system was not universal and couldn't classify every element.
In simple words: Dobereiner's triads could not group all the elements together.

🎯 Exam Tip: Focus on the lack of universality as the key limitation of Dobereiner's classification.

Question 4. What was the basis of elements made by Newlands?
Answer: Newlands classified elements based on their increasing atomic masses. He noticed that every eighth element had similar properties to the first, much like the notes in a musical octave.
In simple words: Newlands put elements in order by their atomic masses.

🎯 Exam Tip: Connect Newlands' Law of Octaves to the musical scale for easy recall of the "every eighth element" rule.

Question 5. Give two limitations of Newlands' law of Octaves.
Answer: Here are two limitations of Newlands' law of Octaves: 1. This law only worked well for elements up to calcium. It did not apply to heavier elements. 2. To make elements fit his table, Newlands sometimes placed two different elements in the same slot. He also grouped elements with very different properties together. These inconsistencies showed that his law was not a perfect system for all elements.
In simple words: The law only worked for lighter elements like calcium, and sometimes he put different elements in the same box.

🎯 Exam Tip: Remember the two main flaws: its applicability only to lighter elements and the forced fitting of unlike elements.

Question 7. Which two chemical properties were considered by Mendeleev for grouping of elements?
Answer: Mendeleev considered two chemical properties to group elements: - The type of compounds elements formed with oxygen. - The type of compounds elements formed with hydrogen. These compounds, oxides and hydrides, are very stable and helped reveal the elements' fundamental reactivities.
In simple words: Mendeleev looked at how elements reacted with oxygen and hydrogen to group them.

🎯 Exam Tip: Focus on oxides and hydrides as these are crucial for understanding an element's bonding behavior.

Question 8. State Mendeleev's Periodic Law.
Answer: Mendeleev's Periodic Law states that the properties of elements repeat periodically when they are arranged according to their atomic masses. This was a groundbreaking idea that showed a pattern in the chemical behavior of elements.
In simple words: Mendeleev's law says element properties come back in a pattern when you list them by their atomic weight.

🎯 Exam Tip: Clearly distinguish Mendeleev's law (based on atomic mass) from the Modern Periodic Law (based on atomic number).

Question 9. Define 'groups and periods'.
Answer: In a periodic table, the vertical columns are known as groups, and the horizontal rows are called periods. Elements in the same group often share similar chemical properties, while elements in the same period show a gradual change in properties.
In simple words: Groups are up-and-down columns, and periods are side-to-side rows in the periodic table.

🎯 Exam Tip: Visualizing the periodic table helps to remember that groups are vertical and periods are horizontal.

Question 10. What is the formula of oxide and hydride of Group I elements?
Answer: For elements in Group I, the oxide formula is \( R_2O \), and the hydride formula is \( RH \). Here, 'R' stands for any element in Group I. Group I elements typically form compounds by losing one electron, leading to these specific formulas.
In simple words: Group I elements form oxides like \( R_2O \) and hydrides like \( RH \).

🎯 Exam Tip: Knowing the valency (usually 1 for Group I) helps deduce these general formulas for compounds.

Question 11. Name three elements discovered later, which filled gaps left by Mendeleev for them.
Answer: Three elements discovered later that filled the gaps left by Mendeleev were Scandium, Gallium, and Germanium. Mendeleev even predicted the properties of these unknown elements with remarkable accuracy.
In simple words: Scandium, Gallium, and Germanium were found later and fit into Mendeleev's table.

🎯 Exam Tip: Mendeleev's ability to predict undiscovered elements was a major success of his periodic table.

Question 13. How many groups and periods are present in the Modern Periodic Table?
Answer: The Modern Periodic Table contains 18 groups and 7 periods. Groups are vertical columns showing similar chemical properties, while periods are horizontal rows showing changing properties.
In simple words: The modern periodic table has 18 groups going down and 7 periods going across.

🎯 Exam Tip: It is crucial to remember the exact number of groups and periods for the Modern Periodic Table.

Question 14. What is the location of metals and non-metals in the Modern Periodic Table?
Answer: In the Modern Periodic Table, metals are typically found on the left side, while non-metals are located on the right side. A zig-zag line separates metals from non-metals, with metalloids found along this line.
In simple words: Metals are on the left of the periodic table, and non-metals are on the right.

🎯 Exam Tip: The position of an element in the periodic table gives a strong clue about whether it is a metal, non-metal, or metalloid.

Question 15. State Modern Periodic Law.
Answer: The Modern Periodic Law states that the properties of elements are a periodic function of their atomic number. This law correctly explains the patterns of element properties, unlike Mendeleev's law which used atomic mass.
In simple words: The modern law says element properties repeat based on their atomic number.

🎯 Exam Tip: Emphasize "atomic number" when stating the Modern Periodic Law to distinguish it from older laws.

Question 16. In Modern Periodic Table, what is common among all the elements in a group.
Answer: All elements in the same group of the Modern Periodic Table have the same number of valence electrons and therefore show similar chemical properties. This shared number of outermost electrons is why elements in a group behave chemically in similar ways.
In simple words: Elements in the same group have the same number of outer electrons and similar chemical behaviors.

🎯 Exam Tip: The number of valence electrons is the primary reason for similar chemical properties within a group.

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 (atomic number 9) has an electron configuration of 2,7. Chlorine (atomic number 17) has an electron configuration of 2, 8,7. Both fluorine and chlorine, being in Group 17, have 7 valence electrons. These 7 valence electrons make them highly reactive non-metals, often seeking one more electron to achieve a stable configuration.
In simple words: Fluorine and Chlorine both have 7 valence electrons, which means they have 7 electrons in their outermost shell.

🎯 Exam Tip: For Group A elements, the group number often directly indicates the number of valence electrons.

Question 18. Describe the electron capacity of different shells in the periodic table.
Answer: This table represents the electron capacity of shells.

The formula for maximum electrons in a shell is \( 2n^2 \), where 'n' is the period number (or principal quantum number). This means that as the period number increases, the electron capacity of the shells also increases.
In simple words: The table shows how many electrons each shell can hold. The first shell holds 2, the second holds 8, the third holds 18, and the fourth holds 32 electrons.

🎯 Exam Tip: Remember the \( 2n^2 \) formula to quickly calculate the maximum electron capacity of any principal shell.

Question 20. What is atomic size?
Answer: Atomic size is defined as the distance from the center of an atom's nucleus to its outermost electron shell. This size is very tiny and is typically measured in picometers, where one picometer is \( 10^{-12} \) meters.
In simple words: Atomic size is how big an atom is, measured from its center to its very outer edge.

