The chapter explains the organization of elements based on their properties.
1. Historical Development:
Dobereiner’s Triads: Groups of three elements with similar properties where the middle element’s atomic weight was roughly the average of the other two.
Newlands’ Law of Octaves: Every eighth element had properties similar to the first, like musical notes. It worked only for lighter elements.
Mendeleev’s Periodic Table: The most significant. He arranged elements by increasing atomic masses and grouped those with similar properties. He boldly left gaps for undiscovered elements (e.g., Eka-silicon, later named Germanium) and predicted their properties.
2. Modern Periodic Law:
The properties of elements are a periodic function of their atomic numbers (not atomic masses). This resolved discrepancies in Mendeleev’s table.
3. Structure of the Table:
Groups: Vertical columns (1 to 18). Elements in a group have the same number of valence electrons and hence, similar chemical properties.
Periods: Horizontal rows (1 to 7). The number of the period signifies the number of electron shells in the elements of that period.
4. Periodic Trends:
Valency: The number of electrons an atom can lose, gain, or share.
Across a period (→): First increases, then decreases.
Down a group (↓): Remains the same.
Atomic Size (Radius):
Across a period (→): Decreases (due to increased nuclear pull).
Down a group (↓): Increases (due to addition of new shells).
Metallic Character:
Across a period (→): Decreases.
Down a group (↓): Increases.
Non-Metallic Character: Opposite of metallic character.
Electronegativity: The tendency to attract electrons.
Across a period (→): Increases.
Down a group (↓): Decreases.
Key Takeaway: The modern periodic table is a powerful tool to predict and compare the properties of elements based on their position.
1) a) State modern periodic law. Name the scientist who stated the law.
b) What is a periodic table? How many groups and periods does the modern periodic table have?
Ans:
The Modern Periodic Law and the Periodic Table
Understanding how elements are organized is one of the most fundamental concepts in chemistry. This order is captured by the Modern Periodic Law and visually represented in the Periodic Table.
a) The Modern Periodic Law: The Modern Periodic Law states that the characteristics of elements, both physical and chemical, repeat in a predictable pattern when the elements are arranged according to their atomic numbers. In simpler terms, an element’s position in the periodic table, which is determined by its atomic number, governs its properties
In simpler terms, this means that when you arrange all the elements in order of increasing atomic number (the number of protons in an atom’s nucleus), you will notice a repeating pattern. Elements with similar properties—like how they react, their conductivity, or their bonding behavior—will appear at regular intervals, or periods.
This crucial update to the older law proposed by Dmitri Mendeleev was established by the English physicist Henry Moseley in 1913. Through his groundbreaking work with the X-ray spectra of elements, Moseley demonstrated that atomic number, not atomic mass, was the true fundamental property governing an element’s characteristics. This discovery resolved inconsistencies in Mendeleev’s table (where elements like argon and potassium were out of order by mass but perfectly in place by atomic number) and provided a rock-solid foundation for the modern table.
b) The Periodic Table
Moving across a horizontal row (a period), each element has one more electron than the last, but they all share the same number of electron shells. Conversely, elements within a vertical column (a group) have an identical number of electrons in their outer shell; this commonality is what gives them their shared chemical properties.
Groups: These are the vertical columns (numbered 1 to 18). Elements within the same group share the same number of electrons in their outermost shell, which is why they have very similar chemical properties. For example, all elements in Group 1 (the alkali metals) are highly reactive and tend to form +1 ions.
Periods: These are the horizontal rows (numbered 1 to 7). As you move from left to right across a period, the atomic number increases by one, and electrons are added to the same outer shell. This leads to a gradual change in properties, moving from metals on the left to non-metals on the right.
For a more dynamic breakdown of these concepts, a helpful video guide can be an excellent resource. It can visually illustrate the crucial differences between periods and groups and provide a clear explanation for why this specific structure so perfectly reflects the periodic nature of the elements.
2) Why does the sodium element of group 1 and chlorine element of group 17 both have valency 1?
Ans: Sodium (Group 1) and chlorine (Group 17) both have a valency of 1, but for opposite reasons related to their need to achieve a stable electron configuration.
Sodium (Na) possesses a single electron in its valence shell. Its valency is 1 because it loses that one electron easily to achieve a stable, full outer shell.
Chlorine (Cl) has 7 electrons in the outermost shell. Its valency is 1 because it gains one electron easily to achieve a stable, full outer shell of 8 electrons.
In short, both elements interact in a way that involves one electron to become stable, which is why their combining capacity or valency is 1.
3) What are horizontal rows and vertical columns in a periodic table known as?
Ans: In the periodic table, the horizontal rows are known as periods, and the vertical columns are known as groups or families. .
- There are seven periods, and each one represents a different electron shell. All elements in the same period have the same number of electron shells, which house their electrons.
- There are 18 groups, and elements in the same group share similar chemical properties because they have the same number of valence electrons (electrons in the outermost shell).
4) Periodicity is observed due to the similar …………………. (Number of valence electrons / atomic number / electronic configuration).
Ans: The periodic table is organized to highlight a fundamental concept in chemistry: elements exhibit repeating patterns in their properties, a principle known as periodicity. This occurs because elements with similar arrangements of their outermost electrons, known as valence electrons, will behave in chemically similar ways.
Valence Electrons: If the periodic table is a beautifully organized mosaic, then the tiles that create its repeating patterns are the valence electrons. The number of these outermost electrons an atom possesses is the master key that unlocks its chemical identity and dictates its placement among the elements. These are the electrons involved in bonding and chemical reactions. Elements within the same vertical group on the table share the same number of valence electrons. For instance, every alkali metal in Group 1 has a single valence electron. This shared characteristic is why they are all extremely reactive and exhibit such comparable chemical behaviors.
Electron Configuration: The broader electron structure of an atom, called its electron configuration, determines its number of valence electrons. Consequently, elements with analogous electron configurations in their outer shell will have similar properties.In chemistry, the gold standard for this type of behavior is perfectly demonstrated by the elements that make up Group 18. These elements, the noble gases, are so stable and unreactive that they have become the fundamental example used to teach this concept in textbooks and classrooms worldwide. Each one has a complete outer electron shell, which is a highly stable arrangement. This recurring, stable configuration is the reason every noble gas is notably unreactive.
5) How does the electronic configuration in atoms change
(i) in a period from left to right?
(ii) in a group top to bottom?
Ans: Moving Across a Period (Left to Right):
- Electron Shells: Stay the same (same principal energy level).
- Atomic Number: Increases by one for each element.
- Valence Electrons: Increase from 1 to 8.
Moving Down a Group (Top to Bottom):
- Valence Electrons: Stay the same.
- Electron Shells: A new shell is added at each step.
- Atomic Number: Increases.
- Chemical Properties: Are similar due to the same number of valence electrons.
6) Correct the statements.
(i) Elements in the same periods have equal valency.
(ii) Valency depends upon the number of shells in an atom.
(iii) Copper and zinc are representative elements.
(iv) Transition elements are placed at the extreme right of the periodic table.
Ans: (i) True. Elements in the same period have different numbers of valence electrons, leading to different valencies.
(ii) True.
(iii) True. Copper and zinc are transition metals (d-block), not main-group representative elements (s and p-block).
(iv) True. Transition elements are located in the central part of the periodic table, not on the far right.
7) Name two elements in each case:
Name two elements in each case:
| (i) Alkali metals | |
| (ii) Akaline earth metals | |
| (iii) halogens | |
| (iv) Inert gas | |
| (v) Transition element | |
| (vi) Lanthanides | |
| (vii) Actinides |
Ans:
| (i) Alkali metals | Sodium and potassium |
| (ii) Akaline earth metals | Calcium and magnesium |
| (iii) halogens | Chlorine and bromine |
| (iv) Inert gas | Neon and Argon |
| (v) Transition element | Iron and Cobalt |
| (vi) Lanthanides | Cerium and Europium |
| (vii) Actinides | Uranium and Neptunium |
8) What do you understand?
