Grade 10 Q&A: Chapter 2: Periodic Classification of Elements

Concept Questions

Q1: What is the modern periodic law?

Answer: The modern periodic law states that the physical and chemical properties of elements are periodic functions of their atomic numbers. This means that when elements are arranged in order of increasing atomic number, elements with similar properties occur at regular intervals.

Q2: How is the modern periodic law different from Mendeleev's periodic law?

Answer: Mendeleev's periodic law stated that the properties of elements are periodic functions of their atomic masses. The modern periodic law, proposed by Henry Moseley in 1913, states that properties are periodic functions of atomic numbers (number of protons). This resolved anomalies in Mendeleev's table where elements arranged by atomic mass didn't align with their chemical properties.

Q3: What are periods and groups in the periodic table?

Answer: Periods are horizontal rows in the periodic table, numbered from 1 to 7. Elements in the same period have the same number of electron shells. Groups are vertical columns, numbered from 1 to 18. Elements in the same group have similar chemical properties due to having the same number of valence electrons.

Q4: What are the four blocks of the periodic table?

Answer: The four blocks of the periodic table are: 1. s-block (Groups 1-2): Elements with valence electrons in s-orbitals 2. p-block (Groups 13-18): Elements with valence electrons in p-orbitals 3. d-block (Groups 3-12): Elements with valence electrons in d-orbitals 4. f-block (Lanthanides and Actinides): Elements with valence electrons in f-orbitals

Q5: What are the main achievements of Mendeleev's periodic table?

Answer: The main achievements of Mendeleev's periodic table include: 1. Prediction of properties of undiscovered elements (like gallium, scandium, and germanium) 2. Correction of atomic masses of some elements 3. Accommodation of noble gases when they were discovered later 4. Providing a systematic arrangement of elements based on their properties 5. Establishing the concept of periodicity in element properties

Q6: How does atomic size vary across a period and down a group?

Answer: Across a period (left to right), atomic size decreases due to increasing nuclear charge pulling electrons closer while the number of electron shells remains the same. Down a group (top to bottom), atomic size increases due to the addition of new electron shells, which outweighs the effect of increased nuclear charge.

Q7: What is ionization energy and how does it vary in the periodic table?

Answer: Ionization energy is the energy required to remove an electron from a gaseous atom in its ground state. Across a period (left to right), ionization energy generally increases due to increasing nuclear charge and decreasing atomic size. Down a group (top to bottom), ionization energy decreases due to increasing atomic size and greater shielding effect.

Q8: What are the characteristics of alkali metals?

Answer: Alkali metals (Group 1) have these characteristics: 1. One valence electron (ns¹ configuration) 2. Soft, silvery metals with low melting and boiling points 3. Highly reactive, stored in oil to prevent reaction with air 4. Form ionic compounds with non-metals 5. React vigorously with water to form hydroxides and hydrogen gas 6. Strong reducing agents 7. Form +1 ions easily

Q9: What are the characteristics of halogens?

Answer: Halogens (Group 17) have these characteristics: 1. Seven valence electrons (ns²np⁵ configuration) 2. Exist as diatomic molecules (F₂, Cl₂, Br₂, I₂) 3. Highly reactive non-metals 4. Form salts with metals 5. Strong oxidizing agents 6. Reactivity decreases down the group 7. Form -1 ions easily

Q10: Why do elements in the same group have similar chemical properties?

Answer: Elements in the same group have similar chemical properties because they have the same number of valence electrons in their outermost shell. Since chemical properties are primarily determined by valence electron configuration, elements with the same number of valence electrons tend to form similar types of compounds and undergo similar reactions.

Application-Based Questions

Q11: How would you identify whether an element is a metal, non-metal, or metalloid based on its position in the periodic table?

Answer: - Metals: Located on the left and center of the periodic table (Groups 1-2, 3-12, and parts of 13-16). They typically have 1-3 valence electrons. - Non-metals: Located on the upper right side of the periodic table (parts of Groups 13-17 and all of Group 18). They typically have 4-8 valence electrons. - Metalloids: Form a diagonal border between metals and non-metals (B, Si, Ge, As, Sb, Te, Po). They show properties of both metals and non-metals.

The "staircase" line running from boron (B) to polonium (Po) roughly separates metals (to the left) from non-metals (to the right), with metalloids along this boundary.

Q12: Explain why sodium (Na) is more reactive than potassium (K) with water, even though both are in Group 1.

