CHAPTER 3 CLASSIFICATION OF ELEMENTS S PERIODICITY IN PROPERTIES

VERY SHORT QUESTIONS ANSWER

Q.1. State Mendeleev‘s periodic law?

Ans. "Elements' properties repeat periodically based on their atomic masses" - Mendeleev's periodic law.

Q.2.What is the basis of long form of periodic table?

Ans. Electronic configuration.

Q.3.What is mean by newland law of octaves?

Ans. Classification.

Q.4.What is meant by periodicity of properties?

Ans. Repetition.

Q.5.Why do elements with similar properties occur in the same group?

Ans. Valence electrons.

Q.6. State modern periodic law?

Ans. "Elements are arranged in increasing order of atomic number, and their properties show a periodic pattern."

Q.7. Define groups and periods?

Ans. Groups: Vertical columns in the periodic table representing elements with similar chemical properties and the same number of valence electrons.

Periods: Horizontal rows in the periodic table representing elements with sequentially increasing atomic numbers.

Q.8. How many groups and how many periods are there in long form of periodic table?

Ans. There are 18 groups and 7 periods in the long form of the periodic table.

Q.9. With which quantum number every period in periodic table begins?

Ans. Principal quantum number (n).

Q.10.What are s- block elements?

Ans. Alkali metals and alkaline earth metals.

Q.11.Give general electronic configuration of s-block elements?

Ans. ns^1 or ns^2 (where "n" represents the principal quantum number).

Q.12.What are p- block elements Give their general electronic configuration?

Ans. P-block elements are the elements in groups 13 to 18 of the periodic table. Their general electronic configuration is ns^2 np^1 to ns^2 np^6 (where "n" represents the principal quantum number).

Q.13.What are representative elements?

Ans. Representative elements, also known as main group elements, are the elements in groups 1, 2, and 13 to 18 of the periodic table. They exhibit a wide range of chemical properties and are often involved in chemical reactions due to the number of valence electrons in their outermost energy level.

Q.14.What are d-block elements why they are called transition metals?

Ans. Transition metals.

Q.15.Give general electronic configuration of d-block elements?

Ans. ns^2 (n-1)d^1 to ns^2 (n-1)d^10 (where "n" represents the principal quantum number).

Q.16.To which series man-made elements belong?

Ans. Actinide series.

Q.17.What is meant by lanthanides and Acitnides?

Ans. Lanthanides: A series of elements in the periodic table, from atomic number 57 (lanthanum) to 71 (lutetium). They are also known as rare earth elements.

Actinides: A series of elements in the periodic table, from atomic number 89 (actinium) to 103 (lawrencium). They are radioactive and have similar properties to actinium.

Q.18.What are inner transition metals why are they called rare earth metals?

Ans. Inner transition metals are the elements in the f-block of the periodic table, and they are called rare earth metals due to their historical difficulty in extraction, not necessarily their abundance.

Q.19.Give general electronic configuration of respective group why are they lest reactive?

Ans. Group 18 (Noble gases): ns^2 np^6; They are the least reactive due to their stable, fully filled valence electron shells, making them chemically inert.

Q.20.Which group elements are known as chalcogens?

Ans. Group 16 elements are known as chalcogens.

Q.21.Name two radioactive s-block elements?

Ans. Polonium and Francium.

Q.22.Name third chalcogen and fifth noble gas?

Ans. Third chalcogen: Sulfur

Fifth noble gas: Krypton

Q.23. Why do noble gases have bigger atomic size than halogens?

Ans. Noble gases have bigger atomic size than halogens because noble gases have a full outermost electron shell (octet), resulting in weaker attractive forces between the electrons and nucleus, leading to larger atomic radii.

Q.24.Define jonisation energy?

Ans. Ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms or ions to form one mole of positively charged ions.

Q.25.Why are lanthanides and actinides place at the bottom of the periodic table?

Ans. To fit them in without disrupting the periodic table's layout.

Q.26.What is the valence of the elements belonging to group 2,16?

Ans. Two

Q.27. Why are cations smaller than neutral atom?

Ans. Increased effective nuclear charge (stronger attraction between protons and electrons) leads to a smaller size due to electron removal in cations.

Q.28.Define electronegativity?

Ans. Electronegativity is the ability of an atom to attract and hold electrons in a chemical bond.

Q.29.Which among the following has the largest radius? Na, Mg2⁺ Al, k?

Ans. k.

Q.30.Out of Na and Mg which has higher second ionization energy?

Ans. Mg has a higher second ionization energy than Na.

SHORT QUESTIONS ANSWER

Q.1.What is periodic table how elements are classified in it?

Ans. The periodic table is a tabular arrangement of chemical elements based on their atomic number, electron configuration, and chemical properties. Elements in the periodic table are classified into periods (horizontal rows) and groups (vertical columns) based on their similar chemical properties and valence electron configurations.