(i) Periodicity:
(ii) Typical elements:
(iii) Orbits:
Ans:(i) Periodicity
Periodicity is the predictable and recurring pattern of properties that elements show when they are arranged in order of increasing atomic number. This is the fundamental principle behind the periodic table. For example, as you move across a period (a horizontal row), properties like atomic size and electronegativity change in a systematic way, and these trends repeat from one period to the next. The cause of this phenomenon is the recurring pattern of valence electron configurations in the elements.
(ii) Typical Elements
Typical elements are the elements of the third period of the periodic table, specifically Sodium (Na) through Chlorine (Cl). These elements are considered “typical” because they effectively summarize and represent the chemical properties of their respective groups without any of the unusual or anomalous behaviors seen in the elements of the second period. By studying these elements, one can easily understand the characteristic properties of their entire group.
(iii) Orbits
In the context of early atomic models (like Bohr’s model), an orbit was understood as a fixed, well-defined circular path around the nucleus where an electron moved.
9) Why are noble gases placed in a separate group?
Ans: Noble gases are placed in a separate group on the far right of the periodic table because of their unique and shared properties, most notably their extreme chemical inertness or non-reactivity. This lack of reactivity is a result of their electron configuration.
The properties of elements, including their reactivity, are primarily determined by the number of electrons in their outermost shell, known as valence electrons.
- Stable Electron Shell: Noble gases (Group 18) all have a completely full outer electron shell. This is a very stable and energetically favorable configuration. For example, neon has 8 valence electrons, completing its second electron shell. Helium is the exception, with a full first shell containing 2 electrons. Because their outer shells are already full, these elements have no desire to gain, lose, or share electrons with other atoms to form chemical bonds.
- Zero Valency: Due to their full outer shells, noble gases have a valency of zero, meaning they are not found in a combined state with other elements under normal conditions. This fundamental difference in their bonding behavior sets them apart from all other groups on the periodic table, which are highly reactive and seek to achieve a stable configuration.
- Historical Discovery: The noble gases were discovered relatively late because of their inert nature. Early chemists didn’t find them in compounds and initially thought they didn’t exist. Their unique properties warranted a new, separate group, initially called Group 0, to reflect their zero valency
10) Name two elements you would expect to show chemical reactions similar to calcium. What is the basis of your choice?
Ans: Two elements that you would expect to show chemical reactions similar to calcium are magnesium (Mg) and strontium (Sr).
The basis for this choice is that all three elements—calcium (Ca), magnesium (Mg), and strontium (Sr)—belong to the same group on the periodic table, specifically Group 2, the alkaline earth metals.
Elements in the same vertical column (group) of the periodic table share the same number of valence electrons. For example, calcium, magnesium, and strontium all belong to the same group and each has two valence electrons
Chemical Behavior: The number of valence electrons determines an element’s chemical properties and reactivity. All three elements tend to lose these two valence electrons to form a positive ion with a charge of +2 (e.g., Ca2+ , Mg 2+, Sr2+) to achieve a stable electronic configuration. Because they all undergo this same fundamental chemical process, their reactions with other substances, such as water or acids, are very similar
11) Name the metal(s) and non-metals in the first twenty elements.
Metals:
Non-metals:
Ans:
Metals: Lithium, Beryllium, Sodium, Magnesium, Aluminium, Potassium, Calcium. Non-metals: Hydrogen, Helium, Carbon, Nitrogen, Oxygen, Fluorine, Neon, Phosphorus, Sulphur, Chlorine, Argon.
12) Name the type of elements, which have their:
(i) Outermost shell complete – ………………..
(ii) Outermost shell incomplete – ………………
(iii) two outermost shell incomplete – ………………
(iv)one electron short of octet – ……………………
(v) two electrons in the outermost orbit – ……………
Ans: (i)Outermost shell complete: Noble gases
These elements (Group 18) are very stable and unreactive because they have a full valence shell, typically with eight electrons (an octet), except for Helium, which has two.
(ii)Outermost shell incomplete: Reactive elements (metals, non-metals, and metalloids)
This includes the majority of elements on the periodic table, which actively gain, lose, or share electrons to achieve a stable, complete outermost shell.
(iii)Two outermost shells incomplete: Transition elements
Also known as d-block elements, these metals have partially filled d-orbitals, which means both their outermost and second-to-last shells are incomplete. This is why they exhibit unique properties like variable oxidation states.
(iv)One electron short of octet: Halogens
Halogens (Group 17):
These non-metals are highly reactive because their atoms have seven valence electrons.
Alkaline Earth Metals (Group 2):
These metals are reactive as they have two valence electrons. They readily lose these two electrons to form stable +2 ions.
13) An element has 2 electrons in its N shell.
(i) What is its atomic number?
(ii) State its position in periodic table
(iii) Is it metal or non-metal?
(iv) State the name assigned to this group?
Ans: (i) Atomic Number: 20. Summing these electrons gives a total of 20. In a neutral atom, the number of electrons is equal to the number of protons, so the atomic number is 20.
The atomic number of an element is defined by the number of protons in the nucleus of a single atom of that element. In a neutral atom, the number of protons is equal to the number of electrons.
Count the electrons: The provided electron configuration is K=2, L=8, M=8, N=2. Adding up the electrons in each shell gives: 2 + 8 + 8 + 2 = 20 electrons.
Relate to protons: Since the atom is neutral, the number of protons must also be 20.
Determine the atomic number: Therefore, the atomic number is 20. The element with an atomic number of 20 is Calcium (Ca).
(ii) Position in Periodic Table: The element is in Period 4, Group 2. The number of electron shells (4) corresponds to its period, and the number of valence electrons (2) corresponds to its group.
(iii) Metal or Non-metal: It is a metal. Elements in Group 2 are metals. Metals typically have a small number of valence electrons which they readily lose to form positive ions.
(iv) Group Name: The name assigned to this group is Alkaline Earth Metals.
14) State the valency of the elements of periods 3 and write the formula of their oxides.
Ans: The valency of Period 3 elements increases from 1 to 4 and then decreases to 0. This trend directly influences the formulas of their highest oxides.
- Sodium (Na) to Aluminium (Al): Valency increases from 1 to 3, as they lose valence electrons. Their oxides are Na2O, MgO, and Al2O3.
- Silicon (Si): Has a valency of 4, sharing electrons to form SiO2.
- Phosphorus (P) to Chlorine (Cl): Valency increases from 5 to 7, as they share or gain electrons. Their highest oxides are P4O10 (or P2O5), SO3, and Cl2O7.
- Argon (Ar): Has a valency of 0 and does not form an oxide due to its stable electron configuration.
15) An element A has atomic number 14. To which period does this element belong and how many elements are there in this period.
Ans: Silicon is classified as a metalloid because it exhibits characteristics of both metals and non-metals. It has a metallic luster but is a poor conductor of electricity and heat on its own. It’s this property that makes it invaluable in the electronics industry.
Why Silicon is Important in Technology
Silicon’s significance in technology stems from its ability to act as a semiconductor. This means that its electrical conductivity can be precisely manipulated, a process called doping, by adding impurities like boron or phosphorus. This control over conductivity allows silicon to be the foundational material for creating:
- Integrated circuits: The fundamental components of computer chips, processors, and memory.
- Transistors: The building blocks of modern electronics.
- Solar cells: Which convert sunlight directly into electricity.
The widespread use of silicon in these critical components has made it a cornerstone of the digital age, earning the San Francisco Bay Area the name “Silicon Valley.”
16) Answer the following in respect of element 𝟑𝟏 𝟏𝟓 𝐏
(i) Give its electronic configuration
(ii) To which group and period does it belong?
(iii) What is its valency?
(iv) Is it a metal or non – metal
(v) Is it a reducing agent or oxidizing agent?
(vi) Give its formula with chlorine.