Answer: This statement is incorrect. Potassium (K) is actually more reactive with water than sodium (Na), not less. Down Group 1, reactivity increases from Li to Cs because: 1. Atomic size increases, making it easier to lose the valence electron 2. The valence electron is farther from the nucleus and experiences more shielding 3. The ionization energy decreases down the group 4. The larger atoms have lower charge density, making the +1 ion more stable

Therefore, potassium reacts more vigorously with water than sodium, producing a more violent reaction with faster hydrogen gas evolution.

Q13: Why does fluorine (F) have a higher electronegativity than chlorine (Cl), even though chlorine has more protons?

Answer: Fluorine has higher electronegativity than chlorine despite having fewer protons because: 1. Fluorine has a smaller atomic radius than chlorine (F is in period 2, Cl is in period 3) 2. The valence electrons in fluorine are closer to the nucleus 3. The nuclear attraction for shared electrons is stronger in fluorine due to less shielding 4. Fluorine has 7 valence electrons in a smaller shell, creating a stronger pull on shared electrons

The effect of atomic size outweighs the effect of nuclear charge in this case, making fluorine the most electronegative element in the periodic table.

Q14: How would you explain the fact that lithium forms more covalent compounds than other alkali metals?

Answer: Lithium forms more covalent compounds than other alkali metals due to: 1. Its small size, which leads to high charge density in the Li⁺ ion 2. Higher polarizing power (ability to distort electron clouds of anions) 3. Lower electropositive character compared to other alkali metals 4. Higher ionization energy, making it relatively harder to lose an electron completely 5. The absence of d-orbitals in its valence shell, which affects its bonding behavior

These factors make lithium exhibit diagonal relationships with magnesium and form covalent bonds with highly electronegative elements more readily than other alkali metals.

Q15: Why are noble gases used in lighting and signs, and what property makes them suitable for this purpose?

Answer: Noble gases are used in lighting and signs because: 1. They are chemically inert and don't react with other materials in the lighting fixture 2. When electricity passes through them, their electrons get excited to higher energy levels 3. When these electrons return to their ground state, they emit photons (light) of specific wavelengths 4. Different noble gases emit different colors: neon (red-orange), argon (blue/lavender), helium (pink), krypton (white-violet), xenon (blue/green) 5. They have low reactivity, ensuring long-lasting performance 6. They exist as monatomic gases, simplifying the emission spectrum

Their complete valence electron shells make them stable and unreactive, while their electron configuration allows for predictable and characteristic light emission when electrically excited.

Higher-Order Thinking Questions

Q16: Compare and contrast the periodic table developed by Mendeleev with the modern periodic table.

Answer:

Mendeleev's Periodic Table: 1. Basis of arrangement: Elements arranged by increasing atomic mass 2. Prediction of elements: Left gaps for undiscovered elements and predicted their properties 3. Anomalies: Some elements were placed out of atomic mass order to match chemical properties 4. Structure: Had 8 groups (vertical columns) and 7 periods (horizontal rows) 5. Noble gases: Not included initially (discovered later) 6. Theoretical basis: Based on empirical observations without understanding of atomic structure 7. Isotopes: Could not explain the existence of isotopes

Modern Periodic Table: 1. Basis of arrangement: Elements arranged by increasing atomic number 2. Prediction of elements: Based on electronic configuration and quantum numbers 3. Anomalies: No anomalies in arrangement by atomic number 4. Structure: Has 18 groups and 7 periods, with s, p, d, and f blocks 5. Noble gases: Included as Group 18 6. Theoretical basis: Based on quantum mechanical model of the atom 7. Isotopes: Can explain isotopes as atoms with same atomic number but different mass numbers

Key Improvements: - The modern table resolves anomalies in Mendeleev's table by using atomic number instead of atomic mass - It provides a theoretical explanation for periodicity based on electronic configuration - It accommodates all known elements in a logical structure - It explains chemical behavior based on valence electron configuration

Q17: Explain why the first ionization energy generally increases across a period but decreases down a group, and discuss any exceptions to this trend.

