The Periodic Table

The Periodic Table is a masterful chart that organizes all known chemical elements based on their fundamental properties, providing a systematic framework for understanding matter. Its foundation lies in the Periodic Law, which states that when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic repetition. This ingenious arrangement was primarily developed by Dmitri Mendeleev, who had the remarkable insight to leave gaps for elements yet to be discovered, accurately predicting their properties. The table is structured into vertical columns called ‘groups’ and horizontal rows known as ‘periods’. Each element is represented by its unique chemical symbol, atomic number (number of protons), and atomic mass, offering a complete identity card for every building block of the universe.

The modern Periodic Table is powerfully divided into metals, non-metals, and metalloids, with a clear segregation that helps predict an element’s behavior. As one moves from left to right across a period, a key trend is observed: the atomic size generally decreases due to the increasing pull of the nucleus on the electron shells, while the tendency to gain electrons (non-metallic character) increases. Conversely, moving down a group, the atomic size increases as new electron shells are added, and the metallic character strengthens because the outer electrons are farther from the nucleus and more easily lost. This predictable variation in properties like valency, electronegativity, and metallic nature across periods and down groups is what makes the table an indispensable tool for chemists, allowing them to anticipate how elements will interact and bond with one another.

This systematic classification is not just theoretical; it has immense practical value. It allows scientists to logically understand the relationships between elements, such as why elements in the same group, like the halogens (Group 17) or alkali metals (Group 1), share strong chemical similarities. By simply knowing an element’s position, one can deduce its likely valency, reactivity, and the type of compounds it will form. For instance, knowing that oxygen is in Group 16 helps us predict that it forms a similar type of compound with hydrogen (H₂O) as sulfur (H₂S), its group member below. Thus, the Periodic Table serves as the ultimate guide map for the entire field of chemistry, elegantly illustrating the order inherent in the natural world.

Exercise 5 (A)

Question 1. 

What is the need for classification of elements?

Ans:

The classification of elements is crucial for three main reasons:

1. Organization from Chaos: It transforms a random list of elements into a logical system. Instead of memorizing each element individually, we can see them as part of a larger, interconnected pattern.
2. Predictive Power: The primary strength of a good classification (like the Periodic Table) is its ability to predict. By knowing an element’s position, we can accurately predict its physical and chemical behavior, how it will react, and what kinds of compounds it will form.
3. Revealing Periodic Trends: It systematically reveals trends in properties like atomic size, reactivity, and electronegativity. This allows us to understand why elements behave differently and how their properties change progressively.

Question 2. 

What was the basis of the earliest attempts made for classification and grouping of elements?

Ans:

The earliest attempts to classify elements were based on the most obvious and observable characteristics available to scientists at the time. Initially, this meant grouping elements into broad categories like metals and non-metals based on their simple physical properties, such as lustre, malleability, and conductivity.

Later, as more elements were discovered, chemists sought a more scientific basis than just appearance. A significant early effort was made by Johann Döbereiner, who grouped elements with similar properties into sets of three, called ‘Triads’. He observed that the atomic weight of the middle element in a triad was roughly the average of the other two. While limited, this was a crucial step as it established a relationship between atomic weight and elemental properties, paving the way for the modern Periodic Table.

Question 3. 

1. A, B and C are the elements of a Dobereiner’s triad. If the atomic mass of A is 7 and that of C is 39, what should be the atomic mass of B? 

2. Why was Dobereiner’s triad discarded?

Ans:

1. Atomic Mass of B

According to Dobereiner’s law of triads, the atomic mass of the middle element (B) in a triad is approximately the average of the atomic masses of the other two elements (A and C).

Given:
Atomic mass of A = 7
Atomic mass of C = 39

Calculation:
Atomic mass of B = (Mass of A + Mass of C) / 2
= (7 + 39) / 2
= 46 / 2
= 23

Therefore, the atomic mass of element B is 23.

2. Reason for Discarding Dobereiner’s Triads

Dobereiner’s classification was ultimately discarded because it was too limited and could not be universally applied. He could identify triads for only a small group of elements, while the majority of known elements did not fit this pattern. As more elements were discovered, the model failed to provide a comprehensive framework for organizing all of them, leading to its rejection in favor of more complete classifications like Newlands’ Law of Octaves and ultimately, Mendeleev’s Periodic Table.

Question 4. 

Explain ‘Newland’s Law of Octaves.’ Why was the law discarded?

Ans:

Explaining Newland’s Law of Octaves

In the mid-19th century, well before Dmitri Mendeleev created the modern periodic table, chemists were searching for patterns to bring order to the growing number of known elements. One of the most intriguing and poetic attempts came from the English chemist John Newlands in 1864.

Newlands proposed that when elements are arranged in order of increasing atomic weight, every eighth element exhibits similar properties, much like the repeating octaves in a musical scale where the eighth note echoes the first.

The Core Idea:

• Newlands took the known elements (there were about 56 at the time) and started listing them by ascending atomic weight.
• He noticed that the first and the eighth element shared chemical similarities, the second and the ninth shared similarities, and so on.
• He called this pattern the “Law of Octaves.”

A Simple Analogy:

Think of a piano. If you start with the note C, the next white keys are D, E, F, G, A, B, and the eighth key is C again, one octave higher. Newlands saw elements in the same way. For example, starting with Lithium (Li), the eighth element was Sodium (Na), which indeed shares very similar, highly reactive metallic properties with lithium. The next, Beryllium (Be), found its “musical match” in the ninth position, Magnesium (Mg).

Why Was the Law of Octaves Discarded?

Despite its initial cleverness, Newland’s Law was ultimately rejected by the scientific community for several critical reasons:

1. The “Calcium Wall”: It Failed with Heavier Elements
The law worked reasonably well for the lightest elements, like lithium, beryllium, and boron. However, as the list progressed into heavier elements like calcium, the pattern broke down completely. Elements that were supposed to be “similar” according to the octave rule had absolutely nothing in common chemically. It was as if the musical scale suddenly started producing random, discordant noises after a certain point.

2. Forced Groupings and Ignored Gaps
To make his pattern fit, Newlands was forced to place two elements in the same slot, much like putting two different notes on one piano key. For instance, he placed cobalt and nickel together in the same position, and did the same for platinum and iridium. This was an admission that the pure “every eighth element” rule didn’t hold. Furthermore, he left no gaps for elements that had not yet been discovered, making the model rigid and inflexible.

3. The Problem of New Elements
The 19th century was a period of rapid discovery in chemistry. As new elements were found, they simply could not be accommodated into Newlands’ rigid octave structure without destroying the entire pattern. The law had no predictive power or room for growth.

4. A Flawed Foundation: Atomic Weight
While Newlands was correct to use a quantitative measure (atomic weight), it is not the fundamental property that governs periodicity. The true key is the atomic number (the number of protons), a concept not understood at the time. Since atomic weight doesn’t always increase uniformly with atomic number (due to isotopes), the octave pattern was built on a shaky foundation.

5. Mockery and Skepticism
When Newlands presented his law to the Chemical Society in London, it was met with ridicule. Some scientists sarcastically asked if he had tried arranging the elements alphabetically instead. This mockery, while harsh, reflected the scientific consensus that the model was more of a curious coincidence for a few elements than a true law of nature.

Question 5. 

Did Dobereiner’s triads also exist in the columns of Newland’s Octaves? Compare and find out.

Ans:

Yes, there was a partial overlap between Dobereiner’s Triads and the columns of Newland’s Octaves, but it was not a perfect match.

Newland’s Law of Octaves arranged elements in increasing order of atomic mass, where every eighth element had properties similar to the first, just like notes in music. Some of Dobereiner’s triads, such as the halogen triad (chlorine, bromine, iodine), did fall within the same column of Newland’s Octaves. This happened because both systems were based on the fundamental idea that elements with similar properties could be grouped together.

However, the fit was not consistent. Newland’s table became disordered with heavier elements and included transition metals, breaking the pattern. Many of Dobereiner’s other triads, like the alkaline earth metal triad (calcium, strontium, barium), did not align perfectly in a single Newland’s column because other elements would fall in between, disrupting the triad relationship. Therefore, while a few triads existed in the octaves, the two schemes were different and neither could systematically organize all known elements.

Question 6. 

1. Lithium, sodium, and potassium elements were put in one group on the basis of their similar properties. What are those similar properties? 

2. The elements calcium, strontium and barium were put in one group or family on the basis of their similar properties. What were those similar properties?

Ans:

1. Similar Properties of Lithium, Sodium, and Potassium

Lithium, sodium, and potassium are grouped together as the Alkali Metals. Their similarities stem from all of them having just a single electron in their outermost shell, which they readily lose. This shared electron configuration leads to the following key similar properties:

• Intense Reactivity: They are famously reactive, especially with water, producing hydrogen gas and a strong alkaline solution (hence the group name). The reactivity increases dramatically as you move down the group from lithium to potassium.
• Soft Texture: They are all soft enough to be cut with a simple knife. Lithium is the hardest of the three, but it is still soft by metal standards, while sodium and potassium are even softer.
• Low Density: They are very light metals. Lithium, sodium, and potassium are all less dense than water, which is why they float and fizz violently during their water reaction.
• Silvery, Shiny Appearance: When freshly cut, they exhibit a characteristic metallic, silvery-white luster. However, they tarnish almost instantly upon exposure to air, forming a dull oxide or nitride layer.
• Forming Strong Alkalis: They all react with water to form strong bases (alkalis), such as sodium hydroxide (NaOH) and potassium hydroxide (KOH).
• Forming Similar Compounds: They each form compounds with a 1:1 ratio, such as chlorides (NaCl, KCl), carbonates (Na₂CO₃, K₂CO₃), and sulfates (Na₂SO₄, K₂SO₄), which often share similar chemical behaviors.
• Colored Flames: When heated in a flame, they each impart a characteristic color: lithium (crimson red), sodium (intense yellow), and potassium (lilac).