Ans:
(i) Electronic configuration:
2, 8, 5 (or 1s² 2s² 2p⁶ 3s² 3p³)
(ii) Group and period:
Group 15 and Period 3
(iii) Valency:
Its common valencies are 3 and 5.
(iv) Metal or Non-metal:
It is non-metal.
(v) Reducing or Oxidizing agent:
It is primarily a reducing agent.
(vi) Formula with chlorine:
Its common chloride is PCl₃ (Phosphorus trichloride) and PCl₅ (Phosphorus pentachloride).
INTEXT – QUESTION – 2
1) Name any five periods properties.
Ans: Atomic Size:
Across a Period (→): Decreases. More protons pull the electron cloud closer.
Down a Group (↓): Increases. More electron shells are added.
Ionization Energy:
Across a Period (→): Increases.
Down a Group (↓): Decreases.
Electronegativity:
Across a Period (→): Increases. Atoms have a stronger pull on shared electrons.
Down a Group (↓): Decreases. More electron shells shield the nucleus’s pull.
Metallic Character:
Across a Period (→): Decreases. Elements are less likely to lose electrons.
Down a Group (↓): Increases. Elements lose electrons more easily.
In short: As you move right on the table, atoms get smaller, hold electrons tighter, and are less metallic. Moving down, atoms get larger, hold electrons looser, and are more metallic.
2) What do you understand about atomic size? State its unit.
Ans: Atomic size is the size of an atom. Since atoms don’t have a hard edge, we measure it as the atomic radius.Because atoms are so incredibly small, we measure them in tiny units:
Picometers (pm)
Angstroms (Å), hvor 1 Å = 100 pm.
3) Give the trends in atomic size on moving:
(i) down the group
(ii) across the period right to left.
Ans: (i) Down a group: Atomic size increases as you move down a group. This is because each new period adds an extra electron shell, placing the outermost electrons farther away from the nucleus. The inner electrons also create a “shielding effect” that reduces the pull of the nucleus on these outer electrons.
(ii) Across a period (right to left): Atomic size increases as you move from right to left across a period. This is the reverse of the typical trend (left to right, size decreases). This leads to a weaker pull on the outermost electrons, allowing the electron cloud to expand and the atomic size to increase.
4) Arrange the elements of second and third periods in increasing order of their atomic size.
(i) Second Period
(ii) Third Period
Ans: (i) Second Period Elements (Li, Be, B, C, N, O, F, Ne)
Increasing atomic size order: Ne < F < O < N < C < B < Be < Li</mark>
Largest atom: Lithium (Li)
(ii) Third Period Elements (Na, Mg, Al, Si, P, S, Cl, Ar)
Increasing atomic size order: Ar < Cl < S < P < Si < Al < Mg < Na</mark>
Largest atom: Sodium (Na)
Reason for the Trend:
As you move from left to right across a period on the periodic table, the atomic size of elements generally decreases. This is because, while elements within the same period have the same number of electron shells, the nuclear charge (the number of protons in the nucleus) steadily increases. This stronger positive charge exerts a greater electrostatic force of attraction on the electrons in the outermost shell, pulling the entire electron cloud closer to the nucleus and causing the atomic radius to shrink. The additional electrons added as you move from one element to the next are placed in the same principal energy level, so they don’t provide significant shielding from the increased nuclear pull. The net effect is a stronger attractive force that makes the atom more compact.
5) Why is the size of
(i) neon greater than fluorine?
(ii) sodium is greater than magnesium?
Ans: (i) Neon’s Size vs. Fluorine’s Size
Neon is larger than fluorine because we measure their sizes differently. Fluorine’s size is its covalent radius (measured when atoms are bonded and electron clouds overlap). Neon, a noble gas, doesn’t form bonds, so its size is its van der Waals radius (measured between separate, non-overlapping atoms). The van der Waals radius is always larger, making neon appear bigger.
(ii) Sodium’s Size vs. Magnesium’s Size
Sodium is larger than magnesium. However, magnesium has more protons in its nucleus (a higher nuclear charge. This stronger positive charge pulls the electron shells inward more tightly, resulting in a smaller atomic size for magnesium.
6) Which is greater in size?
(i) an atom or a cation
(ii) an atom or an anion
(iii) Fe2+ or Fe3+
Ans: (i) Atom or a Cation
The atom is larger than its cation. Losing electrons reduces electron repulsion and allows the nucleus to pull the remaining electrons closer, shrinking the ion.
(ii) Atom or Anion
An anion is larger than its parent atom. This is because adding electrons increases repulsion within the electron cloud, causing it to expand.
(iii) Fe²⁺ or Fe³⁺
The Fe²⁺ ion is larger than Fe³⁺. With fewer electrons, the nucleus exerts a stronger pull on the remaining electrons in Fe³⁺, drawing them closer and reducing the ion’s size.
7) Metallic character and non-metallic character are periodic properties discussed.
Ans: Metallic Character: This is an element’s ability to lose electrons easily and form positive ions (cations). Elements with strong metallic character are typically good conductors and are shiny and malleable.
Non-metallic character is an element’s tendency to gain electrons and form negative ions (anions).
Trend Down a Group: Decreases.
Reason: As atomic size increases down a group, the nucleus’s pull on new electrons weakens. This makes it harder for the atom to gain electrons
Trends in the Periodic Table
These two properties show opposite trends across the periodic table due to changes in atomic size and nuclear charge.
Across a Period (Left to Right)
Metallic character decreases because the effective nuclear charge increases while the number of electron shells remains the same. This stronger pull from the nucleus makes it harder for the atom to lose electrons.
Non-metallic character increases for the same reason—the increased nuclear charge makes it easier for the atom to attract and gain new electrons.
Down a Group (Top to Bottom)
Metallic character increases because the number of electron shells increases, which places the valence electrons farther from the nucleus. This reduces the attraction from the nucleus, making it easier for the atom to lose electrons.
Non-metallic character decreases as the atomic size increases and the nucleus’s pull on incoming electrons weakens, making it less likely for the atom to gain electrons.
8)Give the trend in metallic character:
(i) across the period left to right,
(ii) down the group top to bottom.
Ans: (i) Across a Period (Left to Right)
Metallic character decreases. The increasing nuclear charge pulls electrons closer, making them harder to lose.
(ii) Down a Group (Top to Bottom)
Metallic character increases. More electron shells are added, which shield the outer electrons. This weaker hold makes it easier to lose electrons, increasing metallic character down a group.
9)State the trends in chemical reactivity:
(i) across the periods left to right
(ii) Down the group
Ans:Across a Period (Left to Right):
Metals (e.g., Na, Mg) become less reactive due to decreasing tendency to lose electrons.
Non-metals (e.g., Cl, O) become more reactive due to increasing tendency to gain electrons.
Down a Group:
Metals (e.g., Group 1): Reactivity increases down the group as atomic size increases, making it easier to lose electrons.
Non-metals (e.g., Group 17): Reactivity decreases down the group as atomic size increases, reducing the tendency to gain electrons.
10) State the trends in physical properties on moving down the group. Give an example to illustrate.
Ans: Trends in Physical Properties Down a Group
Atomic Radius Increases: Each element adds a new electron shell, making the atom larger.
Metallic Character Increases: Outer electrons are farther from the nucleus and are more easily lost, making elements more metallic.
Melting and Boiling Points: The trend depends on bonding.
For metals (e.g., Group 1), they generally decrease.
For non-metals with covalent bonding (e.g., Group 17), they generally increase.
Example: Group 1 (Alkali Metals)
Atomic Radius: Increases from Lithium (Li) to Francium (Fr).
Metallic Character: Increases; Lithium is the least metallic.
Melting Point: Decreases noticeably. Lithium has a relatively high melting point (180°C), while Caesium (Cs) melts at just 28.5°C, low enough to melt in your hand.
11) An element X belongs to the 4th period and 17th group, state.
(i) no valence electrons in it
(ii) name of the element.
(iii) name the family to which it belongs.
