Periodic Table - Some Important Questions

 

A. Multiple Choice Questions

1.   Which of the following elements has the highest electronegativity?

a.   Chlorine (Cl)

b.   Fluorine (F)

c.   Oxygen (O)

d.   Sodium (Na)

 

2.   As you move down a group in the periodic table, the atomic radius generally:

a.   Increases

b.   Decreases

c.   Stays the same

d.   Doubles

 

3.   What is the first ionization enthalpy of an element?

a.   Energy released when an electron is added

b.   Energy required to remove an electron

c.   Energy required to add an electron

d.   Energy released when an atom is formed

4.   Which group of elements is known for having low electron gain enthalpy?

a.   Alkali metals

b.   Alkaline earth metals

c.   Noble gases

d.   Halogens

5.   Which of the following trends is true as you move across a period from left to right?

a.   Atomic radius increases

b.   Ionization enthalpy decreases

c.   Electronegativity increases

d.   Reactivity of metals increases

Ans.:

1.   Fluorine (F)

2.   Increases

3.   Energy required to remove an electron

4.   Noble gases

5.   Electronegativity increases

 

B. Fill in the blanks.

1.  The ability of an atom to attract electrons towards itself in a chemical bond is called ………………………………..

2.     As you move ……………………………. in the periodic table, the atomic radius generally decreases.

3.  The energy required to remove an electron from an atom in the gas phase is known as…………………………………….

4.   Elements in Group 1 of the periodic table are known as……………………… and they are highly reactive.

5.  The measure of the energy change when an electron is added to an atom is referred to as………………………………...

 

Ans.: 

1.   electronegativity

2.   from left to right across a period

3.   ionization enthalpy

4.   alkali metals

5.   electron gain enthalpy

 

C. True or False

1.  Electronegativity generally decreases as we move down a group in the periodic table.

2.  The atomic radius increases as we move from top to bottom within a group.

3.  Ionization enthalpy tends to increase as you move across a period from left to right.

4.  Noble gases have a very low electron gain enthalpy because they have full valence shells and do not readily accept electrons.

5.  Alkali metals have lower ionization energies compared to alkaline earth metals.

Ans.:

1.  False

2.  False 

3.  False

4.  False

5.  True

  

D. Some Important Questions

Q1. Why electronegativity decreases down the period?

Ans.: Electronegativity is a measure of an atom’s ability to attract and hold onto electrons in a chemical bond.

Reason: The trend of decreasing electronegativity from top to bottom in a group of the periodic table can be attributed to several factors:

·   Increased Atomic Radius: As we move down a group, the number of electron shells increases, which leads to a larger atomic radius. The increased distance between the nucleus and the valence electrons reduces the nucleus’s ability to attract bonding electrons.

·  Shielding Effect: With more inner electron shells present, the outer electrons experience increased shielding from the positive charge of the nucleus. This shielding effect means that the effective nuclear charge is reduced, making it harder for the nucleus to attract additional electrons.

·   Nuclear Charge: Although the number of protons (nuclear charge) increases as we move down a group, the effect of increased distance and shielding outweighs this increase in charge. As a result, the overall ability of the atom to attract electrons decreases.

 

Conclusion: it can be concluded that electronegativity decreases down the group due to increased atomic radius, increased shielding effect and decreased nuclear charge.

 

Q2. Why electronegativity increases across the period?

Ans.: Electronegativity is a measure of an atom’s ability to attract and hold onto electrons in a chemical bond.

Reason: Electronegativity increases from left to right across the periodic table primarily due to two factors: increasing nuclear charge and decreasing atomic radius.

·   Increasing Nuclear Charge: As  we move from left to  right across a period, the  number of protons in the nucleus increases. This increase in positive charge attracts the electrons more strongly. Since electronegativity is a measure of an atom’s ability to attract and hold onto electrons, a greater nuclear charge results in higher electronegativity.

·   Decreasing Atomic Radius: As the atomic number increases, electrons are added to the same principal energy level while the nuclear charge increases. This leads to a stronger attraction between the nucleus and the electrons, pulling them closer and reducing the size of the atom (atomic radius). A smaller atomic radius means that the valence electrons are closer to the nucleus, enhancing the atom’s ability to attract additional electrons.