2. Similar Properties of Calcium, Strontium, and Barium

Calcium, strontium, and barium are grouped together as the Alkaline Earth Metals. Their similarity arises from having two electrons in their outermost shell. This shared characteristic results in the following common properties:

• High Reactivity: While not as violently reactive as the alkali metals, they are still very reactive and are never found in their pure, elemental form in nature. They readily lose their two outer electrons to form stable compounds.
• Formation of Alkaline Solutions: They react with water, though less vigorously than their alkali metal counterparts. For instance, calcium reacts with water to produce calcium hydroxide, a moderately strong alkali, and hydrogen gas.
• Silvery-White Metallic Luster: In their pure form, they are shiny, silvery-white metals. Like the alkali metals, they tarnish quickly in air, forming an oxide layer on their surface.
• Harder and Stronger: They are harder, denser, and have higher melting points than the alkali metals in the same period. For example, calcium is significantly harder than potassium.
• Forming Insoluble Compounds: A key characteristic is their tendency to form compounds that are often insoluble in water, especially their sulfates and carbonates. This is a classic test used to identify them.
• Colored Flames: They also produce distinctive flame colors, which are used for their identification: calcium (brick-red), strontium (crimson red), and barium (apple green).

Question 7. 

1.What was Mendeleev’s basis for the classification of elements? 

2. Mendeleev’s contributions to the concept of a periodic table laid the foundation for the Modern Periodic Table. Give reasons.

Ans:

1. What was Mendeleev’s basis for the classification of elements?

Mendeleev’s primary basis for classifying elements was their atomic mass. He observed that when elements were arranged in order of increasing atomic mass, their chemical and physical properties showed a repeating, or periodic, pattern. He grouped elements with similar properties together in the same vertical columns. Crucially, he prioritized this periodic pattern over strict atomic mass order in a few cases, swapping elements to maintain family resemblance, and he confidently left gaps for elements he predicted were yet to be discovered.

2. Mendeleev’s contributions laid the foundation for the Modern Periodic Table. Give reasons.

Mendeleev’s work was foundational for two key reasons. First, his successful predictions validated his periodic law; he left gaps for unknown elements like ‘eka-aluminium’ (Gallium) and ‘eka-silicon’ (Germanium) and accurately predicted their properties, which were later confirmed. Second, his table provided a powerful framework that highlighted relationships between elements, allowing chemists to systematically study and correct inaccuracies in atomic masses and properties, directly paving the way for the modern table based on atomic number.

Question 8. 

State Mendeleev’s periodic law.

Ans:

Mendeleev’s groundbreaking contribution was the formulation of the Periodic Law, which established that the properties of elements are not random but follow a repeating pattern when arranged by their atomic weights. He demonstrated that elements with similar chemical behaviors reappear at regular intervals, or periods, creating a natural rhythm in the layout of the elements. This principle allowed him to construct a table where families of elements, like the alkali metals and halogens, were grouped together vertically, revealing a profound order within the chemical world.

The most compelling validation of Mendeleev’s table was its predictive power. With remarkable confidence, he left deliberate gaps in his arrangement for elements that were unknown during his time. For these missing elements, such as gallium and germanium, he precisely forecasted their physical and chemical properties based on the trends observed in their surrounding elements. The subsequent discovery of these elements, with characteristics matching his predictions almost exactly, cemented the credibility and utility of his periodic system.

It is crucial to recognize that the modern Periodic Law has refined this concept by replacing atomic weight with atomic number as the fundamental organizing principle. This shift, which came with a deeper understanding of atomic structure, resolved inconsistencies in Mendeleev’s original table. However, his core insight—that elemental properties are periodic—remains utterly sound, and his original model stands as a testament to brilliant scientific reasoning that correctly deciphered nature’s patterns based on the knowledge available at the time.

Question 9. 

1. Use Mendeleev’s Periodic Table to predict the formula of hydrides of carbon and silicon. 

2. Use Mendeleev’s Periodic Table to predict the formula of oxides of potassium, aluminum, and barium.

Ans:

1. Predicting the Formula of Hydrides of Carbon and Silicon

In Mendeleev’s Periodic Table, elements are arranged in groups based on their valence, which is their combining power with other atoms. Carbon and silicon are both located in Group IV. This means they have a valence of 4 towards hydrogen.

• Carbon (Group IV): The formula of its hydride is CH₄ (Methane).
• Silicon (Group IV): The formula of its hydride is SiH₄ (Silane).

Conclusion: Using Mendeleev’s Table, the predicted hydrides for Group IV elements follow the formula XH₄.

2. Predicting the Formula of Oxides of Potassium, Aluminum, and Barium

Mendeleev’s Table allows us to predict the formula of oxides by using the valency of the element with oxygen, which has a valency of 2.

• Potassium (K): Located in Group I, it has a valence of 1. To combine with oxygen (valence 2), two potassium atoms are needed. The formula of its oxide is K₂O.
• Aluminum (Al): Located in Group III, it has a valence of 3. To combine with oxygen (valence 2), the simplest ratio is two aluminum atoms (2 x 3 = 6) for three oxygen atoms (3 x 2 = 6). The formula of its oxide is Al₂O₃.
• Barium (Ba): Located in Group II, it has a valence of 2. This matches perfectly with oxygen’s valence of 2. The formula of its oxide is BaO.

Question 10. 

1.  Which group of elements was missing from Mendeleev’s original periodic table? 

2. How can the reaction proceed continuously? 

Ans:

1. Which group of elements was missing from Mendeleev’s original periodic table?

The most notable group entirely absent from Mendeleev’s original periodic table was the noble gases (Group 18 in the modern table).

The reason for their omission is simple: they had not yet been discovered. Mendeleev published his periodic law in 1869. The first noble gas, helium, was identified in the spectrum of the sun in 1868 but was not yet recognized as a new element on Earth. The first earthly discovery of a noble gas was argon in 1894 by Lord Rayleigh and William Ramsay. This discovery was met with skepticism because these elements were chemically inert and did not fit into the existing table’s pattern.

The brilliance of the periodic table’s structure, however, allowed for this expansion. Rather than disrupting his model, Ramsay correctly proposed a new, entire column for these inert elements between the highly reactive halogens and alkali metals. This addition strengthened the predictive power of the periodic table rather than breaking it.

2. How can the reaction proceed continuously?

For a chemical reaction to proceed continuously, it cannot be a simple reaction in a closed container where reactants are used up and products accumulate, eventually stopping the process. Continuous operation requires a specific setup that constantly refreshes the reaction mixture. This is achieved primarily in two ways:

1. Using an Open System with Continuous Flow:
This is the most common method in industrial chemistry. Imagine a pipeline or a reactor vessel where the reactant substances are fed in at one end at a steady rate. Inside the reactor, the chemical transformation takes place. Simultaneously, the product mixture is continuously removed from the other end. This creates a dynamic equilibrium where the reaction never runs out of reactants or gets choked by its own products. Examples include the Haber process for ammonia synthesis and the catalytic cracking of oil in a fluidized bed reactor.

2. Employing a Regenerative Catalyst:
In some complex reactions, the catalyst itself can be temporarily consumed. For the reaction to be continuous, the system must include a parallel process that constantly regenerates the active catalyst. A perfect example is the catalytic converter in a car. It uses platinum and rhodium to convert harmful exhaust gases. While these metals facilitate the reaction, they can become “poisoned” or coated. The system is designed so that the constant flow of exhaust and the high temperature, along with secondary reactions, continuously clean and regenerate the catalytic surfaces, allowing the treatment of exhaust to proceed for the life of the vehicle.

Question 11. 

State the merits of Mendeleev’s classification of elements.

Ans:

Dmitri Mendeleev’s formulation of the Periodic Table in 1869 was a revolutionary leap in chemistry. Its profound success was not just due to its organization of known elements, but also its predictive and adaptable nature. The key merits of his classification are as follows:

1. Strategic Gaps for the Unknown
Unlike previous attempts, Mendeleev had the foresight and confidence to leave blank spaces in his table. He did not force known elements to fit a perfect pattern. Instead, he correctly deduced that these gaps represented elements that had not yet been discovered. This transformed the table from a static catalog into a dynamic, predictive map, guiding the quest for new elements.

2. A Foundation for Predicting Properties
Directly linked to the first merit, Mendeleev could accurately predict the physical and chemical properties of these missing elements. For instance, he described the characteristics of “Eka-aluminium” (later named Gallium), “Eka-boron” (Scandium), and “Eka-silicon” (Germanium) with remarkable accuracy. The subsequent discovery of these elements, matching his predictions, was the strongest possible validation of his system.

3. A System that Accommodated Corrections
At the time, atomic mass was the only measurable guide for ordering elements. Mendeleev wisely prioritized the periodicity of chemical properties over a strict, slavish adherence to increasing atomic mass. In a few specific cases (like tellurium and iodine, or cobalt and nickel), he deliberately swapped the order based on their chemical behavior, a bold move that was later vindicated when the concept of atomic number was introduced.