🎯 Exam Tip: Clearly state the definition and the common unit of measurement (picometer) for atomic size.

Question 21. What happens to the size of atom as we move from left to right in a period.
Answer: As we move from left to right across a period in the periodic table, the atomic size of elements generally decreases. This occurs because the nuclear charge increases across a period, pulling the electrons closer to the nucleus.
In simple words: Atoms get smaller as you go from left to right in a period.

🎯 Exam Tip: Explain that the increased nuclear charge (more protons) with the same number of shells causes electrons to be pulled inward.

Question 22. How does the tendency to lose electrons will change in a period.
Answer: As we move across a period from left to right, the tendency of elements to lose electrons decreases. This is because the effective nuclear charge on the valence electrons increases, holding them more tightly.
In simple words: It becomes harder for elements to lose electrons as you move right across a period.

🎯 Exam Tip: Relate the tendency to lose electrons directly to metallic character, which also decreases across a period.

Question 24. How was the anomaly in arrangement of elements in the Mendeleev's Periodic Table removed?
Answer: The anomalies in Mendeleev's Periodic Table were resolved when elements were arranged based on their increasing atomic number, rather than atomic mass. This change, proposed by Moseley, led to the modern periodic table, which more accurately reflects element properties.
In simple words: Mendeleev's table problems were fixed by sorting elements by atomic number instead of atomic mass.

🎯 Exam Tip: Highlight the shift from atomic mass to atomic number as the crucial change that fixed Mendeleev's table.

Question 25. What are noble gases/inert gases?
Answer: Noble gases, also known as inert gases, are elements that are unreactive and do not easily combine with other elements. This is because their outermost electron shell is completely filled. Examples include Helium (He), Neon (Ne), Argon (Ar), and Xenon (Xe), which are very stable.
In simple words: Noble gases don't react much because their outer electron shells are full, like Helium or Neon.

🎯 Exam Tip: The key characteristic of noble gases is their complete outermost electron shell, leading to chemical inertness.

Question 26. Name two alkali metals present in Group I
Answer: Three alkali metals present in Group 1 are Lithium (Li), Sodium (Na), and Potassium (K). Alkali metals are highly reactive and tend to lose one electron to form a positive ion.
In simple words: Lithium, Sodium, and Potassium are alkali metals in Group 1.

🎯 Exam Tip: Alkali metals are always in Group 1 and are known for their high reactivity with water.

Question 27. An element 'X' belongs to II group and 2nd period. Write the atomic number and name of element.
Answer: An element 'X' belonging to Group II and Period 2 has an atomic number of 4. This element is Beryllium. Beryllium has an electron configuration of 2,2, meaning it has two shells and two valence electrons.
In simple words: Element 'X' in Group 2, Period 2 is Beryllium, with atomic number 4.

🎯 Exam Tip: The period number tells you the number of shells, and the group number (for representative elements) tells you the number of valence electrons.

Question 28. An element 'A' has atomic number 11, name the period and group number to which it belongs.
Answer: An element 'A' with atomic number 11 is Sodium (Na). Its electronic configuration is 2, 8, 1. It belongs to Period 3 because it has three electron shells, and to Group 1 because it has one valence electron. Sodium is an alkali metal, known for its high reactivity.
In simple words: Element 'A' with atomic number 11 is Sodium. It is in Period 3 and Group 1.

🎯 Exam Tip: Determine the electronic configuration first to easily identify the period and group of an element.

Question 29. An element 'P' belongs to group = 2 and period = 3, state whether it is a metal or non-metal and predict the nature of its oxides.
Answer: An element 'P' in Group 2 and Period 3 is a metal. Its oxides would be basic in nature. Group 2 elements are alkaline earth metals, which are known to form basic oxides.
In simple words: Element 'P' is a metal in Group 2, Period 3, and its oxides are basic.

🎯 Exam Tip: Elements in Group 1 and 2 are always metals, and their oxides are typically basic.

Question 30. The electronic configuration of an atom is 2, 8, 7. Give its atomic number, nature of oxide.
Answer: An atom with electronic configuration 2, 8, 7 has an atomic number of 17. This element is Chlorine (Cl), which is a non-metal, so its oxide would be acidic. Non-metal oxides typically react with water to form acids.
In simple words: Electronic configuration 2, 8, 7 means atomic number 17. This element forms an acidic oxide.

🎯 Exam Tip: Remember that non-metal oxides are generally acidic, while metal oxides are generally basic.

Question 31. An element belongs to group 13 and period 3, name the element and give its valency.
Answer: An element belonging to Group 13 and Period 3 is Aluminum (Al). Its valency is 3. Aluminum tends to lose its three valence electrons to form a \( +3 \) ion.
In simple words: The element in Group 13, Period 3 is Aluminum, and its valency is 3.

🎯 Exam Tip: For Group 13 elements, the valency is typically 3, as they tend to lose three valence electrons.

Question 32. What are metalloids? Give 2 examples.
Answer: Metalloids are elements that display properties of both metals and non-metals. They are also known as semi-metals. These elements are found along the zig-zag line that separates metals from non-metals in the periodic table.
In simple words: Metalloids are elements that act like both metals and non-metals. Examples are Boron and Silicon.

🎯 Exam Tip: Key examples of metalloids include Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, and Polonium.

Question 33. An element X belongs to group 17 and element Y belongs to group 1. What type of bond will they form?
Answer: Element X, belonging to Group 17, is a non-metal that readily accepts electrons. Element Y, belonging to Group 1, is an alkali metal that readily loses electrons. Therefore, they will form an ionic bond with the formula XY. Ionic bonds typically form between metals and non-metals due to the transfer of electrons.
In simple words: Element X (Group 17) and Element Y (Group 1) will form an ionic bond because one wants to lose electrons and the other wants to gain them.

🎯 Exam Tip: Ionic bonds occur when electrons are transferred, usually between a metal (electron donor) and a non-metal (electron acceptor).

Question 34. The following elements belong to same period arrange them in order. X Y Z Atomic Radius \( \rightarrow \)231 262 242
Answer: Given the atomic radii (Y=262, Z=242, X=231), the elements arranged in decreasing order of atomic radius within the same period are Y, Z, and X. Atomic radius generally decreases from left to right across a period because of the increasing nuclear charge pulling electrons closer.
In simple words: Arranging by size from biggest to smallest, the order is Y, then Z, then X.