(iv) Write the formula of the compound formed when it reacts with 𝟐𝟕 𝟏𝟑 𝐲
Ans: (i) Valence electrons: 7 (Group 17 elements have seven electrons in their outermost shell).
(ii) Name of element: Bromine (Br).
(iii) Family: Halogens. This is the common name for Group 17 elements.
(iv) Formula of compound: AlBr₃. Aluminum (Y, atomic number 13) has a valency of +3 and bromine has a valency of -1, so three bromine atoms are needed to balance one aluminum atom.
12) The given table shows elements with the same number of electrons in its valence shell
State: (i) Whether these elements belong to the same group or period.
(ii) Arrange them in order of increasing metallic character.
| Elements | A | B | C |
| m.p. | 63.0 | 180.0 | 97.0 |
Ans: Based on the melting points and group trends:
(i) These elements belong to the same group because they have the same number of valence electrons.
(ii) Metallic character increases down a group. Since melting points for metals in a group typically decrease downward, the element with the highest melting point (B) is at the top, and the one with the lowest (A) is at the bottom.
Thus, the order of increasing metallic character is:
B < C < A
INTEXT – QUESTION – 3
1) (a) Define the term ‘ionisation potential’
(b) Represent in the form of an equation. In which unit is it measured?
Ans:Ionization Potential (Ionization Energy)
(a) Definition:
Ionization energy is the energy needed to remove the outermost electron from a neutral, gaseous atom, forming a positive ion. A higher ionization energy means it is more difficult to remove an electron, which is key to predicting an element’s reactivity and bonding behavior.
(b) Representation:
The process is shown by the equation:
X(g) + Energy → X⁺(g) + e⁻
X(g) = Neutral gaseous atom
X⁺(g) = Gaseous positive ion
e⁻ = Removed electron
It is measured in kilojoules per mole (kJ/mol) or electron volts per atom (eV/atom).
2) What do you understand about successive ionization energies?
Ans: Successive Ionization Energies
These are the energies needed to remove electrons from an atom one after another.
Key Trend: Each successive ionization energy is larger than the previous one.
Reason: Removing an electron creates a positive ion. The remaining electrons are pulled more strongly by the nucleus (same proton charge, fewer electrons) and experience less electron-electron repulsion, making them harder to remove.
3) State the trends in ionization energy:
(a) across the period
(b) down the group.
Ans: (a) Across a Period (left to right):
Trend: Generally increases.
Reasons:
Nuclear charge increases, pulling electrons closer.
The number of electron shells remains the same, so the stronger pull makes electrons harder to remove.
(b) Ionization Energy Down a Group:
Trend: Ionization energy decreases from top to bottom.
Reason: As you move down a group, each element adds a new electron shell. This puts the outermost electrons farther from the nucleus, weakening the attractive force. Additionally, more inner electrons block, or “shield,” the outer electrons from the nucleus’s positive charge.
4) Name the elements with highest and lowest ionization energies.
Ans:Helium (He) has the highest ionization energy due to its stable, full electron shell and small atomic size.
Francium (Fr) has the lowest ionization energy because its single valence electron is farthest from the nucleus and is heavily shielded by inner electron shells.
5) Arrange the elements of the second and third period in increasing order of ionization energy.
Ans: The ionization energy generally increases across a period and decreases down a group.
Second Period (Li to Ne):
Li < Be > B < C < N < O < F < Ne
Third Period (Na to Ar):
Na < Mg > Al < Si < P < S < Cl < Ar
Note on the exceptions: Be has a higher ionization energy than B, and Mg has a higher ionization energy than Al. This is due to the stable electron configuration of Be and Mg (full s-orbital). Similarly, N has a higher ionization energy than O, and P has a higher ionization energy than S, due to their stable half-filled p-orbitals.
6) (a) Define the term ‘electron affinity’.
(b) Arrange the elements of the second period in increasing order of their electron affinity. Name the elements which do not follow the trend in this period.
Ans: (a) Definition of Electron Affinity
Electron affinity is the energy change that occurs when a neutral atom in its gaseous state gains an electron to form a negative ion. A higher (more exothermic) electron affinity means the atom has a greater tendency to gain an electron.
(b) Order for Second Period & Exception
The general increasing order of electron affinity for the second period is:
Li < Be < B < C < O < N < F < Ne</mark>
Elements that do not follow the trend:
This is because nitrogen has a stable half-filled p-orbital, making it harder to add an electron compared to oxygen.
Beryllium (Be) and Neon (Ne) have very low (endothermic) electron affinities due to their stable full and complete octet configurations, respectively.
7) State the factors on which electron affinity depends.
Ans: 1. Atomic Size: Smaller atoms have a higher electron affinity. The incoming electron is closer to the nucleus and feels a stronger attractive force.
2. Nuclear Charge: A higher effective nuclear charge (the positive charge felt by valence electrons) increases electron affinity. The nucleus attracts the new electron more strongly.
3.Stable Configuration: Atoms with a stable (e.g., full or half-filled) outer shell have low electron affinity. Adding an electron disrupts this stable state.
4.Shielding Effect: More inner electron shells block the nucleus’s attraction.
8)Electron affinity values generally _______ across the periods left to right and ______ down the group top to bottom.
Ans: Electron affinity values generally increases across the periods left to right and decreases down the group top to bottom
9) Give reason:
(a) Electron affinity of halogens is comparatively high,
(b) Electronegativity of chlorine is higher than Sulphur.
Ans:(a) Halogens have a high electron affinity because they are in Group 17 and have a valence electron configuration of ns2np5. This means they are only one electron short of a stable, noble gas configuration (an octet). Gaining this single electron releases a significant amount of energy, making the process highly favorable. Their small atomic size and high effective nuclear charge also contribute to a strong attraction for an incoming electron.
(b) Chlorine (Cl, atomic number 17) has a higher electronegativity than sulfur (S, atomic number 16). This is because both elements are in the same period (Period 3) of the periodic table, and electronegativity generally increases from left to right across a period. As we move from sulfur to chlorine, the number of protons in the nucleus increases from 16 to 17. The additional proton in chlorine’s nucleus exerts a stronger positive pull on the outer electrons. Since the electrons are still in the same main energy shell, there is no significant increase in shielding. This stronger nuclear pull allows the chlorine atom to attract shared electrons in a bond more effectively than a sulfur atom can, giving it a higher electronegativity.
10) Why fluorine has higher E.N. than chlorine?
Ans: Fluorine has a higher electronegativity than chlorine primarily due to its smaller atomic size. Electronegativity is the ability of an atom to attract a shared pair of electrons in a bond. Since fluorine has only two electron shells (a smaller atomic radius), its valence electrons are much closer to the positively charged nucleus. This results in a stronger pull on bonding electrons.
In contrast, the outer electrons of chlorine are further from the nucleus and are also shielded by an additional layer of inner electrons. This increased distance and shielding weaken the nucleus’s attraction to a shared electron pair, giving chlorine a lower electronegativity compared to fluorine.
11) Define the term ‘Electronegativity’ and state its unit.
Ans: Electronegativity is an atom’s ability to attract shared electrons in a bond, influencing a molecule’s polarity. It is a relative, dimensionless value, typically measured on the Pauling scale. Fluorine is the most electronegative element, assigned a value of 4.0, and all other elements are ranked in comparison to it.
12) (a) Name the elements with highest and lowest electronegativity,
(b) State the character of the oxide of period 3.
Ans:(a) Fluorine (F) has the highest electronegativity; Francium (Fr) has the lowest.
(b) Across Period 3, oxides change from basic (e.g., Na₂O, MgO) to amphoteric (Al₂O₃) and then to acidic (e.g., SiO₂, P₄O₁₀, SO₃).
13) Name the periodic property which relates to the:
(a) Amount of energy required to remove an electron from an isolated gaseous atom, (b) character of element which loses one or more electrons when supplied with energy.
(c) tendency of an atom to attract the shared pair of electrons.