4. Grouping by Chemical Kinship
The vertical columns (groups) in Mendeleev’s table brought together elements with strikingly similar properties. For example, the halogens (fluorine, chlorine, bromine, iodine) were all grouped together, clearly displaying their shared traits like high reactivity and tendency to form salts with metals. This systematic grouping allowed chemists to study families of elements rather than just individual entities.

5. A Coherent Framework for a Growing Field
Before Mendeleev, the study of elements was chaotic, with no clear relationship between the 60+ known elements. His table provided a logical and systematic framework that organized all known elements based on a fundamental principle: periodicity. It brought order to the chemical world, making the vast body of chemical knowledge easier to understand, teach, and build upon.

Question 12. 

Why did Mendeleev leave some gaps in his periodic table of elements? Explain your answer with an example.

Ans:

Mendeleev left deliberate gaps in his periodic table because of his firm belief in the periodic law he had discovered. When he arranged the known elements in order of increasing atomic mass, he found that their chemical properties showed a repeating, or periodic, pattern. To maintain this consistent pattern, he had to place some elements in columns (groups) with others that shared similar properties, even if it meant skipping a spot. He was confident that these gaps did not represent flaws in his table, but rather elements that had not yet been discovered. His bold decision to predict the existence and properties of these missing elements was what truly set his work apart from other scientists of his time.

A famous example of this is the gap he left for an element he called “eka-aluminium” (“eka” meaning one in Sanskrit, indicating it was one place below aluminium in the group). Based on the pattern of the group, Mendeleev predicted the properties of this unknown element with remarkable accuracy. He forecasted that eka-aluminium would have an atomic mass of around 68, a low melting point, and a density of about 5.9 g/cm³. He also predicted its oxide formula would be Ea₂O₃ and that it would be discovered using spectroscopy.

In 1875, the French chemist Paul-Émile Lecoq de Boisbaudran discovered a new element, which he named Gallium. The properties of Gallium matched Mendeleev’s predictions for eka-aluminium almost exactly: its atomic mass was found to be 69.7, it had a low melting point (it melts in your hand), and its density was 5.91 g/cm³. This stunning confirmation was a major triumph for Mendeleev’s periodic table and proved that it was not just a cataloguing system, but a powerful tool that could predict the very building blocks of nature.

Question 13. 

The atomic number of an element is more important to the chemist than its relative atomic mass. Why?

Ans:

The atomic number is fundamentally more important to a chemist than the relative atomic mass because it defines the very identity of an element. The atomic number, representing the number of protons in an atom’s nucleus, is an element’s unique and unchangeable fingerprint. This number alone determines an element’s position on the Periodic Table and dictates the number of electrons in a neutral atom. Since the electrons, and particularly the arrangement of electrons in the outermost shell, are responsible for all chemical bonding and reactions, the atomic number ultimately governs an element’s chemical personality and behavior.

In contrast, the relative atomic mass, which is the weighted average mass of all an element’s naturally occurring isotopes, does not define identity. This is perfectly illustrated by the historical problem of the “pair reversal” in elements like argon (atomic number 18) and potassium (atomic number 19). If arranged by atomic mass alone, argon would come before potassium, which disrupts the periodic pattern of chemical properties. However, when arranged by atomic number, potassium correctly comes before argon, perfectly aligning with their chemical families—potassium as a reactive alkali metal and argon as an inert noble gas. This proves that the atomic number, not the mass, is the true key to the Periodic Law.

Therefore, while relative atomic mass is crucial for quantitative calculations like stoichiometry, it is secondary in importance. The atomic number is the master key for a chemist. It allows for the prediction of how an element will react, what kinds of bonds it will form, and its general reactivity, simply by its position on the table. In essence, the atomic number tells you what the element is and how it will behave, while the atomic mass tells you about a quantitative characteristic of a sample of that element.

Question 14. 

Consider the following elements: Be, Li, Na, Ca, K. Name the elements of (a) same group (b) same period.

Ans:

Considering the elements Beryllium (Be), Lithium (Li), Sodium (Na), Calcium (Ca), and Potassium (K):

(a) Elements of the same group:
Elements in the same group (vertical column) have the same number of valence electrons.

• Group 1 (Alkali Metals): Lithium (Li), Sodium (Na), and Potassium (K).
• Group 2 (Alkaline Earth Metals): Beryllium (Be) and Calcium (Ca).

(b) Elements of the same period:
Elements in the same period (horizontal row) have the same number of electron shells.

• Period 2: Lithium (Li) and Beryllium (Be).
• Period 4: Potassium (K) and Calcium (Ca).

Sodium (Na) is in Period 3 by itself from this list.

Question 15. 

1. Name an element whose properties were predicted on the basis of its position in Mendeleev’s periodic table. 

2. Name two elements whose atomic weights were corrected on the basis of their positions in Mendeleev’s periodic table. 

3. How many elements were known at the time of Mendeleev’s classification of elements?

Ans:

1. The element Germanium is a prime example whose properties were predicted by Mendeleev. He called it “Eka-silicon” and accurately forecasted its atomic weight, density, and the properties of its oxide and chloride, years before the element was actually discovered.
2. Mendeleev corrected the atomic weights of two elements based on their positions:
   • Beryllium: Its atomic weight was incorrectly thought to be 13.5, but Mendeleev, considering its properties similar to magnesium, proposed it should be 9.4, which is close to the modern value of 9.012.
   • Indium: Its atomic weight was believed to be 75.6, but to fit its properties correctly in the table, Mendeleev suggested it should be approximately 113, which was later confirmed.
3. At the time of Mendeleev’s original classification of elements in 1869, approximately 63 elements were known to scientists.

Exercise 5 (B)

Question 1. 

1. State the modern periodic law. 

2. How many periods and groups are there in the modern periodic table?

Ans:

1. Modern Periodic Law 

The Modern Periodic Law states that:

The physical and chemical properties of the elements are a periodic function of their atomic numbers.

This law means that when elements are arranged in increasing order of their atomic number (Z), elements with similar properties recur at regular intervals (periods). This resolved the inconsistencies found in Mendeleev’s original table, which was based on atomic mass.

2. Periods and Groups in the Modern Periodic Table 

The Modern Periodic Table is organized into the following number of periods and groups:

• Periods: There are 7 periods (the horizontal rows).
• Groups: There are 18 groups (the vertical columns).

Question 2. 

What is the main characteristic of the last elements in the periods of a periodic table? What is the general name of such elements?

Ans:

The main characteristic of the last elements in the periods of the periodic table is their stable and unreactive nature. This is a direct result of their electron configuration, as their outermost electron shells are completely filled. For most of these elements, this means having a full octet (eight electrons) in their valence shell, making them exceptionally stable and satisfying the octet rule. This complete electron shell makes them highly non-reactive because they have a very low tendency to gain, lose, or share electrons with other atoms.

The general name given to this distinct family of elements is the Noble Gases. This group includes elements like Helium, Neon, Argon, Krypton, Xenon, and Radon, which are all located in Group 18 of the modern periodic table. Their lack of chemical reactivity under standard conditions is their most defining property, which is why they are often used in applications where inertness is required, such as in protective atmospheres, lighting (neon signs), and as a non-reactive shield in welding.

Question 3. 

1. What is meant in the periodic table by a group? 

2. What is meant in the periodic table by a period?

Ans:

1. What is meant in the periodic table by a group?

There are 18 such numbered columns. Elements residing in the same group share a key characteristic: they have the same number of electrons in their outermost shell, known as valence electrons. This identical valence electron configuration is the primary reason why elements within a group exhibit very similar chemical properties and reactivity. For example, all elements in Group 1 (the alkali metals) have one valence electron, which makes them highly reactive and likely to form ions with a +1 charge. Think of a group as a chemical “family” where the members behave in a similar manner.

2. What is meant in the periodic table by a period?

There are 7 periods, corresponding to the number of electron shells an atom possesses. The valence electrons are being added to the same outer shell, but the increasing nuclear pull makes the atomic radius decrease. This gradual filling of the same electron shell results in a steady and predictable change in properties, from metallic on the left to non-metallic on the right. In essence, a period represents a complete cycle of this progressive change in elemental character.

Question 4. 

From the standpoint of atomic structure, what determines which elements will be the first and which the last in a period of the periodic table?

Ans:

From the standpoint of atomic structure, the position of an element within a period is determined by its atomic number (Z) and, consequently, its electronic configuration.

1. First Element in a Period (Group 1) 

The first element in any period (excluding Period 1) will be an Alkali Metal (Group 1).

• Valence Electrons: It is defined by having only one valence electron in its outermost shell.
• Atomic Number: Its atomic number (Z) is such that its electrons just fill a principal quantum shell and begin filling the next new shell.
• Property: It is highly electropositive (metallic) and easily loses this single valence electron to form a unipositive ion (M^+).

2. Last Element in a Period (Group 18) 

The last element in any period is a Noble Gas (Group 18).

• Valence Electrons: It is defined by having a complete outermost shell (a stable octet, or duplet for Helium).
• Atomic Number: Its atomic number (Z) is such that all electron shells are completely filled for that principal quantum number.
• Property: It is highly unreactive due to its stable electron configuration.