🎯 Exam Tip: Remember that atomic radius decreases across a period because the increased nuclear charge pulls the electron shells closer, even with more electrons.

Question 35. What is the valency of magnesium with atomic number 12 and nitrogen with atomic number 7?
Answer: For Magnesium (atomic number 12), the electronic configuration is 2, 8, 2, so its valency is 2 (it loses 2 electrons). For Nitrogen (atomic number 7), the electronic configuration is 2, 5, so its valency is 3 (it gains 3 electrons). Valency is determined by the number of electrons an atom needs to lose or gain to achieve a stable outer shell.
In simple words: Magnesium has a valency of 2, and Nitrogen has a valency of 3.

🎯 Exam Tip: Valency is often the number of electrons gained or lost to achieve a full outermost electron shell (octet rule).

Question 36. How many shells are present in all the elements that belong to period 3?
Answer: All elements belonging to Period 3 of the periodic table have 3 electron shells (K, L, and M shells) where their electrons are distributed. The period number directly indicates the number of electron shells an atom has.
In simple words: All elements in Period 3 have 3 electron shells.

🎯 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: As we move from left to right across a period in the periodic table, the electropositive character of elements decreases. This is because elements on the right side of a period tend to gain electrons rather than lose them, exhibiting more non-metallic characteristics.
In simple words: Elements become less electropositive (less likely to lose electrons) as you move from left to right in a period.

🎯 Exam Tip: Electropositive character is the inverse of electronegativity; as one increases, the other decreases.

Question 39. What are halogens? Where are they located in the periodic table?
Answer: Halogens are a group of elements that readily react with metals to form salts. They are located in Group 17 of the Periodic Table. These elements have 7 valence electrons and are highly reactive non-metals.
In simple words: Halogens are salt-forming elements found in Group 17 of the periodic table.

🎯 Exam Tip: Recall that halogens need only one electron to complete their outer shell, making them very reactive.

Question 40. Atomic number of 4 elements is given below. Which element will belong to the group of Helium. W X Y Z 8 5 36 20
Answer: Among the given elements, Element Y with atomic number 36 (Krypton) will belong to the same group as Helium (He). Both Helium and Krypton are noble gases, located in Group 18, characterized by a stable, full outermost electron shell.
In simple words: Element Y, with atomic number 36, is Krypton. It is a noble gas like Helium, so it belongs to the same group.

🎯 Exam Tip: Noble gases (Group 18) are defined by their stable electron configuration, which is why elements like Helium and Krypton share the same group.

Atomic Theory, Periodic Classification, and Properties of Elements Short Answer Type Questions

Question 1. Why Mendeleev could not assign fixed position to hydrogen in the table?
Answer: Mendeleev could not give hydrogen a fixed place in his periodic table for two main reasons: (a) Hydrogen is similar to alkali metals because it reacts with halogens, oxygen, and sulfur to form compounds with similar formulas. (b) Hydrogen also resembles halogens because it exists as a diatomic molecule (H2) and can form covalent compounds with both metals and non-metals. This dual nature made it difficult to place hydrogen uniquely in his mass-based classification.
In simple words: Hydrogen acted like both alkali metals and halogens, so Mendeleev couldn't put it in just one spot.

🎯 Exam Tip: Remember hydrogen's unique position is due to its ability to both lose and gain an electron, behaving like Group 1 and Group 17 elements.

Question 3. State two characteristics of groups.
Answer: Here are two characteristics of elements within a group in the periodic table: 1. All elements in a group have the same number of valence electrons. This is why they show similar chemical properties. 2. The atomic radius increases as you move from top to bottom within a group, leading to a gradual change in their properties. The increasing atomic radius down a group is due to the addition of new electron shells.
In simple words: In a group, elements have the same outer electrons and similar properties, but their atomic size gets bigger as you go down.

🎯 Exam Tip: The number of valence electrons is the defining feature for chemical similarities within a group.

Question 4. What happens to the valency of elements as we move from left to right in a Periodic Table?
Answer: As we move from left to right across a period in the periodic table, the valency of elements first increases up to 4, and then it decreases down to zero. This trend is observed as elements try to achieve a stable electron configuration by either losing or gaining electrons.
In simple words: Valency first goes up to 4 and then down to zero as you move from left to right in a period.

🎯 Exam Tip: Visualize valency as the combining capacity, which changes systematically across a period.

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: The atomic size decreases as we move from left to right across a period, even though the number of electrons increases. This happens because all elements in a period have the same number of electron shells. As we move right, the number of protons in the nucleus increases, which leads to a stronger attractive force pulling the valence electrons closer to the nucleus, thus shrinking the atomic size. This stronger pull from the nucleus is known as increased effective nuclear charge.
In simple words: Atomic size shrinks across a period because more protons pull the same number of shells closer to the center.

🎯 Exam Tip: Emphasize that for a given period, the number of shells remains constant, so the increased nuclear charge is the dominant factor.

Question 6. What happens to the metallic character as we move from top to bottom in a group?
Answer: As we move from top to bottom in a group, the metallic character of elements increases. This is because the atomic size increases, making it easier for the outermost electrons to be lost, which is a key property of metals.
In simple words: Metallic character gets stronger as you go down a group because it becomes easier for atoms to lose electrons.

🎯 Exam Tip: Relate metallic character to the ease of electron loss; larger atoms lose electrons more readily.

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' with atomic number 17 is Chlorine (Cl). Its electronic configuration is 2, 8, 7. (a) Its valency is 1, as it needs to gain one electron to complete its outermost shell. (b) The formula of its halide would be HX, where H is hydrogen. (c) It forms a negative ion, specifically an anion (\( X^- \)). (d) Its reactivity is higher compared to other elements below it in the same group, as reactivity decreases down the halogen group. Elements in Group 17, like chlorine, are halogens and are highly reactive non-metals.
In simple words: For element X (atomic number 17): (a) Valency is 1. (b) Halide formula is HX. (c) It forms a negative ion. (d) It is the most reactive compared to elements below it in the group.

🎯 Exam Tip: For Group 17 (halogens), reactivity decreases down the group because electron-gaining tendency decreases with increasing atomic size.

Question 9. Why are noble gases placed in the separate group?
Answer: Noble gases are placed in a separate group (Group 18) because they all have the same valency of zero, meaning they are unreactive or inert. This also means they share similar chemical properties due to their complete outermost electron shells. This unique stability sets them apart from other elements that actively participate in chemical reactions.
In simple words: Noble gases are in their own group because they all have zero valency, are unreactive, and have similar chemical properties.