Answer:

General Trend Explanation: - Across a period (left to right): Ionization energy generally increases because: 1. Nuclear charge increases while the number of electron shells remains the same 2. Effective nuclear charge increases as electrons are added to the same shell 3. Atomic radius decreases, bringing valence electrons closer to the nucleus 4. Shielding effect remains relatively constant within the same period

• Down a group (top to bottom): Ionization energy generally decreases because:
• Atomic radius increases due to additional electron shells
• Valence electrons are farther from the nucleus
• Inner electrons provide increased shielding from nuclear charge
• Effective nuclear charge experienced by valence electrons decreases

Exceptions to the Trend: 1. Between Groups 2 and 13: There's a slight decrease in ionization energy from Be to B, Mg to Al, etc., because: - Group 13 elements have their outermost electron in a p-orbital, which is slightly higher in energy and more shielded than s-orbitals - The slight decrease in effective nuclear charge outweighs the effect of increasing atomic number

1. Between Groups 15 and 16: There's a slight decrease in ionization energy from N to O, P to S, etc., because:
2. Group 16 elements have paired electrons in one of their p-orbitals
3. Electron-electron repulsion in the paired orbital makes it slightly easier to remove an electron
4. This effect outweighs the increase in nuclear charge

These exceptions demonstrate that while general trends exist, the actual behavior of elements depends on the complex interplay of nuclear charge, electron configuration, and electron-electron interactions.

Q18: Discuss how the concept of electronegativity helps in predicting the nature of chemical bonds and the properties of compounds.

Answer: Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. This concept is crucial for predicting:

1. Bond Type: - Large electronegativity difference (>1.7): Ionic bond forms (electron transfer) - Moderate electronegativity difference (0.5-1.7): Polar covalent bond forms (unequal sharing) - Small electronegativity difference (<0.5): Non-polar covalent bond forms (equal sharing)

2. Bond Polarity: - The greater the electronegativity difference, the more polar the bond - The direction of the dipole points toward the more electronegative element - Polarity affects intermolecular forces and physical properties like boiling point

3. Compound Properties: - Solubility: Polar compounds dissolve in polar solvents; non-polar in non-polar solvents - Reactivity: Compounds with highly polar bonds are often more reactive - Acid-Base Behavior: Electronegativity helps predict whether a compound will act as an acid or base - Oxidation-Reduction: More electronegative elements tend to be oxidizing agents

4. Molecular Geometry: - Combined with VSEPR theory, electronegativity helps predict molecular shapes - Lone pairs on electronegative atoms influence bond angles

5. Reactivity Patterns: - Helps predict sites of nucleophilic or electrophilic attack in organic reactions - Explains why certain functional groups behave similarly across different molecules

Practical Applications: - Designing catalysts with specific electronic properties - Developing pharmaceuticals with desired solubility and reactivity - Creating materials with specific electrical or thermal properties - Understanding biological processes like enzyme-substrate interactions

The periodic trends in electronegativity (increasing across periods, decreasing down groups) provide a systematic framework for predicting these properties across the periodic table.

Q19: Analyze the anomalous behavior of the first element in Groups 13-17 compared to other elements in their respective groups, and explain the underlying reasons.

Answer: The first elements of Groups 13-17 (B, C, N, O, F) show anomalous behavior compared to other members of their groups due to several factors:

Common Factors for Anomalous Behavior: 1. Small Size: First-period elements have much smaller atomic radii 2. Absence of d-orbitals: Cannot expand their octet or use d-orbitals for bonding 3. High Electronegativity: Generally more electronegative than other group members 4. Strong Bond Formation: Form stronger bonds due to better orbital overlap 5. High Charge Density: Their ions have high charge-to-radius ratios

Group-Specific Anomalies:

Boron (Group 13): - Forms covalent compounds rather than ionic compounds like other group members - Acts as a Lewis acid (electron pair acceptor) due to vacant p-orbital - Forms electron-deficient compounds (e.g., BF₃ with trigonal planar geometry) - Cannot form B³⁺ ions unlike Al³⁺, Ga³⁺, etc. - Shows diagonal relationship with silicon

Carbon (Group 14): - Exceptional ability to catenate (form chains with itself) - Forms multiple bonds readily (C=C, C≡C) unlike other group members - Exists in multiple allotropic forms (diamond, graphite, fullerenes) - Forms stronger bonds with other elements - Cannot form C⁴⁺ ions unlike other group members that can form M⁴⁺ or M²⁺

Nitrogen (Group 15): - Exists as a diatomic gas (N₂) with a triple bond, unlike other group members - Less reactive due to the strong N≡N triple bond - Forms hydrogen bonds in NH₃, unlike other group hydrides - Cannot form pentahalides (NX₅) due to absence of d-orbitals - Shows limited catenation compared to phosphorus