Summary

The start of a new period is defined by the addition of the first electron into a new, higher principal energy shell (n). The end of a period is defined when that shell is completely filled. The atomic number is simply the counter that tracks this sequential filling of electron shells.

Question 5. 

1. What are the following groups known as? Group 1 

2. What are the following groups known as? Group 17 

3. What are the following groups known as? Group 18  

4. Name two elements of each group.

Ans:

1. Group 1 is known as the Alkali Metals.
   • Two elements from this group are Lithium and Sodium.
2. Group 17 is known as the Halogens.
   • Two elements from this group are Chlorine and Iodine.
3. Group 18 is known as the Noble Gases.
   • Two elements from this group are Neon and Argon.

Question 6. 

1. What is the number of elements in the 1st period? 

2. What is the number of elements in the 3rd period of the modern periodic table?

Ans:

The first period of the periodic table is remarkably brief, consisting solely of hydrogen and helium. This minimalistic structure is a direct reflection of the rules governing atomic architecture. The first and innermost electron shell has a very limited capacity, able to accommodate only two electrons to achieve a stable, filled state. Hydrogen, with its single electron, begins the table, while helium perfectly completes this foundational shell, bringing the first period to an immediate and logical conclusion.

In stark contrast, the third period unfolds with a much broader sequence of eight elements. This journey across the period starts with the soft, reactive metal sodium and progresses through a fascinating transition. We encounter other metals like magnesium and aluminium, then move into the metalloid silicon, before arriving at clearly non-metallic elements such as phosphorus, sulfur, and chlorine. The period culminates with the noble gas argon, an element known for its notable inertness and stability. This entire progression, from a highly reactive metal to an unreactive gas, demonstrates the predictable and gradual change in chemical character across a period, driven by the step-by-step filling of the third electron shell which can hold a maximum of eight electrons.

Question 7. 

1.How does the number of valence electrons vary on moving from left to right in the third period of a periodic table? 

2. How does the number of (i) valence electrons (ii) valency; vary on moving from left to right in the second period of a periodic table?

Ans:

1. Variation of Valence Electrons in the Third Period

The third period begins with Sodium (Na) and ends with Argon (Ar). As we move from left to right across this period, the number of valence electrons increases regularly from 1 to 8.

This happens because the atomic number increases by one with each element, meaning one extra proton and one extra electron are added. These new electrons are added to the same outermost shell (the third shell, or ‘M’ shell). The filling order dictates that this shell gets populated before inner shells are filled further.

Here is a breakdown for clarity:

• Sodium (Na), Atomic No. 11: Electron configuration ends with 3s¹. Valence electrons = 1
• Magnesium (Mg), Atomic No. 12: Electron configuration ends with 3s². Valence electrons = 2
• Aluminium (Al), Atomic No. 13: Electron configuration ends with 3s² 3p¹. Valence electrons = 3
• Silicon (Si), Atomic No. 14: Electron configuration ends with 3s² 3p². Valence electrons = 4
• Phosphorus (P), Atomic No. 15: Electron configuration ends with 3s² 3p³. Valence electrons = 5
• Sulphur (S), Atomic No. 16: Electron configuration ends with 3s² 3p⁴. Valence electrons = 6
• Chlorine (Cl), Atomic No. 17: Electron configuration ends with 3s² 3p⁵. Valence electrons = 7
• Argon (Ar), Atomic No. 18: Electron configuration ends with 3s² 3p⁶. Valence electrons = 8

In summary: The number of valence electrons increases sequentially from 1 to 8, which is why the chemical properties of elements change dramatically across a period.

2. Variation in the Second Period

The second period begins with Lithium (Li) and ends with Neon (Ne).

(i) Variation in the Number of Valence Electrons

Similar to the third period, the number of valence electrons increases from 1 to 8 on moving from left to right in the second period.

The electrons are being added to the second shell (the ‘L’ shell), which is the outermost shell for these elements.

• Lithium (Li), Atomic No. 3: Electron configuration ends with 2s¹. Valence electrons = 1
• Beryllium (Be), Atomic No. 4: Electron configuration ends with 2s². Valence electrons = 2
• Boron (B), Atomic No. 5: Electron configuration ends with 2s² 2p¹. Valence electrons = 3
• Carbon (C), Atomic No. 6: Electron configuration ends with 2s² 2p². Valence electrons = 4
• Nitrogen (N), Atomic No. 7: Electron configuration ends with 2s² 2p³. Valence electrons = 5
• Oxygen (O), Atomic No. 8: Electron configuration ends with 2s² 2p⁴. Valence electrons = 6
• Fluorine (F), Atomic No. 9: Electron configuration ends with 2s² 2p⁵. Valence electrons = 7
• Neon (Ne), Atomic No. 10: Electron configuration ends with 2s² 2p⁶. Valence electrons = 8

(ii) Variation in Valency

Valency is the combining capacity of an atom. Unlike the steady increase in valence electrons, valency first increases and then decreases.

This pattern arises because valency depends on whether an atom tends to lose electrons (metals) or gain electrons (non-metals) to achieve a stable octet.

• Lithium (Li): Has 1 valence electron. It loses 1 electron to form Li⁺. Valency = 1
• Beryllium (Be): Has 2 valence electrons. It loses 2 electrons to form Be²⁺. Valency = 2
• Boron (B): Has 3 valence electrons. It can lose 3 electrons (or sometimes share), showing a Valency = 3
• Carbon (C): Has 4 valence electrons. It typically shares 4 electrons, showing a Valency = 4
• Nitrogen (N): Has 5 valence electrons. It needs 3 more to complete its octet, so it can gain 3 electrons (e.g., in N³⁻) or share three pairs. Valency = 3
• Oxygen (O): Has 6 valence electrons. It needs 2 more, so it gains 2 electrons (e.g., in O²⁻) or shares two pairs. Valency = 2
• Fluorine (F): Has 7 valence electrons. It needs 1 more, so it gains 1 electron (e.g., in F⁻). Valency = 1
• Neon (Ne): Has a complete octet of 8 valence electrons. It is inert and does not readily form compounds. Valency = 0

Question 8. 

How do atomic structures (electron arrangements) change in a period with an increase in atomic numbers moving left to right?

Ans:

When moving from left to right across a period in the Periodic Table, the most significant change in the atomic structure is the progressive addition of protons to the nucleus and electrons to the same outermost shell, or valence shell. While a new period begins when a new electron shell starts to fill, throughout a single period, the principal energy level (like n=1, n=2, etc.) remains constant. For instance, in the third period, from Sodium (Na) to Argon (Ar), the electrons are all being added to the third shell (the M-shell). What changes is the number of electrons in this valence shell, starting from one in sodium and increasing to a full complement of eight in argon.

This consistent increase in atomic number, and thus the number of protons in the nucleus, has a profound effect. The positive charge of the nucleus increases with each successive element. Since the additional electrons are entering the same principal energy level, they do not effectively shield each other from this growing nuclear charge. Consequently, the entire electron cloud is pulled progressively closer to the nucleus. This leads to a steady decrease in atomic radius across the period. The atom becomes smaller and more tightly held as you move from left to right.

This changing electron arrangement directly dictates the chemical properties of the elements. On the left side, elements have only one or two electrons in their valence shell. These are held weakly by the weaker net nuclear charge and are easily lost, explaining the strong metallic character of these elements. As we move right, the increasing nuclear charge holds the growing number of valence electrons more and more tightly. Elements in the middle have a balanced ability to share electrons (covalent bonding), while those on the far right, with nearly full valence shells, have a strong tendency to gain electrons to achieve a stable configuration, exhibiting non-metallic character. Thus, the systematic left-to-right change in electron configuration is the fundamental reason behind the observed periodic trends in atomic size and chemical behavior.

Question 9. 

1. This question refers to elements of the periodic table with atomic numbers from 3 to 18. In the table below, some elements are shown by letters, even though the letters are not the usual symbols of the elements.  

3	4	5	6	7	8	9	10
A	B	C	D	E	F	G	H
11	12	13	14	15	16	17	18
I	J	K	L	M	N	O	P

Which of these is:

a noble gas?

a halogen?

an alkali metal?

an element with valency 4? 

2. This question refers to elements of the periodic table with atomic numbers from 3 to 18. In the table below, some elements are shown by letters, even though the letters are not the usual symbols of the elements. 

3	4	5	6	7	8	9	10
A	B	C	D	E	F	G	H
11	12	13	14	15	16	17	18
I	J	K	L	M	N	O	P

If A combines with F, what would be the formula of the resulting compound?

3. This question refers to elements of the periodic table with atomic numbers from 3 to 18. In the table below, some elements are shown by letters, even though the letters are not the usual symbols of the elements.

3	4	5	6	7	8	9	10
A	B	C	D	E	F	G	H
11	12	13	14	15	16	17	18
I	J	K	L	M	N	O	P

What is the electronic configuration of G?