🎯 Exam Tip: The completed octet (or duet for Helium) is the primary reason for the unique placement and properties of noble gases.

Question 10. Following are the 4 elements W, X, Y and Z and their atomic numbers are 9, 10, 16, 17 respectively. Predict the following: (a) Two elements lying in same group. (b) Elements in second period.
Answer: Let's determine the electronic configuration, group, and period for each element: - **W (Atomic Number 9):** Electronic configuration (2, 7), Group 17, Period 2 (Fluorine) - **X (Atomic Number 10):** Electronic configuration (2, 8), Group 18, Period 2 (Neon) - **Y (Atomic Number 16):** Electronic configuration (2, 8, 6), Group 16, Period 3 (Sulfur) - **Z (Atomic Number 17):** Electronic configuration (2, 8, 7), Group 17, Period 3 (Chlorine)
(a) Two elements lying in the same group are **W and Z** (both in Group 17). (b) Elements in the second period are **W and X**. Elements in the same group share similar chemical behavior due to the same number of valence electrons.
In simple words: W and Z are in the same group. W and X are in the second period.

🎯 Exam Tip: The number of valence electrons defines the group, while the number of electron shells defines the period.

Question 11. State the difference between Modern Periodic Table and Mendeleev's Periodic Table.
Answer: Here are the key differences between Mendeleev's Periodic Table and the Modern Periodic Table:

The shift from atomic mass to atomic number as the basis for arrangement solved many inconsistencies observed in Mendeleev's table.
In simple words: Mendeleev's table used atomic mass, but the modern table uses atomic number. The modern table also has more groups and handles isotopes better.

🎯 Exam Tip: Clearly state the primary basis of classification (atomic mass vs. atomic number) and the number of groups/periods for both tables.

Question 12. Write all the elements present in third period of the periodic table and give their electronic configuration.
Answer: The elements present in the third period of the periodic table, along with their electronic configurations, are:

All elements in Period 3 have three electron shells, indicating their similar principle quantum number.
In simple words: The third period includes Sodium, Magnesium, Aluminum, Silicon, Phosphorus, Sulfur, Chlorine, and Argon, each with electrons arranged in three shells.

🎯 Exam Tip: Be able to list elements from any given period and correctly write their electronic configurations.

Question 13. How does electronic configuration helps us to locate the position of element in the Periodic Table?
Answer: The electronic configuration of an atom helps locate its position in the periodic table in two ways: - The number of valence electrons (electrons in the outermost shell) indicates the group number of the element. - The total number of electron shells an atom has determines the period number it belongs to. For example, an element with configuration 2, 8, 1 has 1 valence electron (Group 1) and 3 shells (Period 3).
In simple words: The number of outer electrons tells you the group, and the total number of electron shells tells you the period.

🎯 Exam Tip: Mastering electronic configuration is essential for understanding an element's position and properties in the periodic table.

Question 15. Why does the reactivity of metals increases and that of non-metals decreases as we move down the group?
Answer: The reactivity of metals increases as we move down a group because metallic reactivity depends on the ease of losing electrons. Down a group, atomic size increases, so the valence electrons are further from the nucleus and less tightly held, making them easier to remove and form positive ions. Conversely, the reactivity of non-metals decreases as we move down a group. Non-metallic reactivity depends on the tendency to gain electrons. Down a group, the atomic size increases, and the effective nuclear charge on the valence electrons decreases, making it harder for the atom to attract and gain new electrons. This opposing trend highlights the fundamental difference in how metals and non-metals achieve stability.
In simple words: Metals become more reactive down a group because they lose electrons more easily. Non-metals become less reactive down a group because they find it harder to gain electrons.

🎯 Exam Tip: Remember that metal reactivity is about losing electrons, while non-metal reactivity is about gaining electrons; these tendencies are affected differently by atomic size and nuclear charge down a group.

Question 16. List the elements present in 2nd period. Write their atomic number and electronic configuration.
Answer: The elements present in the 2nd period, along with their atomic numbers and corrected electronic configurations, are:

All elements in the second period have two electron shells.
In simple words: The second period elements are Lithium, Beryllium, Boron, Carbon, Nitrogen, Oxygen, Fluorine, and Neon, with their electron arrangements.

🎯 Exam Tip: Accurately writing electronic configurations is key to understanding elemental properties and positions.

Question 17. For the elements listed below (Al, Cl, K), predict their (a) Valency, (b) Period number, (c) Group number.
Answer: For the given elements, their valency, period number, and group number are:

The electronic configuration clearly shows the number of shells and valence electrons, which directly determine the period and group respectively.
In simple words: The table shows the valency, period, and group for Aluminum, Chlorine, and Potassium based on their atomic numbers.

🎯 Exam Tip: Practice deriving valency, period, and group from electronic configurations, as this is a fundamental skill in periodic table studies.

Question 18. Elements of group 1 are given below with their atomic number: Li (3) Na (11) K (19) (a) Give the order of their Atomic size. (b) Reactivity.
Answer: For the Group 1 elements Lithium (3), Sodium (11), and Potassium (19):

(a) The atomic size increases as we move down the group from top to bottom. This is because a new electron shell is added to each element. So, the order is \( Li < Na < K \). (b) The reactivity also increases as we move down the group. Larger atoms can lose their outermost electrons more easily because the nuclear force on these electrons is weaker. Thus, the reactivity order is also \( Li < Na < K \). Both atomic size and reactivity trends in Group 1 are directly related to the increasing distance between the nucleus and the valence electron down the group.
In simple words: For Group 1 elements, both atomic size and how easily they react increase as you go down the group. So Lithium is smallest and least reactive, and Potassium is largest and most reactive.

🎯 Exam Tip: In Group 1, increased atomic size down the group directly leads to decreased ionization energy and increased metallic reactivity.

Question 20. 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.
Answer: Carbon has an atomic number of 6, which means its electrons are arranged as 2, 4. Silicon has an atomic number of 14, with electrons arranged as 2, 8, 4. Both carbon and silicon have 4 valence electrons (electrons in their outermost shell). Because they have the same number of valence electrons, they are placed in the same group in the periodic table. This explains why they share some chemical properties despite being different types of elements.
In simple words: Carbon and silicon both have 4 electrons in their outer shell. Because of this, they are grouped together in the periodic table.

🎯 Exam Tip: Remember that elements in the same group share similar chemical properties primarily because they have the same number of valence electrons.