Q.2.Which important property did Mendeleev use to classify the elements in his periodic table?

Ans. that's correct. Mendeleev used the atomic mass of elements as the key property to classify and arrange them in his periodic table. He noticed that when elements were arranged in order of increasing atomic mass, their chemical properties exhibited a periodic pattern, which led to the development of the modern periodic table.

Q.3. State the modern periodic law?

Ans. "The physical and chemical properties of elements are periodic functions of their atomic numbers."

Q.4.Why do different periods of the periodic table have different number of elements?

Ans. Different periods of the periodic table have different numbers of elements because each period corresponds to the different principal quantum numbers (n) of the elements' electron shells. As the value of n increases, more electron shells are added, accommodating more elements in each successive period.

Q.5.What is periodicity? what is its cause?

Ans. Periodicity refers to the regular repetition of certain properties or characteristics of elements in the periodic table. The cause of periodicity is the arrangement of electrons in atoms. The electronic configuration, particularly the number of valence electrons, plays a significant role in determining an element's chemical behavior and its position in the periodic table. Elements in the same group have similar outer electron configurations, leading to similar chemical properties and periodic patterns.

Q.6. 3rd period has 8 but not 18 elements why?

Ans. The 3rd period of the periodic table has 8 elements because it includes elements from sodium (Na) to argon (Ar). This period starts with 2 electrons in the 2nd energy level (n=2) and fills the 3rd energy level (n=3) up to 8 electrons. The 3rd energy level can accommodate a maximum of 18 electrons, but the 3rd period does not have enough elements to fill all 18 electron slots. Instead, it only includes elements with atomic numbers 11 to 18, making a total of 8 elements in the period.

Q.7. Define atomic radius why exact size of the atom cannot be determined?

Ans. Atomic radius is the distance from the nucleus of an atom to its outermost electron shell or the boundary of its electron cloud.

The exact size of an atom cannot be determined precisely due to the wave-like nature of electrons in quantum mechanics. In the quantum model, electrons are not fixed particles with definite positions, but rather exist as probability distributions or electron clouds. The position of an electron is described in terms of probabilities of finding it in certain regions around the nucleus. This inherent uncertainty in electron position makes it impossible to pinpoint the exact location of an electron and, consequently, the precise size of the atom. Instead, atomic radius is often estimated based on experimental data and theoretical models.

Q.8.What are isoelectronic ions?

Ans. Isoelectronic ions are ions that have the same number of electrons. These ions can belong to different elements, but they possess the same electron configuration. As a result, isoelectronic ions have similar chemical properties, despite being derived from different elements. For example, Na^+ (sodium ion), Mg^2+ (magnesium ion), and Al^3+ (aluminum ion) are isoelectronic since they all have the electron configuration of neon (1s^2 2s^2 2p^6).

Q.9.Name the different groups in s and p-block write their general configuration?

Ans. In the s-block, there are two groups:

Group 1 (Alkali Metals):

General electronic configuration: ns^1 (where "n" represents the principal quantum number).

Group 2 (Alkaline Earth Metals):

General electronic configuration: ns^2 (where "n" represents the principal quantum number).

In the p-block, there are six groups:

Group 13 (Boron Group):

General electronic configuration: ns^2 np^1 (where "n" represents the principal quantum number).

Group 14 (Carbon Group):

General electronic configuration: ns^2 np^2 (where "n" represents the principal quantum number).

Group 15 (Nitrogen Group):

General electronic configuration: ns^2 np^3 (where "n" represents the principal quantum number).

Group 16 (Chalcogens):

General electronic configuration: ns^2 np^4 (where "n" represents the principal quantum number).

Group 17 (Halogens):

General electronic configuration: ns^2 np^5 (where "n" represents the principal quantum number).

Group 18 (Noble Gases):

General electronic configuration: ns^2 np^6 (where "n" represents the principal quantum number).

Q.10.What is screening effect how does it affect the ionization enthalpies of the elements?

Ans. Screening effect, also known as shielding effect, refers to the phenomenon where inner electrons in an atom repel and shield the outer electrons from the full attractive force of the nucleus. The outer electrons experience a reduced effective nuclear charge due to the presence of inner electrons between them and the nucleus.

The screening effect affects the ionization enthalpies of elements in the following way:

Across a Period: As we move from left to right across a period, the number of protons (positive charge) in the nucleus increases, resulting in a stronger attractive force on the outer electrons. However, the screening effect remains relatively constant because the number of inner electrons also increases, balancing the effect. As a result, the ionization enthalpy generally increases across a period due to the increasing effective nuclear charge.

Down a Group: As we move down a group, the number of energy levels (shells) increases, and the outer electrons are farther away from the nucleus. The inner electrons are less effective in screening the outer electrons from the nucleus's positive charge. As a result, the screening effect decreases down a group, and the ionization enthalpy generally decreases due to the weaker attractive force experienced by the outer electrons.