 

Conclusion: Together, these factors contribute to the trend of increasing electronegativity from left to right across a period in the periodic table.

 

Q3.Why does ionization energy increase across a period?

Ans.: Ionization energy increases across a period due to several key factors related to atomic structure:

·   Nuclear Charge: As we move from left to right across a period, the number of protons in the nucleus increases. This increased positive charge attracts the negatively charged electrons more strongly.

·   Electron Shielding: Although additional electrons are added as we move across a period, they are added to the same principal shell. This means that the shielding effect does not increase significantly. As a result, the effective Nuclear charge felt by the outermost electrons Increases.

·   Atomic Radius: The increased nuclear charge pulls the electrons closer to the nucleus, which reduces the atomic radius. A smaller atomic radius means that outer electrons are held more tightly by the nucleus, making it harder to remove them.

·   Stability of Electron Configuration: As you approach noble gases at the end of a period, atoms become more stable due to their electron configurations. Removing an electron from a stable configuration requires more energy.

 

Conclusion: Overall, these factors combine to make it increasingly difficult to remove an electron as you move across a period, resulting in higher ionization energies.

 

Q4. Compare the reactivity of non-metals as you move down Group 17 of the periodic table.

Ans.: As we move down group 17 of the periodic table, the reactivity of non-metals decreases: 

Fluorine is the most reactive element and Iodine is the least reactive element. The reactivity of the halogens decreases from fluorine to iodine. The order is F > Cl > Br > I. 

Reason: This is because the atomic radius increases as we move down the group, which decreases the attraction for valence electrons of other atoms. This is also due to a decrease in electronegativity, which means there is less electron "pulling.

 

Q5. Why reactivity of metals decreases down the group and of non-metals increase?

Ans.: The trends in reactivity for metals and non-metals as you move down their respective groups in the periodic table are rooted in their atomic structure and electron configurations.

Metals: Reactivity decrease down the group:

·   Atomic Size: As you  move down a group, the atomic size increases due to the addition of electron shells. This means that the outermost electrons are farther from the nucleus.

·   Shielding Effect: Increased electron shells lead to greater electron shielding, which reduces the effective nuclear charge felt by the outermost electrons. As a result, these electrons are held less tightly.

·    Easier to Lose Electrons: While it may seem counterintuitive, for metals (which tend to lose electrons), the increased distance and shielding make it easier to lose outer electrons, but the overall ability to react decreases because the energy needed to remove electrons can still be high due to the nuclear attraction being weaker than in lighter metals.

 

Non-Metals: Reactivity increase down the group:

·    Atomic Size: Similar to metals, non-metals also experience an increase in atomic size down the group.

·  Electron Affinity: Non-metals gain electrons to achieve a stable electronic configuration. As the atomic size increases, the increased distance from the nucleus can make it easier for these atoms to attract additional electrons, as they are less tightly held by the nucleus.

·    Increased Reactivity: For non-metals, the ability to gain electrons increases down the group, enhancing their reactivity.

 

Conclusion: while metals become less reactive down the group primarily due to the increased difficulty in losing electrons effectively, non-metals become more reactive as their ability to gain electrons improves with increased atomic size.

 

Q6. Why the electron gain enthalpy of noble gases is typically positive or less negative compared to other elements?

The electron gain enthalpy of noble gases is typically positive or less negative compared to other elements due to several key factors:

·   Stable Electron Configuration: Noble gases have a complete valence shell, which means they are already in a very stable electronic state. Their outermost shells are filled (e.g., helium has 2 electrons, while others like neon and argon have 8), making them less inclined to gain additional electrons.

·   High Ionization Energy: The ionization energies of noble gases are very high because it requires a significant amount of energy to remove an electron from their stable configuration. Consequently, the addition of an electron (which would lead to an unstable anion) does not release as much energy as it does for elements that are closer to achieving a stable electron configuration.

·   Positive Electron Gain Enthalpy: For noble gases, adding an electron can lead to a situation where the electron experiences repulsion from the filled valence shell, which is already energetically favorable. This repulsion can result in a positive or less negative electron gain enthalpy.