Ans: (a) Energy to remove an electron from a gaseous atom.
(b) Tendency of an element to lose electrons.
(c) Atom’s ability to attract shared electrons in a bond.
14) Explain the following:
(a) Group 17 elements are strong non-metals, while group 1 elements are strong metals
(b) Metallic character of elements decreases from left to right in a period while it increases in moving down a group.
(c) Halogens have a high electron affinity.
(d) The reducing power of an element increases down in the group while decreasing in a period.
(e) Size of the atom progressively becomes smaller when we move from sodium (Na) to chlorine(CI) in the third period of the periodic table.
Ans: (a) Alkali metals (Group 1) are strong metals as they easily lose their one valence electron to form positive ions. Halogens (Group 17) are strong non-metals as they readily gain an electron to form negative ions
(b)Metallic character decreases across a period due to increasing nuclear charge, which holds electrons more tightly and makes them harder to lose.
It increases down a group because a larger atomic size weakens the hold on outer electrons, making them easier to lose.
(c) Halogens have a high electron affinity because they readily gain one electron to achieve a stable noble gas configuration, releasing a significant amount of energy in the process.
(d) Reducing power increases down a group as the atomic size increases and ionization energy decreases, making electron loss easier. It decreases across a period as ionization energy increases, making electron loss harder.
(e) Size of the atom becomes smaller from sodium (Na) to chlorine (Cl) in the third period.
This occurs due to two factors acting across a period:
Increasing Effective Nuclear Charge (Zeff): The number of protons in the nucleus increases from Na (11) to Cl (17).
Constant Shielding: The number of inner electron shells remains constant (both are in period 3).
The increasing positive charge of the nucleus pulls the electron cloud closer without a significant increase in shielding, causing the atomic radius to decrease.
EXERCISE: 1
1)
(a) How does the electronic configuration of an atom relate to its position in the modern periodic
table?
(b) Write the number of protons, neutrons and electronic configuration of 39 19 K, 31 15 P. Also state their position in the periodic table.
Ans: (a) Relation between Electronic Configuration and Position
The electronic configuration of an atom directly determines its position in the modern periodic table:
Period Number: The highest principal quantum number (n) of an element’s configuration equals its period number.
Group Number: The number of electrons in the outermost shell (valence electrons) determines its group number. For elements in groups 1, 2, and 13-18, the group number is equal to the number of valence electrons.
Example: An atom with configuration 2,8,6 has n=3 (so Period 3) and 6 valence electrons (so Group 16).
(b) Details for given elements
1. For Potassium, ³⁹₁₉K
Protons: 19
Neutrons: 20 (Mass No. – Atomic No. = 39 – 19)
Electronic Configuration: 2, 8, 8, 1
Position: Period 4, Group 1 (It has 4 shells and 1 valence electron).
2. For Phosphorus, ³¹₁₅P
Protons: 15
Neutrons: 16 (Mass No. – Atomic No. = 31 – 15)
Electronic Configuration: 2, 8, 5
Position: Period 3, Group 15 (It has 3 shells and 5 valence electrons).
2) Fluorine, chlorine and Bromine are put in one group on the basis of their similar properties.
(a) what are those similar properties?
(b) What is the common name of this group or family?
Ans: (a) Similar properties:
Highly reactive nonmetals
Seven valence electrons
Form similar compounds (e.g., NaF, NaCl, NaBr)
Form −1 ions
Diatomic molecules (F₂, Cl₂, Br₂, I₂)
(b) Common name: Halogens
3) What is the main characteristic of the last element in each period of the periodic table? What is the general name of such elements?
Ans: The final element in every period of the periodic table is a noble gas. These elements, found in Group 18, are known for being chemically unreactive or inert. This stability is a direct result of their electron configuration: they have a complete, stable outer electron shell, which means they don’t need to gain, lose, or share electrons to achieve a stable state. Consequently, noble gases like Neon, Argon, and Helium don’t readily form chemical bonds with other elements.
4)According to atomic structure, what determines which element will be the first and which will be the last in a period?
Ans: Elements in a period are arranged by increasing atomic number.A period starts with an alkali metal, which has one electron in a new outer shell. Each subsequent element adds one proton and one electron.The period ends with a noble gas, which has a stable, full outer electron shell.
5) How does the number of:
(i) valence electrons and
(ii) valency varies on moving from left to right in the second period of the periodic table?
Ans:(i) Number of Valence Electrons
The number of valence electrons increases steadily from 1 to 8.
Lithium (Li) has 1 valence electron.
Beryllium (Be) has 2.
Boron (B) has 3.
This pattern continues, adding one electron per element, until Neon (Ne), which has a full outer shell with 8 valence electrons.
This happens because the atomic number increases, meaning each subsequent element has one more proton and one more electron than the previous one. This new electron is added to the same outer shell (the second shell for period 2).
(ii) In the second period (Li to Ne), valency first increases and then decreases.
It rises from Lithium (valency 1) to Carbon (valency 4) as these elements can lose their few valence electrons.
It then falls from Nitrogen (valency 3) to Fluorine (valency 1) because these elements gain electrons to achieve a full octet. Their valency is calculated as (8 – valence electrons).
6) Fill in the blanks:
(a) The horizontal rows in the periodic table are called …………
(b) On moving across a period from right to left in the periodic table, the atomic size of the atom ……….
(c) on moving from right to left in the second period, the number of valence electrons……….
Ans: (a) The horizontal rows in the periodic table are called periods.
(b) On moving across a period from right to left in the periodic table, the atomic size of the atom increases.
(c) On moving from right to left in the second period, the number of valence electrons decreases.
7) An element barium has atomic number 56. Look up its position in the periodic table and answer the following questions.
(a) Is it a metal or a non – metal?
(b) Is it more or less reactive than calcium?
(c) What is its valency?
(d) What will be the formula of its phosphate?
(e) Is it larger or smaller than caesium (Cs)?
Ans: (a) Is it a metal or a non-metal?
Based on the information provided in the subsequent answers, the element in question is a metal. This is confirmed by its reactivity and its tendency to form positive ions with a valency of 2, which are typical characteristics of metallic elements.
(b) It is more reactive than calcium. Being lower in Group 2 of the periodic table, it loses its outer electrons more easily.
(c) What is its valency?
The valency of this element is 2. Valency refers to the combining power of an element and is determined by the number of electrons it needs to lose, gain, or share to achieve a full outer shell. As a member of Group 2, this element has two electrons in its outermost shell. It achieves a stable electron configuration by losing these two electrons to form an ion with a 2+ charge. Consequently, its valency is 2.
(d) The formula of its phosphate is Ba₃(PO₄)₂.
(e) Is it larger or smaller than caesium (Cs)?
Barium (Ba) is smaller than caesium (Cs).
While both atoms are very large, two main periodic trends explain this difference:
Trend Down a Group: Atomic size increases as you move down a group because each successive element has an additional electron shell.
Trend Across a Period: Atomic size decreases as you move from left to right across a period. This is because the increasing positive charge in the nucleus pulls the electron cloud closer, a effect known as increasing effective nuclear charge.
Caesium is in Period 6, Group 1. Barium is its immediate neighbor to the right in Period 6, Group 2. Although they are in the same period, the pull from barium’s stronger nuclear charge is greater than the effect of its slightly larger number of protons. This results in a tighter hold on its electrons, making its atomic radius slightly smaller than that of caesium. In fact, caesium is known to have the largest atomic radius of any naturally occurring element on the periodic table.
8) How do the following change on moving from left to right in a period of the periodic table? Give examples in support of your answer.
(a) atomic structure (electron arrangements)
(b) Chemical reactivity of elements.
(c) Nature of oxides of the elements.
Ans:(a) Atomic Structure (Electron Arrangements)
Across a period, each subsequent element gains one more valence electron than the one before it. However, every element within the same period shares the exact same number of electron shells.