Oxygen (Group 16): - Exists as a diatomic gas (O₂), unlike other group members - Forms stronger hydrogen bonds than other group members - Cannot expand its octet to form compounds like SO₃ or SF₆ - More electronegative and forms more ionic compounds - Forms peroxides (O₂²⁻) more readily than other group members

Fluorine (Group 17): - Most electronegative element in the periodic table - Strongest oxidizing agent among halogens - Forms only -1 oxidation state, unlike other halogens that can form multiple oxidation states - Cannot act as a central atom in compounds due to small size - Forms the strongest hydrogen bonds among halogens

These anomalies demonstrate how atomic size, electronic configuration, and orbital availability fundamentally affect chemical behavior, creating exceptions to general group trends.

Q20: How has the periodic table evolved over time, and what might future versions of the periodic table look like as new elements are discovered or synthesized?

Answer:

Historical Evolution of the Periodic Table:

1. Early Classifications (1800s):
2. Döbereiner's Triads (1829): Elements grouped in threes with similar properties
3. Newlands' Law of Octaves (1864): Every eighth element had similar properties

4. Lothar Meyer's work (1869): Arranged elements by atomic volume showing periodicity

5. Mendeleev's Table (1869):

6. Arranged elements by atomic mass
7. Left gaps for undiscovered elements
8. Predicted properties of unknown elements

9. Occasionally placed elements out of atomic mass order

10. Discovery of Noble Gases (1890s):

11. Added as a new group to Mendeleev's table

12. Challenged existing arrangements

13. Moseley's Contribution (1913):

14. Established atomic number as the fundamental ordering principle
15. Resolved anomalies in Mendeleev's table

16. Led to the modern periodic law

17. Modern Periodic Table Development:

18. Glenn Seaborg's actinide concept (1940s)
19. Recognition of the f-block elements
20. Standardization by IUPAC
21. Extension to include synthesized superheavy elements

Current Challenges and Future Possibilities:

1. Extension Beyond Current Elements:
2. Theoretical "island of stability" around elements 114-126
3. Potential for elements with atomic numbers >118
4. Increasing difficulty in synthesis and detection

5. Extremely short half-lives limiting practical study

6. Alternative Arrangements:

7. Three-dimensional periodic tables to better represent electron configurations
8. Spiral arrangements to emphasize periodicity
9. "Long-form" tables that better represent the f-block elements

10. Interactive digital formats allowing multiple views and data layers

11. Theoretical Considerations:

12. Relativistic effects becoming dominant in superheavy elements
13. Potential breakdown of periodic trends at extreme atomic numbers
14. Quantum mechanical limitations on possible electron configurations

15. Possible new types of orbitals or electron behavior

16. Practical Innovations:

17. Integration of more data (isotopes, nuclear properties, applications)
18. Customizable views based on specific properties
19. Incorporation of sustainability and supply risk information

20. Educational versions with different levels of complexity

21. Interdisciplinary Approaches:

22. Connections to astrophysics (element formation in stars)
23. Integration with materials science databases
24. Quantum computing applications for predicting properties
25. Artificial intelligence to identify new patterns and relationships

The periodic table remains a living scientific tool that continues to evolve as our understanding of atomic structure deepens and as new elements are discovered or synthesized. Future versions will likely maintain the core organizational principles while incorporating new dimensions of information and possibly extending to accommodate elements with properties we cannot yet fully predict.

Numerical Problems

Q21: Calculate the number of valence electrons in the following elements: Na, Mg, Al, Si, P.

Answer: To determine the number of valence electrons, we need to identify the group number or examine the electron configuration:

• Sodium (Na): Electronic configuration is [Ne]3s¹ Number of valence electrons = 1 (Group 1 element)

• Magnesium (Mg): Electronic configuration is [Ne]3s² Number of valence electrons = 2 (Group 2 element)

• Aluminum (Al): Electronic configuration is [Ne]3s²3p¹ Number of valence electrons = 3 (Group 13 element)

• Silicon (Si): Electronic configuration is [Ne]3s²3p² Number of valence electrons = 4 (Group 14 element)

• Phosphorus (P): Electronic configuration is [Ne]3s²3p³ Number of valence electrons = 5 (Group 15 element)

Q22: If the first ionization energy of sodium is 496 kJ/mol and that of magnesium is 738 kJ/mol, calculate the percentage increase in ionization energy from sodium to magnesium.