In summary, the screening effect reduces the effective nuclear charge experienced by outer electrons, leading to lower ionization enthalpies down a group and higher ionization enthalpies across a period in the periodic table.

Q.11.Electron affinity of chlorine is more than fluorine why?

Ans. The electron affinity of chlorine is more than fluorine because chlorine has a higher effective nuclear charge. The effective nuclear charge is the net positive charge experienced by an electron in an atom after considering the shielding effect of inner electrons. As we move from left to right across a period in the periodic table, the atomic number (number of protons) increases, resulting in a higher effective nuclear charge.

Chlorine (Cl) is in the 3rd period, and fluorine (F) is in the 2nd period. Since they are in the same group (Group 17 or Halogens), they have the same number of valence electrons. However, chlorine has an additional energy level (shell) compared to fluorine, and its valence electrons are farther from the nucleus. As a result, the attraction between the valence electrons and the nucleus is weaker in chlorine, leading to a higher electron affinity as it can more easily accept an additional electron to achieve a stable electron configuration.

Therefore, chlorine has a higher electron affinity than fluorine due to its higher effective nuclear charge and larger atomic size.

Q.12.Beryllium and magnesium atoms do not impart colour to flame whereas alkaline earth metals do so why?

Ans. Beryllium and magnesium atoms do not impart color to the flame because they have completely filled valence electron shells. When these atoms are heated in a flame, their electrons are excited to higher energy levels, and as they return to their ground state, they release energy in the form of light. However, since beryllium and magnesium have full valence electron shells (ns^2 np^6 configuration), their electrons do not undergo any electronic transitions, and they do not emit any visible light.

On the other hand, alkaline earth metals, such as calcium, strontium, and barium, have partially filled valence electron shells. When these metals are heated in a flame, their valence electrons can be excited to higher energy levels and then emit characteristic colors of light as they return to their ground state. The specific colors produced are due to the electronic transitions within their partially filled electron configurations.

In summary, the lack of visible color in the flame of beryllium and magnesium is due to their fully filled valence electron shells, which prevents them from undergoing electronic transitions and emitting visible light. Alkaline earth metals, with partially filled valence electron shells, emit characteristic colors in a flame due to the electronic transitions of their valence electrons.

Q.13.Why inert gases have higher ionization enthalpy but lower electron gain enthalpy than halogens?

Ans. Inert gases (noble gases) have higher ionization enthalpy but lower electron gain enthalpy than halogens due to their electron configurations.

Ionization Enthalpy: Inert gases have completely filled valence electron shells (ns^2 np^6 configuration), making them very stable. To remove an electron from a noble gas atom, a significant amount of energy is required, resulting in higher ionization enthalpies. The removal of an electron from a noble gas would lead to an unstable configuration, which is energetically unfavorable.

Electron Gain Enthalpy: Halogens, such as fluorine and chlorine, have one less electron in their valence electron shell, making them highly reactive and eager to gain an additional electron to achieve a stable, fully filled valence shell (ns^2 np^6 configuration). When a halogen gains an electron, it attains a stable electron configuration, resulting in a release of energy and lower electron gain enthalpy.

In summary, inert gases have higher ionization enthalpy due to the stability of their fully filled valence shells, making it difficult to remove an electron. On the other hand, halogens have lower electron gain enthalpy due to their high reactivity and the strong tendency to gain an electron to achieve a stable electron configuration.

Q.14.What is newland law of octaves?

Ans. The Newlands' law of octaves, proposed by John Newlands in 1865, was an early attempt to organize the known elements into a periodic table. According to this law, when elements are arranged in order of increasing atomic masses, every eighth element displays similar properties to the first element, much like musical notes that repeat every octave in music. However, Newlands' law had limitations and could not accommodate all known elements, leading to its eventual replacement by Mendeleev's more successful periodic table based on atomic number.

Q.15.Why do halogens have high electron gain enthalpies?

Ans. Halogens have high electron gain enthalpies because of their electron configuration and the desire to achieve a stable, fully filled valence electron shell.

Halogens belong to Group 17 of the periodic table, and they have seven valence electrons in their outermost energy level (ns^2 np^5 configuration). To achieve a stable electron configuration like the noble gases (ns^2 np^6 configuration), halogens need to gain one more electron. Since they are only one electron away from achieving this stable configuration, they have a strong tendency to attract an additional electron.

When a halogen gains an electron, it forms a negatively charged ion (anion), and this process releases energy. The energy released is the electron gain enthalpy. Due to their high electronegativity and the relatively small size of their outermost electron shell, halogens have a strong ability to attract and capture an extra electron, resulting in high electron gain enthalpies. This high electron affinity makes halogens highly reactive and readily forms negatively charged ions in chemical reactions.