For example, every element in period 3 has three electron shells. Sodium (Na) has 1 valence electron, magnesium (Mg) has 2, aluminium (Al) has 3, and so on, all the way to argon (Ar), which has a complete outer shell with 8 electrons.
(b) Chemical Reactivity of Elements
Reactivity first decreases in metals as it becomes harder to lose the increasing number of valence electrons. Then, reactivity increases in non-metals as it becomes easier to gain electrons to complete the octet.
Example: In period 2, the highly reactive metal Lithium (Li) is on the left. The reactivity decreases to the less reactive metal Beryllium (Be). On the right, the highly reactive non-metal Fluorine (F) is more reactive than the less reactive non-metal Oxygen (O).
(c) Nature of Oxides of the Elements
The nature of oxides changes from basic on the left, through amphoteric in the centre, to acidic on the right.
Example: In period 3:
Sodium oxide (Na₂O) is basic.
Aluminium oxide (Al₂O₃) is amphoteric (can act as both acid and base).
Sulphur dioxide (SO₂) is acidic.
9) This question refers to the elements of the periodic table with atomic number from 3 to 18. Some of the elements are shown by letters, but the letters are not the usual symbols of the elements.
| 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:
(a) have the most electronegative element.
(b) is a halogen?
(c) is an alkali metal?
(d) is an element with valency 4?
(e) have least ionization energy?
(f) has the least atomic size in period 3.
Ans: (a) Noble gases- H and P
(b) Halogens- G and O
(c) Alkali metals – A and I
(d) D and L have valency of 4
(e) I with atomic number 11.
(f) Cl has the least atomic size in period 3 with atomic number 17.
10) In group I of the periodic table, three elements X,Y and Z have ionic radii 1.33 A˚, 0.95 A˚ and 0.60 A˚ respectively. Given a reason, arrange them in the order of increasing atomic number in the group.
Ans: The correct order of increasing atomic number is Z, Y, X.
This order is determined by analyzing the given ionic radii. In the periodic table, a fundamental trend within a group (a vertical column) is that the ionic radius increases as you move down the group. This happens because each subsequent element has an extra full shell of electrons, making its ion larger.
Since the provided ionic radii show that Z is the smallest (0.60 Å) and X is the largest (1.33 Å), it indicates that Z is positioned highest in its group, meaning it has the smallest atomic number. Conversely, X is positioned lowest in the group and has the largest atomic number. Therefore, the sequence from smallest to largest atomic number follows the sequence from smallest to largest ionic size: Z, then Y, then X.
11) How does the chemical reactivity of:
(a) alkali metals vary?
(b) halogens vary?
Ans: (a) The chemical reactivity of alkali metals increases as you go down the group because it gets progressively easier for them to lose their single valence electron. This trend is due to two main factors: atomic size and electron shielding.
With each step down the group, a new electron shell is added, making the atoms larger. This increased distance means the outermost electron is farther from the positive pull of the nucleus. Furthermore, the inner electron shells act as a barrier, shielding the outer electron from the full attractive force of the nucleus.
As a result, the hold of the nucleus on the outermost electron becomes weaker. This is measured as ionization energy, which decreases down the group. Since alkali metals react primarily by losing this one electron to form positive ions, a weaker hold directly translates to a more reactive metal. This is why potassium reacts more vigorously with water than sodium, which is in turn more reactive than lithium.
(b) The chemical reactivity of halogens decreases down the group because their ability to gain an electron to complete their outer shell becomes more difficult. This is also a consequence of increasing atomic size and the effect of electron shielding.
As the size of the halogen atoms increases down the group, the outer shell is located much farther from the nucleus. The attractive force that the nucleus exerts on an incoming electron weakens significantly over this greater distance. Additionally, the increasing number of inner electron shells provides more shielding, further reducing the effective nuclear charge experienced by an external electron.
This weaker attraction for electrons means that larger halogens like bromine and iodine are less able to pull an electron away from another atom to form a negative ion. Therefore, their tendency to undergo reactions—which for halogens involves gaining an electron—decreases. This makes fluorine the most reactive halogen and iodine considerably less reactive.
12) An element X belong to 3rd periods and group II of the periodic table state:
(a) the number of valence electrons,
(b) the valency,
(c) name of the element,
(d) whether it is a metal or a non-metal.
Ans:(a) The number of valence electrons:
The element has 2 valence electrons. This is because all elements in group II have two electrons in their outermost shell.
(b) An element’s valency, or combining power, is determined by the number of electrons it needs to lose, gain, or share to achieve a stable electron configuration. In this case, the element has two valence electrons in its outermost shell. To become stable, it will readily lose these two electrons, resulting in a positive charge of 2+ and a valency of 2. This behavior is characteristic of alkaline earth metals like magnesium or calcium
(c) Name of the element:
The element is Magnesium (Mg). The group number (II) tells us it has 2 valence electrons, and the period number (3) tells us its outermost electrons are in the third energy shell. Tracing these coordinates on the periodic table leads to magnesium.
(d) Whether it is a metal or a non-metal:
It is a metal. All elements in group II (the alkaline earth metals) are classified as metals. They are solid, shiny, good conductors of heat and electricity, and tend to lose electrons to form positive ions (cations).
13) The electronic configuration of an element T IS 2, 8, 8, 1.
(a) What is the group number of T?
(b) What is the period number of T?
(c) How many valence electrons are there in an atom of T?
(d) What is the valency of T?
(e) Is it a metal or a non-metal?
Ans: (a) The group number of T is 1.
Element T is part of Group 1 in the periodic table, which is the group of alkali metals. Atoms of elements in this group have a single electron in their outermost shell. When they lose that electron, they form a positively charged ion with a +1 charge, denoted as T⁺.
Due to this strong tendency to lose an electron, Group 1 elements are highly reactive. They are generally soft metals with relatively low melting points. A well-known characteristic of these elements is their vigorous and often explosive reaction with water, similar to the behavior of sodium or potassium.
(b) The period number of T is 4.
T is in the 4th period, meaning its atoms have electrons filling up to the fourth energy level. For a Group 1 element, this suggests it is potassium.
(c) There is 1 valence electron in an atom of T.
Since T belongs to Group 1, it has one valence electron in its outermost shell.
(d) The valency of T is 1.
Valency is the combining capacity. T loses its single valence electron to form a positive ion, so its valency is 1.
(e) It is a metal.
Group 1 elements are metals. They are shiny, conduct heat and electricity, and form positive ions by losing electrons..
14) Arrange the elements of group 17 and group 1 according to the given conditions. (a) Increasing order of atomic size,
(b) Increasing non – metallic character
(c) Increasing ionization potential
(d) Increasing electron affinity
(e) Decreasing electro negativity.
Ans: (a) Increasing order of atomic size
Group 1 (Li, Na, K, Rb, Cs, Fr): Li < Na < K < Rb < Cs < Fr
Group 17 (F, Cl, Br, I, At): F < Cl < Br < I < At
Explanation: Atomic size increases down a group because new electron shells are added, increasing the distance between the nucleus and the outermost electrons.
(b) Increasing non-metallic character
Group 1 (Li, Na, K, Rb, Cs, Fr): Fr < Cs < Rb < K < Na < Li
Group 17 (F, Cl, Br, I, At): At < I < Br < Cl < F
Explanation: Non-metallic character decreases down a group. Fluorine (F) is the most non-metallic element in the entire periodic table. In Group 1, Lithium (Li) is the least metallic and has the most non-metallic character within its group.
(c) Increasing ionization potential (ionization energy)
Group 1 (Li, Na, K, Rb, Cs, Fr): Fr < Cs < Rb < K < Na < Li
Group 17 (F, Cl, Br, I, At): At < I < Br < Cl < F
Explanation: Ionization energy is the energy needed to remove an electron. It decreases down a group because the outer electrons are farther from the nucleus and are more shielded, making them easier to remove.