Answer: Given: - First ionization energy of sodium (Na) = 496 kJ/mol - First ionization energy of magnesium (Mg) = 738 kJ/mol

To calculate the percentage increase: Percentage increase = [(Final value - Initial value) / Initial value] × 100% Percentage increase = [(738 - 496) / 496] × 100% Percentage increase = [242 / 496] × 100% Percentage increase = 0.4879 × 100% Percentage increase = 48.79%

Therefore, the percentage increase in first ionization energy from sodium to magnesium is approximately 48.79%.

Q23: The atomic radii of three elements X, Y, and Z are 186 pm, 104 pm, and 143 pm respectively. If Y is in period 2 and Z is in period 3, determine the period in which element X is likely to be found.

Answer: Given: - Atomic radius of X = 186 pm - Atomic radius of Y = 104 pm (in period 2) - Atomic radius of Z = 143 pm (in period 3)

Atomic radius generally increases down a group and decreases across a period. Since: - Z (143 pm) is in period 3 and has a larger radius than Y (104 pm) in period 2, this confirms the trend of increasing radius down a group. - X has the largest radius (186 pm), which suggests it is in a period below period 3.

Since X's radius (186 pm) is significantly larger than Z's radius (143 pm), and Z is in period 3, element X is likely to be in period 4.

Q24: Elements A, B, C, and D have atomic numbers 9, 11, 17, and 19 respectively. Which of these elements would have the highest electronegativity, and which would have the lowest? Explain your reasoning.

Answer: Given: - Element A: Atomic number 9 (Fluorine, F) - Element B: Atomic number 11 (Sodium, Na) - Element C: Atomic number 17 (Chlorine, Cl) - Element D: Atomic number 19 (Potassium, K)

To determine electronegativity, we need to consider the element's position in the periodic table:

• Fluorine (F): Group 17, Period 2
• Sodium (Na): Group 1, Period 3
• Chlorine (Cl): Group 17, Period 3
• Potassium (K): Group 1, Period 4

Electronegativity generally increases across a period (left to right) and decreases down a group (top to bottom).

Highest Electronegativity: Fluorine (F) would have the highest electronegativity because: 1. It is in Group 17 (highly electronegative group) 2. It is in Period 2 (smaller atomic radius) 3. It has a high effective nuclear charge pulling on its valence electrons 4. Fluorine is actually the most electronegative element in the periodic table

Lowest Electronegativity: Potassium (K) would have the lowest electronegativity because: 1. It is in Group 1 (alkali metals have low electronegativity) 2. It is in Period 4 (larger atomic radius) 3. Its valence electron is far from the nucleus and well-shielded 4. It readily loses its valence electron to form a +1 ion

Therefore, the order of electronegativity from highest to lowest is: Fluorine (F) > Chlorine (Cl) > Sodium (Na) > Potassium (K)

Q25: If element X has 3 valence electrons and element Y has 6 valence electrons, predict the formula of the compound they would form and describe the nature of the bond.

Answer: Given: - Element X has 3 valence electrons - Element Y has 6 valence electrons

Step 1: Identify the likely groups of these elements. - Element X with 3 valence electrons is likely from Group 13 (like Al, Ga) - Element Y with 6 valence electrons is likely from Group 16 (like O, S)

Step 2: Determine how many electrons each element needs to achieve a stable octet. - Element X needs to lose 3 electrons or gain 5 electrons to achieve stability - Element Y needs to gain 2 electrons to achieve stability

Step 3: Determine the most energetically favorable way to achieve stability. - It's more favorable for X to lose 3 electrons and for Y to gain 2 electrons - This means X will form a +3 ion and Y will form a -2 ion

Step 4: Determine the formula based on charge balance. - To balance the charges (+3 for X and -2 for Y), we need 2 X atoms (+6 total) and 3 Y atoms (-6 total) - Therefore, the formula would be X₂Y₃

Step 5: Describe the nature of the bond. - The large difference in electronegativity between a Group 13 element and a Group 16 element suggests an ionic bond - The bond would be predominantly ionic, with electrons being transferred from X to Y - There may be some covalent character due to polarization effects

Therefore, the compound would have the formula X₂Y₃ (like Al₂O₃ or aluminum oxide) with predominantly ionic bonding.

References

1. Maharashtra State Board 10th Standard Science Syllabus 2025-26
2. NCERT Science Textbook for Class 10
3. Royal Society of Chemistry - Periodic Table Resources
4. American Chemical Society - Periodic Table Information
5. Journal of Chemical Education - Periodic Table Trends and Properties