Question 21. What physical and chemical properties of elements were used by Mendeleev in creating his periodic table? List two observations which posed a challenge to Mendeleev's Periodic Law.
Answer: Mendeleev used two main things to make his periodic table. First, he used the atomic mass of each element as a physical property. Second, he looked at the chemical properties, specifically how elements formed compounds with oxygen (oxides) and hydrogen (hydrides). He saw that elements with similar chemical behaviors were grouped together.
However, two problems challenged his law. First, he could not always arrange elements strictly by increasing atomic mass, as some elements with higher mass came before those with lower mass to keep similar properties in a group. Second, he could not find a place for isotopes (atoms of the same element with different atomic masses) because his table was based on a single atomic mass for each element, but isotopes have different masses.
In simple words: Mendeleev used atomic mass and how elements react with oxygen and hydrogen to make his table. The problems were that he couldn't always put elements in order of their mass, and he had no space for isotopes.

🎯 Exam Tip: When discussing Mendeleev's challenges, focus on how his atomic mass-based classification sometimes conflicted with observed chemical similarities and the new discovery of isotopes.

Question 22. Table given below shows a part of the periodic table. Using this table explain why?
(a) Li and Na are considered as active metals.
(b) Atomic size of Mg is less than that of Na.
(c) Fluorine is more reactive than chlorine.
Answer:
(a) Lithium (Li) and Sodium (Na) are considered active metals because their atoms are relatively large. This large size means their outermost electrons are far from the nucleus and are not held very tightly, making it easy for them to lose these electrons. Losing electrons is a key characteristic of active metals, and these elements readily form positive ions.
(b) Magnesium (Mg) has a smaller atomic size than Sodium (Na). This is because both are in the same period, but Mg has more protons in its nucleus than Na. The stronger positive charge in Mg's nucleus pulls its electrons closer, making the atom smaller.
(c) Fluorine (F) is more reactive than Chlorine (Cl). This is because fluorine is smaller than chlorine, which means its nucleus is closer to its outer electrons. This strong attraction makes fluorine more eager to gain an electron to complete its outer shell, thus making it more reactive.
In simple words:
(a) Lithium and Sodium lose electrons easily because their atoms are big.
(b) Magnesium is smaller than sodium because it has more protons, pulling electrons in tighter.
(c) Fluorine is more reactive than chlorine because it's smaller and pulls electrons in more strongly.

🎯 Exam Tip: For reactivity and atomic size questions, always link your explanation to nuclear charge, number of shells, and the effective pull on valence electrons.

Question 23. The position of three elements A, B and C in the periodic table are shown below:

Giving reasons, explain the following:
(a) Element A is a non-metal.
(b) Element B has a larger atomic size than element C.
(c) Element C has a valency of 1.
Answer:
(a) Element A is a non-metal. This is because it is in Group 17, which means it has 7 electrons in its outermost shell. Elements in this group easily gain one electron to become stable and form negative ions. This tendency to gain electrons is typical of non-metals.
(b) Element B has a larger atomic size than element C. Both elements are in the same period (horizontal row), but B is to the left of C. As you move from left to right across a period, the atomic number increases, leading to a stronger pull from the nucleus on the electrons, which makes the atomic size decrease. So, B, being to the left, has fewer protons and a weaker nuclear pull, making its atom larger.
(c) Element C has a valency of 1. Since C is in Group 17, it has 7 valence electrons. To achieve a stable electron configuration (like a noble gas), it needs to gain just one more electron. The number of electrons an atom needs to gain or lose to become stable is its valency, which in this case is 1.
In simple words:
(a) Element A is a non-metal because it's in Group 17 and easily gains an electron.
(b) Element B is bigger than C. This is because B is to the left of C in the same row, meaning C has more protons pulling electrons tighter.
(c) Element C has a valency of 1. It has 7 outer electrons and just needs one more to be stable.

🎯 Exam Tip: Always remember that elements in Group 17 are halogens, highly reactive non-metals that readily gain one electron, defining their valency of 1.

Question 24. The position of three elements A, B and C

🎯 Exam Tip: When given a table of elements, use its structure to identify group and period numbers, which directly relate to valence electrons and number of shells.

Question 25. 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 a non-metal because it is in Group 17. Elements in Group 17 have 7 valence electrons and tend to gain one electron to achieve a stable electron configuration, forming negative ions. This behavior is characteristic of non-metals.
(b) The atom of element C has a larger size than A. Element C is in Period 3 (meaning it has 3 electron shells), while element A is in Period 2 (meaning it has 2 electron shells). Since C has one more electron shell than A, its atomic size is larger.
(c) Element B has a valency of 1. Element B is in Group 1. Elements in Group 1 have 1 valence electron, which they readily lose to become stable. The valency is determined by the number of electrons an atom loses, gains, or shares, so for Group 1 elements, it is 1.
In simple words:
(a) Element A is a non-metal because it's in Group 17 and easily gains an electron.
(b) Element C is bigger than A because C has 3 electron shells while A has only 2.
(c) Element B has a valency of 1 because it's in Group 1 and easily loses its one outer electron.

🎯 Exam Tip: Pay close attention to the row (period) and column (group) of an element as they directly indicate its electron shell structure and valence electron count, which are key to understanding its properties.

Question 26. The elements of the second period of the periodic table are given below:
Li Be B C N O F
(a) Give reason to explain why atomic radii decreases from Li to F.
(b) Identify the most
(i) Metallic and
(ii) Non-metallic element.
Answer:
(a) As we move from Lithium (Li) to Fluorine (F) across the second period, the atomic radii decrease. This happens because the number of electron shells stays the same, but the number of protons in the nucleus increases. The stronger positive charge of the nucleus pulls the outermost electrons closer, making the atom smaller. This stronger pull helps to hold the atom together more tightly.
(b)
(i) The most metallic element is Lithium (Li). Metallic character is highest on the left side of the periodic table, where elements easily lose electrons.
(ii) The most non-metallic element is Fluorine (F). Non-metallic character is highest on the right side of the periodic table (excluding noble gases), where elements strongly gain electrons.
In simple words:
(a) Atoms get smaller from Li to F because more protons pull the electrons closer, even though the number of electron shells stays the same.
(b) (i) Lithium is the most metallic. (ii) Fluorine is the most non-metallic.

🎯 Exam Tip: When asked about trends across a period, think about how increasing nuclear charge affects electron pull and atomic size, and how that influences metallic/non-metallic character.