(d) Increasing electron affinity
Group 17 (F, Cl, Br, I, At): At < I < Br < F < Cl
Explanation: Electron affinity is the energy change when an electron is added. While the trend generally decreases down the group, Chlorine (Cl) has a higher electron affinity than Fluorine (F). This is because fluorine’s small size leads to significant electron-electron repulsion in its compact 2p orbital, making it slightly harder to add an electron compared to the larger chlorine atom.
For Group 1 (alkali metals), electron affinity values are very low and do not follow a clear, strong trend useful for simple comparisons like this. Their tendency is to lose electrons, not gain them.
(e) Decreasing electro negativity
Group 1 (Li, Na, K, Rb, Cs, Fr): Li > Na > K > Rb > Cs > Fr
Group 17 (F, Cl, Br, I, At): F > Cl > Br > I > At
Explanation: Electronegativity is the ability of an atom to attract electrons. It decreases down a group because the increased atomic size and shielding effect mean the nucleus has a weaker pull on bonding electrons. Fluorine is the most electronegative element.
15) Complete the following sentences choosing the correct word or words from those given in brackets at the end of each sentence:
(a) The properties of the elements are a periodic function of their …………… (atomic number, mass number, reative atomic mass).
(b) Moving across a ………….. of the periodic table the elements show increasing ………………..character (group, period, metallic, non-metallic).
(c) The elements at the bottom of a group would be expected to show …….. metallic character than the element at the top. (less, more).
(d) The similarities in the properties of a group of elements are because they have the same ………………… (electronic configuration, number of outer electrons, atomic numbers).
Ans:
Complete the following sentences choosing the correct word or words from those given in brackets at the end of each sentence:
(a) The properties of the elements are a periodic function of their atomic number (atomic number, mass number, reative atomic mass).
(b) Moving across a period of the periodic table the elements show increasing non-metallic character (group, period, metallic, non-metallic).
(c) The elements at the bottom of a group would be expected to show more metallic character than the element at the top. (less, more).
(d) The similarities in the properties of a group of elements are because they have the same number of outer electrons (electronic configuration, number of outer electrons, atomic numbers).
16) Give reasons for the following:
(a) The size of the anion is greater than the size of the parent atom.
(b) argon atoms are bigger than chlorine atoms.
(c) Ionisation potential of the element increases across a period.
Ans: (a) Anions are larger than their parent atoms because adding an electron increases electron-electron repulsion. The nuclear charge remains the same but is now shared over more electrons, reducing the pull on each electron and allowing the cloud to expand.
(b) Argon atoms are bigger than chlorine atoms.
This is a common observation when comparing a halogen and a noble gas in the same period. While atomic size generally decreases across a period due to increasing nuclear charge pulling the electron shells closer, this trend applies to atoms that form bonds. Argon is a noble gas and has a full octet, making it chemically inert. Its atomic size is measured by its van der Waals radius, which is determined by the distance between the nuclei of two non-bonded atoms in a solid state. Chlorine, however, is a halogen and its atomic size is measured by its covalent radius, which is half the distance between the nuclei of two bonded atoms. The van der Waals radius is typically larger than the covalent radius because it accounts for the repulsion between non-bonded electron clouds. Therefore, even though chlorine has a higher nuclear charge, argon’s larger van der Waals radius makes its atomic size appear greater in comparison.
(c) This stronger positive charge pulls electrons closer, making them harder to remove. Since electrons are added to the same shell, they don’t shield each other effectively from this growing nuclear pull.
17) Which element has:
(a) two shells, both of which are completely filled with electrons?
(b) the electronic configuration 2, 8, 3?
(c) a total of three shells with five electrons in its valence shell?
(d) a total of four shells with two electrons in its valence shell?
(e) twice as many electrons in its second shell as in its first shell?
Ans: (a) Neon
(b) Aluminum
(c) Phosphorus
(d) Calcium
(e) Carbon
18) Name
(a) An alkali metal in period 3 and halogen in period 2.
(b) The noble gas with 3 shells.
(c) The non-metals present in period 2 and metals in period 3.
(d) The element of period 3 with valency 4
(e) The element in period 3 which does not form oxide
(f) The element of lower nuclear charge out of Be and Mg.
(g) Which has higher E.A. fluorine or Neon.
(h) Which has a maximum metallic character Na, Li or K.
Ans:
(a) SOL: Na and F
(b) SOL: Argon
(c) SOL: C, N, O and F are non-metals present in period 2 while Na, Mg and Al are metals in period 3.
(d) SOL: Silicon
(e) SOL: Argon
(f) SOL: Mg
(g) SOL: Fluorine
(h) SOL: K
19) Chlorine in the periodic table is surrounded by the elements with atomic numbers 9, 16, 18 and 35.
(a) Which of these have physical and chemical properties resembling chlorine.
(b) Which is more electronegative than chlorine
Ans: (a) Elements resembling chlorine: Fluorine (atomic number 9) and Bromine (atomic number 35) share similar properties with chlorine. They are all in Group 17 (halogens), have comparable valence electron configurations, and are highly reactive nonmetals that form diatomic molecules and metal salts.
(b) Element more electronegative than chlorine: Fluorine (atomic number 9) is more electronegative. Electronegativity increases up a group, and fluorine, being above chlorine in Group 17, has the highest electronegativity of all elements.
20) (a) State the number of elements in periods 1, Periods 2, and Period 3, of the periodic table.
(b) name the elements in period 1.
(c) What is the common feature of the electronic configuration of the elements at the end of period 2, and period 3?
(d) if an element is in group 17, it is likely to be ………….. (Metallic / non-metallic) in character while with one electron in its outermost energy level (shell), then it is likely to be ……………(Metallic / Non-metallic)
Ans: (a) Period 1 has 2 elements while period 2 and period 3 have 8 elements each.
(b) Hydrogen and helium
(c) The elements at the end of period 2 and Period 3 have 8 electrons in its outermost shell.
(d) if an element is in group 17, it is likely to be Non metallic (Metallic / non-metallic) in character while with one electron in its outermost energy level (shell), then it is likely to be metallic (Metallic / Non-metallic)
21) First ionization enthalpy of two elements X and Y are 500KJ mol-1 and 375KJ mol-1 respectively. Comment about their relative position in a group as well as in a period.
Ans: Relative Position in a Period:
Since ionization enthalpy increases from left to right across a period, the element with the higher value (X, 500 kJ/mol) would be located to the right of the element with the lower value (Y, 375 kJ/mol) in the same period.
Relative Position in a Group:
Since ionization enthalpy decreases down a group, the element with the higher value (X, 500 kJ/mol) would be placed above the element with the lower value (Y, 375 kJ/mol) in the same group.
In summary: X is either to the right of Y in a period or above Y in a group.
22) A metal M forms as oxide having the formula M2O3. It belongs to the third period. Write the atomic number and valency of the metal.
Ans:The metal M belongs to the third period and forms an oxide with the formula M₂O₃.
Atomic number: 13
Valency: 3
The metal is Aluminium (Al). Its oxide, Al₂O₃, confirms its valency of 3.
23) Explain why are the following statements not correct:
(a) All groups contain metals and non metals.
(b) Atoms of elements in the same group have the same number of electron(s)
(c) Non- metallic character decreases across a period with increase in atomic number
(d) Reactivity increases with atomic number in a group as well as in a period.
Ans: (a) False. While groups can contain a mix of elements, this is not a universal rule. A clear example is Group 1, known as the Alkali Metals, which is composed exclusively of metals, showing that an entire group can share the same general classification.
(b) Incorrect. While elements in the same group share the same number of valence electrons (outer shell electrons), their total number of electrons increases with atomic number down the group.
(c) The statement is false. For instance, elements on the period’s left, like sodium, are metallic. Moving right, elements become metalloids and then strong non-metals, such as chlorine. This happens because the ability to gain electrons strengthens across a period.
(d) False. Reactivity does not follow one single pattern down all groups. The trend depends entirely on whether the group contains metals or non-metals. For metals, reactivity increases down a group. For non-metals, reactivity decreases down a group.