Ans:

Step 1: Identify the Elements from Atomic Numbers

The atomic numbers 3 to 18 correspond to the following elements:

• 3: Lithium (Li) → A
• 4: Beryllium (Be) → B
• 5: Boron (B) → C
• 6: Carbon (C) → D
• 7: Nitrogen (N) → E
• 8: Oxygen (O) → F
• 9: Fluorine (F) → G
• 10: Neon (Ne) → H
• 11: Sodium (Na) → I
• 12: Magnesium (Mg) → J
• 13: Aluminium (Al) → K
• 14: Silicon (Si) → L
• 15: Phosphorus (P) → M
• 16: Sulfur (S) → N
• 17: Chlorine (Cl) → O
• 18: Argon (Ar) → P

Step 2: Answer Question 1

• (a) Noble Gas: Noble gases are in Group 18 (full valence shell).
  → H (Neon) and P (Argon).
• (b) Halogen: Halogens are in Group 17 (7 valence electrons).
  → G (Fluorine) and O (Chlorine).
• (c) Alkali Metal: Alkali metals are in Group 1 (1 valence electron).
  → A (Lithium) and I (Sodium).
• (d) Element with Valency 4: Elements in Group 14 have 4 valence electrons.
  → D (Carbon) and L (Silicon).

Step 3: Answer Question 2

• A (Lithium) has valency 1.
• F (Oxygen) has valency 2.
• To balance charges, two lithium atoms combine with one oxygen atom.
• Formula: A₂F (which corresponds to Li₂O).

Step 4: Answer Question 3

• G is Fluorine (atomic number 9).
• Electronic configuration: 2,7 (i.e., 2 electrons in the first shell, 7 in the second).

Final Answers

1. (a) Noble Gas: H, P
   (b) Halogen: G, O
   (c) Alkali Metal: A, I
   (d) Valency 4: D, L
2. Formula when A combines with F: A₂F
3. Electronic configuration of G: 2,7

Question 10. 

Sodium and aluminum have atomic numbers 11 and 13, respectively. They are separated by one element in the periodic table and have valencies 1 and 3 respectively. Chlorine and potassium are also separated by one element in the periodic table (their atomic numbers being 17 and 19, respectively) and yet both have valency 1. Explain.

Ans:

The difference in valency behavior between the two pairs of elements you mentioned stems from their distinct positions within the Periodic Table, specifically whether they are located in the same period or the same group. Let’s first consider sodium (atomic number 11) and aluminum (atomic number 13). These elements reside in the same third period. As we move from left to right across a period, the number of valence electrons—the electrons available for bonding—increases sequentially. Sodium, being in Group 1, has a single electron in its outermost shell, leading to a valency of 1. The element between them, magnesium (atomic number 12, Group 2), has two valence electrons and a valency of 2. Finally, aluminum, in Group 13, possesses three valence electrons, which gives it a valency of 3. Their valencies change because they are filling the same electron shell, and their reactivity is defined by completing that shell.

In contrast, the case of chlorine (atomic number 17) and potassium (atomic number 19) is fundamentally different. They are not in the same period; chlorine is at the end of the third period (Group 17), while potassium is at the beginning of the fourth period (Group 1). The element separating them is argon (atomic number 18), a noble gas. The key principle here is that elements within the same group have an identical number of valence electrons. Chlorine, as a halogen in Group 17, has seven valence electrons. Its tendency is to gain one electron to achieve a stable, full outer shell, resulting in a valency of 1. Potassium, as an alkali metal in Group 1, has one valence electron. Its tendency is to lose that single electron to achieve its own stable configuration, also resulting in a valency of 1.

Therefore, while the atomic number gap is the same, the underlying reason for their valencies is not their proximity in atomic number, but their group identity. Sodium and aluminum are neighbors across different groups in a single period, explaining their different valencies. Chlorine and potassium, though separated by only one element in sequence, belong to the first and last groups of the table, respectively. Both groups have distinct electronic configurations that, for different reasons (gaining one electron versus losing one electron), lead to the same net combining power or valency of 1.

Question 11. 

Helium is an unreactive gas and neon is a gas of extremely low reactivity. What, if anything, do their atoms have in common.

Ans:

The atoms of helium and neon share a key structural feature that directly explains their lack of reactivity: a full and stable outer shell of electrons. This stable electron configuration is known as an octet for most elements, or a duet in the specific case of helium. For an atom, having this complete outer shell is the most energetically favorable state, meaning it has very little tendency to gain, lose, or share electrons with other atoms.

Helium, the first noble gas, has only one electron shell, which is completely filled with its two electrons. Neon, further along the same group, has two electron shells; its outermost shell is filled with a stable octet of eight electrons. Because their outer electron shells are already full, their atoms possess no “free” spaces to accept extra electrons and no easily lost electrons to donate. This inherent electronic stability is the fundamental commonality that makes both helium and neon so chemically unreactive and inert.

Question 12. 

1. In which part of a group would you separately expect the elements to have the greatest metallic character 

2. In which part of a group would you separately expect the elements to have the largest atomic size?

Ans:

1. Greatest Metallic Character

You would expect to find the elements with the greatest metallic character at the bottom of a group on the periodic table.

Reasoning: Metallic character is essentially how easily an atom can lose its outer electrons. As you move down a group, each new element has an extra full shell of electrons. These outer electrons are located farther from the attractive pull of the nucleus. Furthermore, the inner electron shells act as a “shield,” weakening the nucleus’s grip on the outermost electrons. This combination of increased distance and shielding makes it much easier for the atom to lose its outer electrons and exhibit strong metallic properties, such as conducting electricity and forming positive ions.

For example, in Group 14, the element at the bottom, Lead (Pb), is a recognizably soft and malleable metal, while the element at the top, Carbon (C), in the form of diamond, is a non-metal and one of the hardest substances known.

2. Largest Atomic Size

You would also expect to find the elements with the largest atomic size at the bottom of a group.

Reasoning: Atomic size is determined by the distance of the outermost electrons from the nucleus. As you move down a group, the number of occupied electron shells (energy levels) increases with each period. Each new shell is located farther from the nucleus. Even though the positive charge of the nucleus is also increasing, the dominant factor is the addition of these new, outer shells. This steady increase in the number of electron shells ensures that atoms get progressively larger as you move down any given group.

For instance, in Group 1 (the alkali metals), Francium (Fr) at the bottom has a much larger atomic radius than Lithium (Li) at the top.

Question 13. 

What happens to the number of valence electrons in atoms of elements as we go down a group of the periodic table?

Ans:

Question 14. 

1. The position of elements A, B, C, D and E in the periodic table are shown below:

Group 1	Group 2	Group 17	Group 18
			D
	B	C
A			E

State which is metals, non-metals, and noble gas in this table.

2. The position of elements A, B, C, D and E in the periodic table are shown below:

Group 1	Group 2	Group 17	Group 18
			D
	B	C
A			E

State which is most reactive (i) metal (ii) non-metal

3. The position of elements A, B, C, D and E in the periodic table are shown below:

Group 1	Group 2	Group 17	Group 18
			D
	B	C
A			E

Which type of ion will be formed by elements A, B, and C.

4. The position of elements A, B, C, D and E in the periodic table are shown below:

Group 1	Group 2	Group 17	Group 18
			D
	B	C
A			E

Which is larger in size (i) D or E (ii) B or C

Ans:

1. Classification of Elements 

• Metals: Elements in Groups 1 and 2 are metals.
  • D, B, E
• Non-metals: Elements in Group 17 are non-metals.
  • A
• Noble Gas: Elements in Group 18 are noble gases.
  • C

2. Most Reactive Elements 

• (i) Most Reactive Metal:
  • Reactivity of metals increases down a group.
  • The most reactive metal is E (it is lower than D in Group 1).
• (ii) Most Reactive Non-metal:
  • Reactivity of non-metals (Group 17) decreases down a group.
  • The most reactive non-metal is A (it is the highest element in Group 17 shown).

3. Type of Ion Formed 

• A (Group 17): Forms a negative ion (Anion, A^-) by gaining one electron to complete its octet.
• B (Group 2): Forms a positive ion (Cation, B^2+) by losing two valence electrons.
• C (Group 18): Forms no ion (it is a stable noble gas).

4. Comparison of Atomic Size 

(i) D or E

• E is larger. Atomic size increases as you move down a group because a new, larger electron shell is added.

(ii) B or C

• B is larger. Atomic size generally decreases across a period (from left to right) due to increasing nuclear charge pulling the valence electrons closer to the nucleus. (Note: Noble gas C often appears slightly larger due to different measurement methods, but based on the general periodic trend across Period 2, B is typically considered larger in terms of metallic/covalent radii).

Question 15. 

Write the electronic configuration of element 17T35.

1. What is the group number of T?
2. What is the period number of T?
3. How many valence electrons are there in an atom of T?
4. What is the valency of T?
5. Is it metal or non-metal?
6. State number of protons and neutrons in T.

Ans:

The symbol 17T35 represents an atom of the element with:

• Atomic Number (Z) = 17 (This defines the element).
• Mass Number (A) = 35 (This is the sum of protons and neutrons for this specific isotope).

The element with atomic number 17 is Chlorine (Cl). The “T” in the question is a placeholder for its symbol.

Here are the answers to your questions:

1. Electronic Configuration of 17T35

The electronic configuration is written by filling electrons into orbitals in the order of increasing energy (Aufbau principle). Chlorine has 17 electrons.

• Detailed Configuration: 1s² 2s² 2p⁶ 3s² 3p⁵
• Short-hand (Noble Gas) Configuration: [Ne] 3s² 3p⁵

2. Group Number of T

The element is in Group 17 (also known as the Halogens) of the modern periodic table. This is determined by the valence shell configuration (ns² np⁵).

3. Period Number of T

The element is in Period 3. This is determined by the highest principal quantum number (n=3) of its valence electrons.