Question 27. The elements of the third period of the periodic table are given below:
Group I II III IV V VI VII
Na Mg Al Si P S Cl
(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 (Na) atom is bigger in size than the Magnesium (Mg) atom. Both are in the same period. As we move from left to right across a period, the atomic number increases. This means there are more protons in the nucleus, which creates a stronger pull on the outermost electrons. This stronger pull makes the atomic size decrease. So, Na, being to the left of Mg, has fewer protons and a larger atomic size.
(b)
(i) The most metallic element in Period 3 is Sodium (Na). Metallic character is strongest for elements that easily lose electrons, which are found on the far left of the periodic table.
(ii) The most non-metallic element in Period 3 is Chlorine (Cl). Non-metallic character is strongest for elements that easily gain electrons, found on the far right of the periodic table (before noble gases).
In simple words:
(a) Sodium is bigger than Magnesium because it has fewer protons pulling its electrons, even though they are in the same row.
(b) (i) Sodium is the most metallic. (ii) Chlorine is the most non-metallic.

🎯 Exam Tip: Understanding atomic radii trends is crucial: they decrease across a period (due to increased nuclear charge) and increase down a group (due to added electron shells).

Question 28. (a) What is meant by periodicity in properties of elements with reference to the periodic table? (b) Why do all the elements of the same group have similar properties? (c) How will the tendency to gain electrons change as we go from left to right across a period? Why?
Answer:
(a) Periodicity in properties means that elements show similar chemical and physical properties at regular intervals when arranged in the periodic table. This repeating pattern helps us understand how elements behave.
(b) All elements in the same group (vertical column) have similar properties because they have the same number of valence electrons (electrons in their outermost shell). It is these valence electrons that largely determine an element's chemical behavior.
(c) As we move from left to right across a period, the tendency of elements to gain electrons increases. This is because the atomic size generally decreases, and the nuclear charge (number of protons) increases. A smaller atom with a stronger nuclear pull can attract and hold extra electrons more easily.
In simple words:
(a) Periodicity means element properties repeat in a pattern.
(b) Elements in the same group have similar properties because they have the same number of outer electrons.
(c) Elements gain electrons more easily when moving left to right in a period. This is because atoms get smaller and the nucleus pulls harder.

🎯 Exam Tip: When explaining periodicity, focus on the repeating pattern of valence electrons, which directly impacts chemical reactivity and other properties.

Atomic Theory, Periodic Classification, and Properties of Elements Long Answer Type Questions

Question 1. The atomic number of element X is 17 predict its physical state, its name, its hydride formula, its nature (metal/non-metal), and its molecular formula as a diatomic molecule.
Answer: Element X has an atomic number of 17. Its electronic configuration is 2, 8, 7. This means it has 7 electrons in its outermost shell, placing it in Group 17 and Period 3 of the periodic table.
Based on this, we can predict its properties:
(a) **Physical state:** It is a gas at room temperature, like other elements in its group (e.g., chlorine).
(b) **Name:** The element with atomic number 17 is Chlorine.
(c) **Formula of hydride:** It forms a hydride with the formula \( \text{HCl} \).
(d) **Nature:** It is a non-metal, as Group 17 elements are known non-metals.
(e) **Molecular formula:** It exists as a diatomic molecule, \( \text{Cl}_2 \). This element is a good example of how electron configuration dictates chemical behavior.
In simple words: Element X, with atomic number 17, is Chlorine. It's a gas, a non-metal, forms \( \text{HCl} \), and exists as \( \text{Cl}_2 \). It has 7 electrons in its outer shell.

🎯 Exam Tip: Knowing the atomic number allows you to determine electron configuration, group, and period, which are all crucial for predicting an element's properties.

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: Let's compare elements A (Group 1) and B (Group 2) from the same period:
(a) **Valency:** Element A has 1 valence electron, so its valency is 1. Element B has 2 valence electrons, so its valency is 2.
(b) **Size of atom:** Element A has a bigger atomic size than element B. As we move across a period from left to right, the atomic size decreases due to an increase in nuclear charge pulling the electron shells closer.
(c) **Formula of oxide:** Element A forms an oxide with the formula \( \text{A}_2\text{O} \). Element B forms an oxide with the formula \( \text{BO} \).
(d) **Nature of oxide:** The oxides of both A and B are basic in nature. This is because elements in Group 1 and 2 are metals, and metal oxides are typically basic.
(e) **Metallic character:** Element A is more metallic than element B. Metallic character generally decreases as you move from left to right across a period because elements become less likely to lose electrons.
In simple words: Comparing element A (Group 1) and B (Group 2) in the same row:
(a) A has a valency of 1, and B has a valency of 2.
(b) Atom A is bigger than atom B.
(c) Oxides are \( \text{A}_2\text{O} \) and \( \text{BO} \).
(d) Both oxides are basic.
(e) Element A is more metallic than B.

🎯 Exam Tip: When comparing elements in the same period, focus on how increasing nuclear charge affects properties like atomic size, valency, and metallic character from left to right.

Question 3. Give all characteristics of group.
Answer: Here are the main characteristics of elements within the same group (vertical column) of the periodic table:
1. **Valence electrons:** All elements in a group have the same number of valence electrons. These are the electrons in their outermost shell, which are key for chemical reactions.
2. **Valency:** Because they have the same number of valence electrons, all elements in a group show the same valency. This means they tend to gain, lose, or share the same number of electrons when forming chemical bonds.
3. **Atomic size:** The atomic size generally increases as you move down a group. This happens because new electron shells are added as you go down the group, making the atom larger.
4. **Metallic character:** The metallic character of elements also increases as you move down a group. This is because the outermost electrons are further from the nucleus and less strongly held, making them easier to lose, which is a key feature of metals.
In simple words: In a group, elements have the same number of outer electrons and the same valency. As you go down, atoms get bigger, and they become more metallic.

🎯 Exam Tip: Focus on how adding electron shells down a group influences atomic size and the ease of losing electrons, which directly impacts valency and metallic character.

Question 4. Explain how various properties like valence electrons, valency, atomic size, metallic character, and non-metallic character change as we move from left to right across a period in the periodic table.
Answer: In a period, as we move from left to right:
(a) **Valence electrons:** The number of valence electrons increases from 1 to 8. For example, it goes from 1, 2, 3, 4, 5, 6, 7, to 8.
(b) **Valency:** The valency first increases from 1 to 4, and then it decreases from 4 to 0. For example, 1, 2, 3, 4, then 3, 2, 1, 0.
(c) **Size of atom:** The atomic size generally decreases. This is due to the increasing nuclear charge pulling the electron shells closer.
(d) **Metallic character:** The metallic character decreases. Elements on the left are metals, and those on the right are non-metals.
(e) **Non-metallic character:** The non-metallic character increases. Elements become more likely to gain electrons. This organized change helps in predicting an element's behavior.
In simple words: When moving left to right in a period:
(a) Outer electrons increase from 1 to 8.
(b) Valency first goes up (1 to 4), then down (4 to 0).
(c) Atom size gets smaller.
(d) Metallic nature goes down.
(e) Non-metallic nature goes up.