24) Arrange the following in order of increasing radii:
(a) CI- , CI
(b) Mg2+, Mg, Mg+
(c) N, O, P
Ans: (a) Cl < Cl¯
(b) Mg2+ < Mg+ < Mg
(c) O < N < P
25) Which element from the following has the highest ionization energy? Explain your choice.
(a) P, Na, CI
(b) F, O, Ne
(c) Ne, He, Ar
Ans: (a) Among the elements P (Phosphorus), Na (Sodium), and Cl (Chlorine), chlorine has the highest ionization energy. This trend occurs because ionization energy increases as you move from left to right across a period on the periodic table. Sodium, positioned on the left, easily loses its outer electron, giving it a low ionization energy. In contrast, chlorine, found on the right, has a strong hold on its electrons, requiring significantly more energy to remove one.
(b) Among F, O, and Ne, neon (Ne) has the highest ionization energy. Neon is a noble gas with a full and stable electron shell, requiring the most energy to remove an electron, even more than its neighbors in the period.
(c) Among Ne, He, and Ar, helium (He) exhibits the highest ionization energy. Despite all being noble gases, helium’s electron is in the first shell, closest to the nucleus. This results in the strongest attractive force and the greatest energy requirement to remove an electron.
26) Choose the correct answer.
(a) An element in period 3 whose electron affinity is zero
(i) Sulphur
(ii)Sodium
(iii) Neon
(iv) Argon
(b) An alkaline earth metal
(i) Lead
(ii) potassium
(iii) calcium
(iv) Copper
(c) An element with highest ionization potential
(i) Calesium
(ii) Fluorine
(iii) Helium
(iv) Neon
Ans:
(a) (iv) Argon
(b) (iii) Calcium
(c) (iii) Helium
27) The table given below represents the first three periods. Study the table and answer the question as given below
(a) Write the formula of the sulphate of the element with atomic number 13
(b) what type of bonding will be present in the oxide of the element with atomic number 17?
(c) Which feature of the atomic structure accounts for the similarities in the chemical properties of the elements in group 7A of the periodic table?
(d) Name the element which has the highest ionization potential.
(e) How many electrons are present in the valency shell of the element with atomic number 18?
(f) What is the name given to the energy released when an atom in its isolated gaseous state accepts an electron to form an anion?
(g) Fill in the blanks: The atomic size ………. as we move from left to right across the periods, because the ……… increases but the …………… remains the same
Ans:
(a) (Al)2(SO4)3
(b) Covalent bonding
(c) Same number of valence electrons
(d) Helium
(e) 8
(f) Electron affinity
(g) The atomic size Decreases as we move from left to right across the periods, because the atomic number increases but the number of shells remains the same
28) The electro negativities (according to pauling) of the elements in periodic table are as follows with the elements arranged in alphabetical order:
| AI | CI | Mg | Na | P | S | Si |
| 1.5 | 3.0 | 1.2 | 0.9 | 2.1 | 2.5 | 1.8 |
(a) Arrange the elements in the order in which they occur in the periodic table from left to right. (The group 1 element first, followed by the group 2 element and so on, up to group 7)
(b) Choose the word or phrase from the brackets which correctly completes each of the following statements:-
(i) The element below sodium in the same group would be expected to have a………. (lower/higher) electro-negativity than sodium and the element above chlorine would be expected to have a ……. (lower/ higher) ionization potential than chlorine.
(ii) On moving from left to right in a given period, the number of shells (remains the same/ increases/ decreases).
(iii) On moving down a group, the number of valence electrons (remains the same/ increases/ decreases).
Ans: (a) The elements in the third period, arranged from left to right by increasing electronegativity, are: Sodium (Na), Magnesium (Mg), Aluminium (Al), Silicon (Si), Phosphorus (P), Sulfur (S), and Chlorine (Cl).
(b) Completing the statements:
(i) The element below sodium in the same group is expected to have a lower electronegativity than sodium. The element above chlorine is expected to have a higher ionization energy than chlorine.
(ii) On moving from left to right in a given period, the number of electron shells remains the same.
(iii) On moving down a group, the number of valence electrons remains the same.
29) Parts (a) to (e) refer to changes in the properties of elements on moving from left to right across a period of the periodic table. For each property, choose the correct answer.
(a) The non-metallic character of the elements:
(i) decreases
(ii) increases,
(iii)remains the same,
(iv) depends on the period
(b) The electronegativity:
(i) depends on the number of valence electrons,
(ii) remains the same,
(iii)decreases,
(iv) increases.
(c) The ionization potential:
(i) goes up and down
(ii) decreases
(iii) increases
(iv) remains the same
(d) The atomic size:
(i) decreases,
(ii) increases,
(iii)remains the same,
(iv) sometimes increases and sometimes decreases.
(e) The electron affinity of the elements in groups 1 to 7:
(i) goes up and then down.
(ii) decreases and then increases,
(iii) increases,
(iv) decreases.
Ans: (a) Increases
(b) Increases
(c) Increases
(d) Decreases
(e) Increases
30) The elements of one short period of the periodic table are given below in order from left to right:
Li Be B C O F Ne
(a) To which period do these elements belong?
(b) One element of this period is missing. Which is the missing element and where should it be placed?
(c) Which one of the elements in this period shows the property of catenation?
(d) Place the three elements fluorine, beryllium and nitrogen in the order of increasing electronegativity.
(e) Which one of the above elements belongs to the halogen series?
Ans: (a) Period 2
(b) Nitrogen (N), between carbon and oxygen
(c) Carbon
(d) Be< N< F
(e) Fluorine
31) A group of elements in the periodic table are given below (boron is the first member of the group and thallium is the last). Boron, Aluminium, Gallium, Indium, Thallium. Answer the following questions in relation to the above group of elements:
(a) Which element has the most metallic character?
(b) Which element would be expected to have the highest electronegativity?
(c) If the electronic configuration of aluminium is 2, 8, 3, how many electrons are there in the outer shell of thallium
(d) The atomic number of boron is 5. Write a chemical formula of the compound formed when boron reacts with chlorine.
(e) Will the elements in the group to the right of this boron group be more metallic or less metallic in character? Justify your answer.
Ans: (a) Of the elements listed, thallium possesses the most metallic character. This is because metallic character increases down a group in the periodic table. As the lowest element in its group, thallium exhibits the strongest metallic properties.
(b) Which element would be expected to have the highest electronegativity?
Boron would be expected to have the highest electronegativity. Electronegativity generally decreases down a group as the atomic size increases, reducing the nucleus’s ability to attract bonding electrons.
(c) If the electronic configuration of aluminium is 2, 8, 3, how many electrons are there in the outer shell of thallium?
All elements in this group (Group 13) have 3 electrons in their outermost shell. Therefore, thallium also has 3 electrons in its outer shell.
(d) The atomic number of boron is 5. Write a chemical formula of the compound formed when boron reacts with chlorine.
Boron and chlorine form the compound BCl₃ (Boron trichloride). Boron has a valency of 3 and chlorine has a valency of 1, so they combine in a 1:3 ratio.
(e) Elements to the right will be less metallic. This is because metallic character decreases moving from left to right across a period. Elements in Group 14 have a greater tendency to gain, rather than lose, electrons compared to Group 13 elements.
32) Select the correct answer from the choice A, B, C, D which are given. Write down only the letter corresponding to the correct answer.
With reference to the variation of properties in the periodic table, which of the following is generally true?
A. Atomic size increases from left to right across a period.
B. Ionization potential increases from left to right across a period.
C. Electron affinity increases going down a group.
D. electro-negativity increases going down a group.
Ans: The correct answer is B.
Explanation: This is because, within the same period, the atomic size decreases while the nuclear charge increases. As a result, the outermost electrons are held more tightly, making them harder to remove.
Option A is false: Atomic size decreases from left to right across a period.
Option C is false: Electron affinity generally decreases going down a group.
Option D is false: Electronegativity decreases going down a group.