4. Number of Valence Electrons in an atom of T

There are 7 valence electrons. Valence electrons are those in the outermost shell (the n=3 shell). The configuration 3s² 3p⁵ gives a total of 2 + 5 = 7 valence electrons.

5. Valency of T

The valency is 1. Halogens have 7 valence electrons and need one more electron to achieve a stable octet. Therefore, they commonly have a valency of 1.

6. Is it metal or non-metal?

It is a non-metal. All elements in Group 17 (Halogens) are typical non-metals.

7. Number of Protons and Neutrons in T

• Number of Protons = Atomic Number (Z) = 17
• Number of Neutrons = Mass Number (A) – Atomic Number (Z) = 35 – 17 = 18

Exercise 5 (C)

Question 1. 

Element P has atomic number 19. To which group and period, does P belong? Is it metal or non-metal? Why?

Ans:

Element P, with an atomic number of 19, belongs to Group 1 and Period 4 of the Periodic Table.

It is a metal.

The reason for this classification is based on its electronic configuration, which is 2, 8, 8, 1. With only one electron in its outermost shell, it has a strong tendency to lose this single electron to achieve a stable, noble gas configuration. This characteristic of readily losing electrons to form positive ions (cations) is a fundamental property of metals. Elements in Group 1, known as the Alkali Metals, all share this property and are highly reactive, typical of metallic behavior.

Question 2. 

1. An element belongs to the third period and Group IIIA (13) of the periodic table. State: the number of valence electrons, 

2. An element belongs to the third period and Group IIIA (13) of the periodic table. State: the valency, 

3. An element belongs to the third period and Group IIIA (13) of the periodic table. State: if it is a metal or non-metal? 

4. An element belongs to the third period and Group IIIA (13) of the periodic table. State: the name of the element.

Ans:

1. The number of valence electrons is three. This is determined by its position in Group 13, which means its atoms have three electrons in their outermost shell.
2. Its valency is 3. Since it has three valence electrons, it tends to lose them to achieve a stable electronic configuration, leading to a valency of three.
3. This element is definitively a metal. It displays the characteristic shiny appearance and excellent electrical conductivity common to metals.
4. The element is Aluminum. It is the only element that fits the description of being in the third period and Group 13 of the periodic table.

Question 3. 

1. Name and state the following with reference to the elements of the first three periods of the periodic table. Noble gas with duplet arrangement of electrons. 

2. Name and state the following with reference to the elements of the first three periods of the periodic table. Metalloid in Period 3. 

3. Name and state the following with reference to the elements of the first three periods of the periodic table. Valency of elements in Group 14 and 15. 

4. Name and state the following with reference to the elements of the first three periods of the periodic table. Noble gas having electronic configuration: 2, 8, 8. 

5. Name and state the following with reference to the elements of the first three periods of the periodic table. Group whose elements have zero valencies. 

6. Name and state the following with reference to the elements of the first three periods of the periodic table. A covalent compound formed by an element in period 2 and a halogen.

7. Name and state the following with reference to the elements of the first three periods of the periodic table. Non-metallic elements present in Period 3 of Groups 15 and 16. 

8. Name and state the following with reference to the elements of the first three periods of the periodic table. An electrovalent compound formed by an alkaline earth metal and a halogen. 

9. Name and state the following with reference to the elements of the first three periods of the periodic table. Bridge elements of Period 3 of Group 1,2 and 3. 

10. Name and state the following with reference to the elements of the first three periods of the periodic table. Alkali metal in period 3 dissolves in the water giving a strong alkali. 

11.  Name and state the following with reference to the elements of the first three periods of the periodic table. Typical elements of Groups 14 and 15. 

12. Name and state the following with reference to the elements of the first three periods of the periodic table. Alkaline earth metal in period 3.

Ans:

1. Noble gas with duplet arrangement of electrons.

• Name: Helium (He)
• State: It is the only noble gas present in the first period and has a complete duplet (2 electrons) in its K-shell, making it stable and chemically inert.

2. Metalloid in Period 3.

• Name: Silicon (Si)
• State: Located in Group 14, it exhibits properties intermediate between metals and non-metals, such as having a shiny appearance but being a poor conductor of electricity.

3. Valency of elements in Group 14 and 15.

• Name: The common valencies for these groups are 4 and 3, respectively.
• State: Elements in Group 14 have 4 valence electrons, leading to a valency of 4 (e.g., Carbon, Silicon). Elements in Group 15 have 5 valence electrons, so they typically gain 3 electrons to achieve a stable octet, resulting in a valency of 3 (e.g., Nitrogen, Phosphorus).

4. Noble gas having electronic configuration: 2, 8, 8.

• Name: Argon (Ar)
• State: It is the noble gas located in the third period. Its atom has three electron shells with a stable octet in the outermost shell, confirming the configuration 2, 8, 8.

5. Group whose elements have zero valencies.

• Name: Group 18 (The Noble Gases)
• State: The elements of this group possess a stable electronic configuration with a complete valence shell (duplet or octet). As a result, they have no tendency to gain, lose, or share electrons, leading to a valency of zero.

6. A covalent compound formed by an element in period 2 and a halogen.

• Name: Carbon Tetrachloride (CCl₄)
• State: This compound is formed when carbon (from Period 2, Group 14) shares its four valence electrons with four chlorine atoms (halogens), resulting in a covalent bond and a stable molecular structure.

7. Non-metallic elements present in Period 3 of Groups 15 and 16.

• Name: Phosphorus (P) in Group 15 and Sulfur (S) in Group 16.
• State: These are the non-metallic elements found in the third period. Phosphorus is known for its tendency to form covalent compounds, while sulfur is a multivalent non-metal.

8. An electrovalent compound formed by an alkaline earth metal and a halogen.

• Name: Magnesium Chloride (MgCl₂)
• State: This ionic compound is formed when magnesium (an alkaline earth metal from Period 3) donates two electrons to two chlorine atoms (halogens), resulting in the formation of oppositely charged ions that are held together by a strong electrostatic force of attraction.

9. Bridge elements of Period 3 of Group 1, 2 and 3.

• Name: Sodium (Na) from Group 1, Magnesium (Mg) from Group 2, and Aluminium (Al) from Group 3.
• State: These three elements are the first members of their respective groups in the third period. They act as a “bridge” by showing a clear gradation in properties from strongly metallic (Na) to less metallic (Al), linking the properties of the second and third periods.

10. Alkali metal in period 3 dissolves in water giving a strong alkali.

• Name: Sodium (Na)
• State: It is the alkali metal found in the third period. It reacts vigorously with water to form sodium hydroxide (NaOH), which is a strong alkali (base).

11. Typical elements of Groups 14 and 15.

• Name: Carbon (C) and Silicon (Si) for Group 14; Nitrogen (N) and Phosphorus (P) for Group 15.
• State: These elements are the first two members of their groups in the second and third periods. They are called ‘typical elements’ because they distinctly represent the characteristic properties of their respective groups, showing a clear gradation in properties without the complication of an inner transition series.

12. Alkaline earth metal in period 3.

• Name: Magnesium (Mg)
• State: It is the only alkaline earth metal (Group 2) located in the third period. It is a light, silvery-white metal that forms a basic oxide.

Question 4. 

Match column A with column B.

Column A	Column B
(a) Element short by 1 electron in octet	(i) Transition elements
(b) Highly reactive metals	(ii) Noble gases
(c) Non-reactive elements	(iii) Alkali metals
(d) Elements of groups 3 to 12	(iv) Alkaline earth metals
(e) Radioactive elements	(v) Halogens
(f) Elements with 2 electrons in the outermost orbit	(vi)Actinides

Ans:

Column A	Column B	Rationale
(a) Element short by 1 electron in octet	(v) Halogens	Halogens are in Group 17 and have 7 valence electrons, needing 1 more to complete their octet.
(b) Highly reactive metals	(iii) Alkali metals	Alkali metals (Group 1) have only 1 valence electron, which they easily lose, making them the most reactive group of metals.
(c) Non-reactive elements	(ii) Noble gases	Noble gases (Group 18) have a complete octet (or duplet for Helium), making them chemically stable and non-reactive.
(d) Elements of groups 3 to 12	(i) Transition elements	The d-block elements located in Groups 3 through 12 of the Periodic Table are collectively known as the Transition Elements.
(e) Radioactive elements	(vi) Actinides	Actinides (elements Z=89 to Z=103) are nearly all synthetic and radioactive; they are often listed separately below the main table.
(f) Elements with 2 electrons in the outermost orbit	(iv) Alkaline earth metals	Alkaline earth metals (Group 2) all have 2 valence electrons, which they easily lose to form a 2+ ion.

Question 5. 

Complete the table:

Atomic number	Element	Electronic configuration	Select an element of the same group
11	Sodium		(Ca/N/K)…
15	Phosphorus		(Ba/N/Rb)…
16	Sulphur		(F/Cl/O)…
9	Fluorine		(Ca/Cl/K)…

Ans:

Atomic Number (Z)	Element	Electronic Configuration	Select an Element of the Same Group
11	Sodium (Na)	2, 8, 2001	K (Potassium, Z=19, Group 1)
15	Phosphorus (P)	2, 8, 2005	N (Nitrogen, Z=7, Group 15)
16	Sulphur (S)	2, 8, 2006	O (Oxygen, Z=8, Group 16)
9	Fluorine (F)	2, 7	Cl (Chlorine, Z=17, Group 17)

Question 6. 