·   Lack of Tendencies to Form Anions: Unlike halogens or other elements that readily gain electrons to achieve a stable octet, noble gases do not have a strong tendency to form anions because they already possess a stable electron arrangement.

 

Conclusion: the positive or less negative electron gain enthalpy of noble gases reflects their inherent stability and reluctance to accept additional electrons.

 

Q7. Arrange Boron, carbon, nitrogen & oxygen elements in order of increasing first ionization energy?

Ans.: The order of increasing first ionization energy for boron (B), carbon (C), nitrogen (N), and oxygen (O) is as follows:

·  Boron (B): The first ionization energy is the lowest because it has a relatively low nuclear charge and fewer protons to attract the outermost electron.

·  Carbon (C): Carbon has a higher ionization energy than boron due to an increased nuclear charge and a more stable electron configuration.

·   Oxygen (O): Oxygen has a higher  ionization  energy than carbon. The  increased  nuclear charge and the fact that oxygen has a half-filled p subshell makes it more stable.

·   Nitrogen (N): Nitrogen has the highest first ionization energy among these four elements, despite being next to oxygen in the periodic table. This is due to its half-filled p subshell, which provides extra stability, requiring more energy to remove an electron.

So, the order is: B < C < O < N

Q8. Explain why atomic radius increases down the group?

Ans.: The atomic radius increases down a group in the periodic table primarily due to two key factors:

·   Addition of Electron Shells: As you  move down a group, each successive element has an additional electron shell. For example, moving from lithium (Li) to sodium (Na) to potassium (K), each element has one more shell of electrons. This increase in the number of shells leads to a greater distance between the outermost electrons and the nucleus.

·  Shielding Effect: With the addition of electron shells, inner electrons shield the outermost electrons from the attractive force of the nucleus. This "shielding effect" reduces the effective nuclear charge that the outermost electrons experience, allowing them to be held less tightly. As a result, the outermost electrons can spread out more, contributing to an increase in atomic radius.

 

Overall, the combination of added electron shells and increased shielding leads to a larger atomic radius as you move down a group in the periodic table.

 

Q9. Explain why atomic radius decreases across the period?

Ans.: The atomic radius decreases across a period in the periodic table due to the following reasons:

·   Increasing Nuclear Charge: As you move from left to right across a period, the number of protons in the nucleus increases. This leads to a greater positive charge, which attracts the electrons more strongly.

·   Constant Shielding Effect: While the number of electrons also increases across the period, they are added to the same principal energy level (shell). Therefore, the shielding effect does not increase significantly since these additional electrons do not shield each other effectively from the nucleus.

·   Stronger Attraction: The  increased  nuclear  charge  and the  relatively constant  shielding result in a stronger attraction between the nucleus and the electrons. This pulls the electrons closer to the nucleus, leading to a decrease in atomic radius.

 

Overall, the combination of increasing nuclear charge and constant shielding causes the atomic radius to decrease as you move from left to right across a period.

 

Q10. Atomic radius of noble gases is comparatively larger than halogens. Why?

Ans.: The atomic radius of noble gases is comparatively larger than that of halogens when considering van der Waals radius due to the following reasons:

·   Complete Electron Shells: Noble gases have completely filled outer electron shells, resulting in a stable and larger electron cloud. This contributes to a larger van der Waals radius, which represents the distance at which the outer electron cloud can interact with neighbouring atoms.

·   Weak Intermolecular Forces: Noble gases primarily exist as monatomic gases and have weaker van der Waals forces compared to halogens. This means that their atoms can come closer together when forming van der Waals interactions, allowing their electron clouds to expand more without significant repulsion.

·   Size and Atomic Structure: The noble gases (e.g., neon, argon) have larger van der Waals radii because their filled electron shells create a more diffuse electron cloud. In contrast, halogens (e.g., fluorine, chlorine) have smaller van der Waals radii due to their higher effective nuclear charge and strong intermolecular forces, which pull their outer electrons closer to the nucleus.

Conclusion: The larger van der Waals radius of noble gases, compared to halogens, is primarily due to their stable, filled electron configurations, which allow for a larger and more diffuse electron cloud that results in greater atomic size.