🎯 Exam Tip: For trends across a period, remember that the number of electron shells remains constant while the nuclear charge increases, impacting all other properties.

Question 5. You are given five elements with some description of each element, place them in the Modern Periodic Table.
Answer: Here is how the five elements described can be placed in the Modern Periodic Table:

These descriptions help identify where each element fits based on its properties and atomic structure.
In simple words: The table shows five elements with their descriptions and where they fit in the periodic table by period and group. For example, Oxygen is in Period 2, Group 16.

🎯 Exam Tip: Connect an element's atomic number to its electron configuration, which then helps determine its period (number of shells) and group (number of valence electrons).

Question 6. An element 'P' belongs to group = 2 and period = 3, state whether it is a metal or non-metal and predict the nature of its oxides.
Answer: An element 'P' in Group 2 and Period 3 is Magnesium (Mg).
- **Metal or Non-metal:** Elements in Group 2 are always metals. So, element 'P' is a metal. Group 2 elements have two valence electrons, which they easily lose.
- **Nature of its oxides:** Since 'P' is a metal, its oxides will be basic. Basic oxides react with acids.
In simple words: Element P is a metal because it's in Group 2. Its oxides will be basic, which means they can react with acids.

🎯 Exam Tip: Remember the general rule: metal oxides are typically basic, while non-metal oxides are typically acidic or neutral.

Question 30. The electronic configuration of an atom is 2, 8, 7. Give its atomic number, nature of oxide.
Answer: An atom with the electronic configuration 2, 8, 7 has 2 + 8 + 7 = 17 electrons. For a neutral atom, the number of electrons is equal to the number of protons, so its **atomic number is 17**.
This element is Chlorine (Cl), which is a non-metal. Non-metals typically form **acidic oxides**. For example, chlorine forms oxides like \( \text{Cl}_2\text{O}_7 \), which is acidic.
In simple words: The atom has 17 electrons, so its atomic number is 17. It's a non-metal (Chlorine), and its oxides will be acidic.

🎯 Exam Tip: The sum of electrons in the electronic configuration gives the atomic number for a neutral atom, which then helps identify the element and its properties.

Question 31. An element belongs to group 13 and period 3, name the element and give its valency.
Answer: An element in Group 13 and Period 3 is Aluminium (Al).
- **Name of the element:** The element is Aluminium.
- **Valency:** Elements in Group 13 have 3 valence electrons. They tend to lose these 3 electrons to achieve a stable electron configuration. Therefore, its valency is 3. This means it usually forms positive ions with a +3 charge.
In simple words: The element is Aluminium. It is in Group 13 and Period 3, and its valency is 3 because it loses 3 electrons.

🎯 Exam Tip: For main group elements, the group number often directly relates to the number of valence electrons, which helps determine valency.

Question 32. What are metalloids? Give 2 examples.
Answer: Metalloids (also called semi-metals) are a special type of element that have properties that are in between those of metals and non-metals. For instance, they might look like metals but conduct electricity less efficiently. They can sometimes act like metals and sometimes like non-metals, depending on the chemical reaction.
Two examples of metalloids are:
1. Boron (B)
2. Silicon (Si)
Other examples include Germanium (Ge) and Arsenic (As).
In simple words: Metalloids are elements that act like both metals and non-metals. Boron and Silicon are two examples.

🎯 Exam Tip: Remember that metalloids are found along the zig-zag line separating metals and non-metals in the periodic table, showing intermediate properties.

Question 33. An element X belongs to group 17 and element Y belongs to group 1. What type of bond will they form?
Answer: Element X belongs to Group 17 (like chlorine), which means it has 7 valence electrons and will readily gain 1 electron to become stable. Element Y belongs to Group 1 (like sodium), meaning it has 1 valence electron and will readily lose 1 electron to become stable.
When an element that easily loses electrons (metal from Group 1) reacts with an element that easily gains electrons (non-metal from Group 17), they will form an **ionic bond**. Element Y will donate its single valence electron to element X.
The chemical formula for the compound formed will be **XY**. For example, if Y is Sodium (Na) and X is Chlorine (Cl), they form NaCl.
In simple words: Element X (Group 17, gains 1 electron) and Element Y (Group 1, loses 1 electron) will form an ionic bond. The formula will be XY.

🎯 Exam Tip: Ionic bonds typically form between metals (which lose electrons) and non-metals (which gain electrons), resulting in the transfer of electrons.

Question 34. The following elements belong to same period arrange them in order.
X Y Z
Atomic Radius → 231 262 242
Answer: The elements X, Y, and Z belong to the same period. We are given their atomic radii as 231, 262, and 242 (units like picometers are implied).
To arrange them in order, we use the trend of atomic radius across a period: atomic size decreases from left to right. This means the element with the largest atomic radius is on the left, and the smallest is on the right.
Given radii: Y (262), Z (242), X (231).
Arranging them in decreasing order of atomic radius (which corresponds to increasing atomic number across a period):
\( \text{Y} > \text{Z} > \text{X} \)
So the order from left to right in the period would be **Y, Z, X**.
In simple words: In a period, atoms get smaller from left to right. So, arranging them from largest to smallest atomic radius means putting them in the order Y, then Z, then X.

🎯 Exam Tip: Always remember that atomic radius decreases across a period because increased nuclear charge pulls the valence electrons closer to the nucleus.

Question 35. What is the valency of magnesium with atomic number 12 and nitrogen with atomic number 7?
Answer: Let's find the valency for both elements:
1. **Magnesium (Mg):** Its atomic number is 12. This means its electronic configuration is 2, 8, 2. Since it has 2 electrons in its outermost shell, it tends to lose these 2 electrons to become stable. Therefore, the **valency of magnesium is 2**.
2. **Nitrogen (N):** Its atomic number is 7. This means its electronic configuration is 2, 5. Since it has 5 electrons in its outermost shell, it tends to gain 3 electrons to complete its octet (become stable with 8 outer electrons). Therefore, the **valency of nitrogen is 3**. Valency tells us how many bonds an atom typically forms.
In simple words: Magnesium (atomic number 12) has 2 outer electrons, so its valency is 2. Nitrogen (atomic number 7) has 5 outer electrons, so its valency is 3 (it needs 3 more).