1. Relative atomic mass of a light element up to calcium is approximately ______ its atomic number. 

2. The horizontal rows in a periodic table are called _____. 

3. Going across a period left to right, atomic size _____. 

4. Moving left to right in the second period, the number of valence electrons _____. 

5. Moving down in the second group, number of valence electrons ______.

Ans:

1. Relative atomic mass of a light element up to calcium is approximately double its atomic number.
2. The horizontal rows in a periodic table are called periods.
3. Going across a period left to right, atomic size decreases.
4. Moving left to right in the second period, the number of valence electrons increases (from 1 to 8).
5. Moving down in the second group, the number of valence electrons remains the same (always 2).

Question 7. 

1. Name the alkali metals, How many electron(s) they have in their outermost orbit. 

2. Take any one alkali metal and write its reaction with (i)oxygen (ii)water (iii)acid.

Ans:

1. The alkali metals are Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs), and Francium (Fr). All these elements possess a single electron in their outermost shell. This solitary valence electron is located in an ‘s’ orbital, giving them a uniform configuration of ns¹. It is this single, loosely held electron that is primarily responsible for their high reactivity and strong metallic character, as they tend to lose it readily to achieve a stable noble gas configuration.

2. Let us take Sodium (Na) as our example and observe its vigorous reactions:

(i) Reaction with Oxygen: Sodium metal tarnishes rapidly when exposed to air, burning with a characteristic bright yellow flame to form sodium oxide and sodium peroxide.
The reaction is: 4Na + O₂ → 2Na₂O

(ii) Reaction with Water: Sodium reacts violently with cold water, skimming across the surface and producing hydrogen gas. The reaction is so exothermic that the heat generated often ignites the hydrogen, which may burn with an orange-yellow flame. It forms a strong alkaline solution of sodium hydroxide.
The reaction is: 2Na + 2H₂O → 2NaOH + H₂

(iii) Reaction with Acid: The reaction with an acid, such as hydrochloric acid, is even more violent and explosive than with water. It produces the corresponding salt and hydrogen gas.
The reaction is: 2Na + 2HCl → 2NaCl + H₂

Question 8. 

1. Name the method by which alkali metals can be extracted. 

2. What is the colour of the flame of sodium and potassium?

Ans:

1. The technique used to obtain alkali metals from their compounds is electrolysis of their molten salts, such as chlorides. This is necessary because alkali metals are the most reactive group of metals and cannot be extracted by common reducing agents like carbon.
2. When burned, sodium produces an intense, persistent yellow flame. Potassium, on the other hand, burns with a lilac or pale violet flame, which can sometimes be difficult to see without a blue cobalt glass filter to block the yellow light often present from sodium impurities.

Question 9. 

1. An element A has 2 electrons in its fourth shell. State: its atomic number 

2. An element A has 2 electrons in its fourth shell. State its electronic configuration. 

3. An element A has 2 electrons in its fourth shell. State: its valency 

4. An element A has 2 electrons in its fourth shell. State: position in the periodic table 5. An element A has 2 electrons in its fourth shell. State: is it a metal or non-metal 

6. An element A has 2 electrons in its fourth shell. State: is it an oxidising or reducing agent

Ans:

1. Its atomic number: 20
   (The element has a total of 20 electrons. The electronic configuration (detailed below) adds up to this number.)
2. Its electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
   (This is the standard notation, showing electrons filling orbitals in order of increasing energy. The “4s²” explicitly states the two electrons in the fourth shell.)
3. Its valency: 2
   (The element has two electrons in its outermost shell (the 4s orbital). It tends to lose these two electrons to achieve a stable, full outer shell configuration, giving it a valency of 2.)
4. Its position in the periodic table: Period 4, Group 2
   *(The highest shell number is 4, so it is in Period 4. Having two electrons in its outermost s-orbital places it in Group 2.)*
5. Is it a metal or non-metal: Metal
   (Elements in Group 2 of the periodic table are known as Alkaline Earth Metals, which are all metallic in nature.)
6. Is it an oxidising or reducing agent: Reducing Agent
   (As a metal, it readily loses electrons (it gets oxidized) during a chemical reaction. A species that loses electrons and causes another to be reduced is defined as a reducing agent.)

Question 10. 

1. Name the first three alkaline earth metals.

2. Write the reactions of the first three alkaline earth metals with dilute hydrochloric acid.

Ans:

The initial members of the alkaline earth metal family are Beryllium (Be), Magnesium (Mg), and Calcium (Ca). A revealing way to explore their properties is by observing how they interact with dilute hydrochloric acid. In each case, a single displacement reaction occurs, where the metal displaces hydrogen from the acid, resulting in a metal chloride salt and hydrogen gas. A clear trend in reactivity becomes apparent as we move down this group.

1. Beryllium (Be): The Slow Starter

Beryllium’s interaction with dilute hydrochloric acid is notably subdued. This is not due to a lack of reactivity, but because the metal is protected by a strong, adherent oxide film on its surface. This passivating layer must be breached before the reaction can proceed. Once this happens, the metal slowly dissolves to form a solution of beryllium chloride while releasing hydrogen gas at a leisurely pace.

• Descriptive Summary: Beryllium + Hydrochloric Acid → Beryllium Chloride + Hydrogen
• Balanced Chemical Equation: Be(s) + 2HCl(aq) → BeCl₂(aq) + H₂(g)

2. Magnesium (Mg): The Vigorous Participant

In stark contrast, magnesium reacts with immediate and obvious vigor. Upon contact with the acid, a rapid effervescence occurs as bubbles of hydrogen gas stream from the metal. The reaction is exothermic, releasing a noticeable amount of heat, and the magnesium ribbon or powder will completely dissolve, leaving behind a solution of magnesium chloride.

• Descriptive Summary: Magnesium + Hydrochloric Acid → Magnesium Chloride + Hydrogen
• Balanced Chemical Equation: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

3. Calcium (Ca): The Energetic Reactor

Calcium takes this reactivity to a more intense level. The reaction is rapid and can be described as violent. Hydrogen gas is produced in a vigorous, frothing effervescence. The heat generated is often substantial; in some cases, it can be enough to ignite the escaping hydrogen, which may burn with a characteristic reddish-orange flame due to calcium ions in the flame. The product of this energetic reaction is calcium chloride.

• Descriptive Summary: Calcium + Hydrochloric Acid → Calcium Chloride + Hydrogen
• Balanced Chemical Equation: Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g)

Question 11. 

1. Write the electronic configuration of the first two alkaline earth metals. 

2. How do alkaline earth metals occur in nature?

Ans:

1. Electronic Configuration of the First Two Alkaline Earth Metals

The first two elements in the alkaline earth metal group are Beryllium (Be) and Magnesium (Mg). Their electronic configurations are as follows:

• Beryllium (Atomic Number 4): The distribution of its 4 electrons is 2, 2. This can also be written in the shell notation as K=2, L=2.
• Magnesium (Atomic Number 12): The distribution of its 12 electrons is 2, 8, 2. In shell notation, this is represented as K=2, L=8, M=2.

A key characteristic of all alkaline earth metals is that they have two electrons in their outermost shell, which is why they are found in Group 2 of the periodic table.

2. Occurrence of Alkaline Earth Metals in Nature

Alkaline earth metals are highly reactive due to their tendency to lose two valence electrons to achieve a stable noble gas configuration. Because of this reactivity, they are never found in their pure, native metallic state in nature. Instead, they always occur in the form of stable compounds, making them mineral constituents within the Earth’s crust.

These elements are commonly found as silicate and carbonate minerals. For example, Magnesium is abundant in minerals like dolomite and magnesite, while Calcium is a major component of limestone, chalk, and marble (all forms of calcium carbonate). Other important sources include gypsum (calcium sulfate) and beryl (a beryllium silicate). Their widespread presence in mineral forms makes them commercially significant for various industrial and construction purposes. Radium, the last member of the group, is radioactive and is found in trace amounts within uranium ores.

Question 12. 

1. Give reason: Alkali metals are kept in an inert solvent.

2. Give reason: Alkali metals and halogens do not occur free in nature. 

3. Give reason: Alkali and alkaline earth metal compounds usually form electrovalent compounds. 

4. Give reason: Inert gases do not form compounds.

Ans:

1. Give reason: Alkali metals are kept in an inert solvent.
Alkali metals, such as sodium and potassium, are exceptionally reactive. They have a strong tendency to react vigorously with both oxygen and moisture present in the air. When exposed, they can quickly form oxides or hydroxides and may even catch fire. To prevent these dangerous and uncontrolled chemical reactions, they are stored submerged in an inert solvent like kerosene oil. This solvent acts as a protective barrier, effectively shielding the metals from coming into contact with air or water vapour, thereby ensuring their safe storage.

2. Give reason: Alkali metals and halogens do not occur free in nature.
This is a direct consequence of their high chemical reactivity. Alkali metals possess just one electron in their outermost shell, which they tend to lose easily. Conversely, halogens are just one electron short of a stable configuration, giving them a powerful tendency to gain an electron. Because of this intense drive to react, they are never found in their pure, elemental (or “free”) state in nature. Instead, they are always found combined with other elements, forming stable compounds like salts (e.g., sodium chloride for the alkali metal sodium and the halogen chlorine).