🎯 Exam Tip: Valency is the combining capacity of an element. For elements with 1, 2, or 3 valence electrons, it's usually the number of electrons lost; for 5, 6, 7, it's 8 minus the number of valence electrons.

Question 36. How many shells are present in all the elements that belong to period 3?
Answer: In the periodic table, the period number indicates the number of electron shells an atom has. For all elements that belong to **Period 3**, their atoms have **3 electron shells**. These shells are typically labeled as K, L, and M, and electrons are distributed among them. This is a fundamental rule for the organization of the periodic table.
In simple words: All elements in Period 3 have 3 electron shells (K, L, M) where their electrons are located.

🎯 Exam Tip: The period number of an element in the periodic table directly corresponds to the number of electron shells its atoms possess.

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: As we move from left to right across a period in the periodic table, the electropositive character of elements **decreases**. Electropositive character refers to an element's tendency to lose electrons.
This decrease happens because, across a period, the atomic number (number of protons) increases, leading to a stronger effective nuclear charge. This stronger positive pull from the nucleus holds the valence (outermost) electrons more tightly, making it harder for the atom to lose them. Therefore, elements on the left side (metals) are more electropositive than elements on the right side (non-metals).
In simple words: As you go from left to right in a period, elements become less electropositive. This means they are less likely to lose electrons because the nucleus pulls them more strongly.

🎯 Exam Tip: Electropositivity and metallic character follow the same trend: they decrease across a period and increase down a group.

Question 39. What are halogens? Where are they located in the periodic table?
Answer: Halogens are a group of highly reactive non-metallic elements. They are known for forming salts when they react with metals. Common examples include fluorine, chlorine, bromine, iodine, and astatine.
In the periodic table, halogens are located in **Group 17**. This group is found on the far right side of the table, just before the noble gases. Their position indicates they have 7 valence electrons, making them eager to gain one electron to achieve stability.
In simple words: Halogens are reactive non-metals that form salts with metals. They are found in Group 17 of the periodic table.

🎯 Exam Tip: Remember the specific groups: Group 1 (Alkali Metals), Group 2 (Alkaline Earth Metals), Group 17 (Halogens), and Group 18 (Noble Gases).

Question 40. Atomic number of 4 elements is given below. Which element will belong to the group of Helium.
W X Y Z
8 5 36 20
Answer: We are given four elements with their atomic numbers: W(8), X(5), Y(36), and Z(20). We need to find which one belongs to the same group as Helium (He), which is a noble gas (Group 18).
- **W (Atomic number 8):** Oxygen, electronic configuration 2, 6. Group 16.
- **X (Atomic number 5):** Boron, electronic configuration 2, 3. Group 13.
- **Y (Atomic number 36):** Krypton, electronic configuration 2, 8, 18, 8. This element has a completely filled outermost shell and is a noble gas.
- **Z (Atomic number 20):** Calcium, electronic configuration 2, 8, 8, 2. Group 2.
Therefore, **Element Y, with atomic number 36 (Krypton), will belong to the same group as Helium**. Both are inert gases, meaning they are very unreactive.
In simple words: Element Y, with atomic number 36 (Krypton), belongs to the same group as Helium (noble gases) because both are unreactive gases with full outer electron shells.

🎯 Exam Tip: Noble gases (Group 18) are characterized by a complete outermost electron shell, making them very stable and unreactive.

Atomic Theory, Periodic Classification, and Properties of Elements Short Answer Type Questions

Question 1. Why Mendeleev could not assign fixed position to hydrogen in the table?
Answer: Mendeleev faced a challenge in giving Hydrogen a fixed position in his periodic table because Hydrogen shows properties similar to both alkali metals (Group 1) and halogens (Group 17).
(a) **Similarity to Alkali Metals:** Hydrogen can form compounds with halogens, oxygen, and sulfur, just like alkali metals do. For example, it forms \( \text{HCl} \) (like \( \text{NaCl} \)) and \( \text{H}_2\text{O} \) (like \( \text{Na}_2\text{O} \)).
(b) **Similarity to Halogens:** Hydrogen exists as a diatomic molecule (\( \text{H}_2 \)), similar to halogens like \( \text{Cl}_2 \). It can also combine with metals and non-metals to form covalent compounds. This dual behavior made it difficult to place it uniquely.
In simple words: Mendeleev couldn't place hydrogen easily because it acts like both alkali metals and halogens, forming compounds and existing as a two-atom molecule like them.

🎯 Exam Tip: Hydrogen's unique dual nature (similarities to both Group 1 and Group 17 elements) is a common point of confusion in early periodic table classifications.

Question 2. (a) Why did Mendeleev leave gaps for undiscovered elements in his periodic table? (b) State any three limitations of Mendeleev's classification of elements. (c) How does the electronic configuration of atoms change in a period with an increase in atomic number?
Answer:
(a) Mendeleev left some gaps for undiscovered elements in his periodic table because he believed that these elements existed and would be discovered in the future. He even predicted the properties of these unknown elements, showing his foresight in developing the periodic law.
(b) Three limitations of Mendeleev's classification are:
* **Position of Hydrogen:** He could not assign a fixed or unique position to hydrogen, as it showed properties similar to both alkali metals and halogens.
* **Increasing Order of Atomic Mass:** The strict increasing order of atomic mass could not always be maintained. For example, cobalt (atomic mass 58.9) was placed before nickel (atomic mass 58.7).
* **Position of Isotopes:** Isotopes (atoms of the same element with different atomic masses) were not given separate places in the table, even though they have different masses, because they share the same chemical properties.
(c) In a period, as the atomic number increases from left to right, the inner core electron shells remain the same. However, the number of valence electrons (outermost shell electrons) increases progressively from left to right. This change in valence electrons is what primarily drives the change in an element's chemical properties across a period.
In simple words:
(a) Mendeleev left gaps for elements he knew would be found later, and he even guessed their properties.
(b) Three problems with his table were: hydrogen had no clear spot, elements weren't always in mass order, and isotopes (atoms of same element with different mass) had no separate space.
(c) Moving left to right in a period, atoms add more outer electrons, but the inner shells stay the same.

🎯 Exam Tip: When listing limitations of Mendeleev's table, ensure you include hydrogen's position, atomic mass anomalies, and the lack of space for isotopes.