3. Give reason: Alkali and alkaline earth metal compounds usually form electrovalent compounds.
The formation of an electrovalent (or ionic) compound involves the complete transfer of electrons from a metal atom to a non-metal atom. Alkali and alkaline earth metals have very low ionization energies, meaning they can lose their outermost electrons with great ease to achieve a stable noble gas configuration. When they encounter non-metals (like oxygen or chlorine) that have a high affinity to gain electrons, this electron transfer occurs readily. The resulting positively charged metal ions and negatively charged non-metal ions are then held together by strong electrostatic forces of attraction, forming a stable electrovalent compound.

4. Give reason: Inert gases do not form compounds.
Inert gases, also known as noble gases, possess a highly stable electron configuration. Their outermost electron shell is completely filled, having the maximum number of electrons it can hold. This full valence shell gives them no tendency to lose, gain, or share electrons with other atoms. Since chemical bonding fundamentally relies on the rearrangement of electrons to achieve stability, and noble gases are already stable by nature, they generally do not participate in chemical reactions or form compounds under normal conditions.

Question 13. 

1. Arrange the following: Elements of group 1, in increasing order of reactivity 

2. Arrange the following: Elements of group 17, in decreasing order of reactivity 

3. Arrange the following: He, Na, Mg ( increasing order of melting point) 

4. Arrange the following: Chlorine, sodium, magnesium (increasing reducing character)

Ans:

1. Elements of group 1, in increasing order of reactivity
Lithium (Li) < Sodium (Na) < Potassium (K) < Rubidium (Rb) < Cesium (Cs)
Reason: Reactivity in Group 1 alkali metals increases down the group. This is because the single outer electron is further from the nucleus and more easily shielded as we go down, making it far easier to lose in chemical reactions.

2. Elements of group 17, in decreasing order of reactivity
Fluorine (F) > Chlorine (Cl) > Bromine (Br) > Iodine (I)
Reason: Reactivity in Group 17 halogens decreases down the group. This is because the atomic size increases, making it harder for the nucleus to attract an extra electron to complete the outer shell.

3. He, Na, Mg (increasing order of melting point)
Helium (He) < Sodium (Na) < Magnesium (Mg)
Reason: Helium is a noble gas with only very weak London dispersion forces between its atoms, so it has an extremely low melting point. Sodium and Magnesium are metals held by metallic bonding. Magnesium has a stronger metallic bond than Sodium due to its smaller atomic size and two valence electrons (vs. one for Na), leading to a higher melting point.

4. Chlorine, Sodium, Magnesium (increasing reducing character)
Chlorine (Cl) < Magnesium (Mg) < Sodium (Na)
Reason: A reducing agent loses electrons. Chlorine is a non-metal that tends to gain electrons, so it is a very poor reducing agent. Between the metals, Sodium is a stronger reducing agent than Magnesium because it more readily loses its single valence electron compared to Magnesium, which must lose two electrons and has a higher ionization energy.

Question 14. 

1. State the nature of compounds formed when group 17 elements combine with (i) metals (ii) non-metals. 

2. Why are group 17 elements highly reactive?

Ans:

1. What is the name given to group 17 elements? Why are they called so?

The collective name for the elements in Group 17 is the Halogens.

They are called this because the name is derived from Greek words: ‘hals’ meaning “salt” and ‘gen’ meaning “to produce.” Therefore, “Halogen” literally translates to “salt-producing.” This name is perfectly fitting because all of these elements have a strong tendency to react with metals to form a class of ionic compounds we commonly know as salts. A classic and familiar example is sodium chloride (NaCl), or common table salt, which is produced when chlorine reacts with sodium.

2. Comment on the (i) reactivity (ii) colour (iii) physical state of group 17 elements.

Going down Group 17, from Fluorine to Astatine, we observe clear and predictable trends in their properties.

(i) Reactivity
Fluorine (F) is the most reactive non-metal of all the elements, being extremely aggressive and violent in its reactions. This trend continues with Bromine (Br) and Iodine (I), which are progressively less reactive.

The reason for this is linked to atomic size. As you move down the group, the number of electron shells increases, making the atoms larger. The outer electrons are further from the attractive pull of the nucleus. Because of this, it becomes harder for a larger atom to gain an extra electron to achieve a stable noble gas configuration. Fluorine, being the smallest, attracts and captures an electron most easily, making it the most reactive.

(ii) Colour
The halogens become progressively darker in colour as you move down the group. This is a very visible trend:

• Fluorine (F): A very pale yellow gas.
• Chlorine (Cl): A greenish-yellow gas.
• Bromine (Br): A reddish-brown liquid that easily forms a brownish-orange vapour.
• Iodine (I): A dark grey, almost metallic-looking solid that produces a beautiful violet-purple vapour when heated.

The deepening colour is due to the electrons in the larger atoms being more easily excited to higher energy levels by visible light, absorbing specific wavelengths and resulting in the perception of a darker, more intense colour.

(iii) Physical State
There is a clear change in physical state at room temperature as you descend the group, moving from gases to a solid.

• Fluorine (F) and Chlorine (Cl) are both gases.
• Bromine (Br) is one of only two elements that is a liquid at room temperature (the other being mercury).
• Iodine (I) is a solid.

Question 15. 

1. How many electrons do inert gases have in their valence shells? 

2. Name an element of group 18 which can form compounds.

Ans:

1. Inert gases, also known as noble gases, have eight electrons in their valence shells. The sole exception is helium, which has only two electrons in total and thus two in its valence shell.
2. Xenon is an element from Group 18 that is known to form a variety of compounds, such as xenon tetrafluoride (XeF₄) and xenon trioxide (XeO₃).

Question 16. 

1. Name the gas used in: Filling balloons 

2. Name the gas used in: Light bulbs 

3. Name the gas used in: Bright colour advertising light works

Ans:

1. The gas traditionally used for filling balloons is Helium. It is much lighter than air, which provides buoyancy, and is chemically inert, making it a safer alternative to flammable hydrogen.
2. The gas commonly used in light bulbs is Argon. This inert gas prevents the thin tungsten filament inside the bulb from oxidizing and burning away when it gets extremely hot, thereby extending the bulb’s lifespan.
3. The gas famously used in bright colour advertising light works (neon signs) is Neon, which produces a characteristic bright red-orange glow when an electric current is passed through it. Other gases like argon (for blue or purple) or helium (for gold) are also used to create different colors.

Question 17. 

1. What is the name given to group 17 elements? Why are they called so? 

2. Comment on the (i) reactivity (ii) colour (iii) physical state of group 17 elements.

Ans:

1. State the nature of compounds formed when group 17 elements combine with (i) metals (ii) non-metals.

When group 17 elements, known as halogens, combine with metals, they form ionic compounds. This happens because metals tend to lose electrons, while halogens have a very strong tendency to gain a single electron. The metal atom donates an electron to become a positively charged cation, and the halogen atom accepts it to become a negatively charged anion (a halide ion). The resulting electrostatic attraction between these oppositely charged ions forms an ionic bond, creating compounds commonly known as salts, such as sodium chloride (NaCl).

In contrast, when halogens react with non-metals, they form covalent compounds. Since non-metals also have a tendency to gain electrons, they cannot transfer electrons to the halogen like a metal would. Instead, the two atoms share a pair of electrons to achieve stable electron configurations. This mutual sharing of electrons results in a covalent bond. A common example is hydrogen chloride (HCl), where a hydrogen atom and a chlorine atom share an electron pair.

2. Why are group 17 elements highly reactive?

Group 17 elements are highly reactive due to their strong tendency to gain a single electron. Their atomic structure is just one electron short of a stable, full outer electron shell, which is the electron configuration of the noble gases next to them in the periodic table. This “electron hunger” is driven by their high electronegativity, which is the ability to attract a bonding pair of electrons.

Because they are so close to achieving a stable octet, they readily undergo chemical reactions to acquire that one missing electron. They do this either by forming ionic bonds with metals, where they completely capture an electron, or by forming covalent bonds with non-metals, where they share an electron pair. This powerful drive to complete their valence shell makes them among the most reactive non-metal families in the periodic table.

Question 18. 

Two elements P and Q belong to the same period of the modern periodic table and are in group 1 and group 2, respectively. Compare the following characteristics in the tabular form.

1. Number of electrons in their atoms
2. Their tendency to lose electrons
3. Their metallic characters
4. Formation of their oxides
5. Formulae of their chlorides

Ans:

Characteristic	Element P (Group 1)	Element Q (Group 2)	Comparison (Across the Period)
Number of Valence Electrons	1 (e.g., Na: 2, 8, 1)	2 (e.g., Mg: 2, 8, 2)	Q has one more valence electron than P.
Tendency to Lose Electrons	Higher (requires less energy to lose 1 electron).	Lower (requires more energy to lose 2 electrons).	P has a higher tendency to lose electrons (lower Ionization Energy).
Metallic Character	Higher (more reactive metal).	Lower (less reactive metal).	P is more metallic than Q (Metallic character decreases across a period).
Formation of Oxides	Forms a basic oxide with the formula P2​O (e.g., Na2​O).	Forms a basic oxide with the formula QO (e.g., MgO).	Both form basic oxides, but P’s oxide is typically more basic.
Formulae of Chlorides	Forms a chloride with the formula PCl (e.g., NaCl).	Forms a chloride with the formula QCl2​ (e.g., MgCl2​).	The formula depends on the cation charge (P+ vs. Q2+).