Q.16.How ionization enthalpy or ionization energy vary along period and group?

Ans. Ionization enthalpy, or ionization energy, refers to the energy required to remove one mole of electrons from one mole of gaseous atoms or ions to form one mole of positively charged ions.

Along a Period: Ionization enthalpy generally increases as you move from left to right across a period in the periodic table. This is because the number of protons in the nucleus increases from left to right, resulting in a stronger attractive force between the nucleus and the electrons. As a result, it becomes more difficult to remove an electron, requiring more energy, and leading to higher ionization enthalpies.

Along a Group: Ionization enthalpy generally decreases as you move down a group in the periodic table. This is due to the increase in the number of energy levels (shells) as you move down a group. The outermost electrons are farther from the nucleus and are shielded by the inner electrons, resulting in a weaker effective nuclear charge experienced by the outermost electron. As a result, the outermost electron is more easily removed, requiring less energy, and leading to lower ionization enthalpies.

In summary, ionization enthalpy increases across a period due to the increasing effective nuclear charge, and it decreases down a group due to the increasing atomic size and weaker effective nuclear charge experienced by the outermost electrons.

Q.17.Why is the first I.E. of transition elements almost same?

Ans. The first ionization energy (I.E.) of transition elements is almost the same because of their similar electron configurations.

Transition elements are located in the d-block of the periodic table, and they have partially filled d-orbitals in their electron configuration. The first ionization energy refers to the energy required to remove the first electron from a neutral atom to form a positively charged ion.

Since the transition elements have similar electron configurations with partially filled d-orbitals, the removal of the first electron involves breaking a weakly filled subshell. As a result, the energy required to remove the first electron is relatively similar among the transition elements. The variation in their first ionization energies is not as significant as in elements with completely filled or empty valence electron shells.

It's important to note that the first ionization energy of transition elements may still show slight variations due to factors like atomic size, effective nuclear charge, and electron shielding. However, compared to the main group elements, the first ionization energy of transition elements is relatively more uniform.

Q.18. Why sodium ion is smaller than sodium atom while fluoride ion is bigger than fluorine atom?

Ans. The size of an ion compared to its neutral atom depends on the gain or loss of electrons during the ionization process.

Sodium Ion (Na+): When sodium (Na) loses one electron to become a sodium ion (Na+), it forms a cation. The loss of an electron reduces the electron-electron repulsion in the electron cloud, making the remaining electrons more strongly attracted to the nucleus. This results in a decrease in the electron cloud's size and a smaller ionic radius compared to the neutral sodium atom.

Fluoride Ion (F-): When fluorine (F) gains one electron to become a fluoride ion (F-), it forms an anion. The addition of an extra electron increases the electron-electron repulsion in the electron cloud, causing the electron cloud to expand. This leads to a larger ionic radius compared to the neutral fluorine atom.

In summary, the sodium ion (Na+) is smaller than the sodium atom (Na) because of the loss of one electron, reducing the electron cloud's size, while the fluoride ion (F-) is bigger than the fluorine atom (F) due to the gain of one electron, causing the electron cloud to expand.

Q.19.What do you know about diagonal relationship?

Ans. The diagonal relationship is a unique similarity observed between certain pairs of elements in the periodic table, despite their apparent differences in group and period. The elements that exhibit diagonal relationships are located diagonally across the periodic table from each other.

The most well-known example of a diagonal relationship is between beryllium (Be) and aluminum (Al) in Group 2 and Group 13, respectively. Beryllium, an alkaline earth metal in Group 2, shares many similar properties with aluminum, a post-transition metal in Group 13. Some of the similarities between these elements include:

Similar atomic and ionic radii.

The ability to form covalent compounds with similar ligands.

Similar electronegativity values.

Formation of amphoteric oxides (capable of acting as both acidic and basic oxides).

The diagonal relationship is attributed to the comparable charge/radius ratios between the elements, which results in similar bonding characteristics and chemical behavior. This phenomenon is also observed in other pairs of elements, such as lithium (Li) and magnesium (Mg), as well as boron (B) and silicon (Si).

The diagonal relationship plays a significant role in understanding the chemical properties of certain elements and provides valuable insights into the periodic trends in the periodic table.

Q.20.What is the relationship between the first ionisation enthalpies and metallic and non –metallic properties?

Ans. The first ionization enthalpy is the energy required to remove one mole of electrons from one mole of neutral atoms in the gaseous state to form positively charged ions. The metallic and non-metallic properties of elements are closely related to their first ionization enthalpies.

Metallic Properties:

Metallic elements are found on the left side of the periodic table, typically in the s-block and partially in the d-block.

They have low first ionization enthalpies, meaning it requires relatively less energy to remove an electron from their outermost shell.

This low ionization energy allows metallic elements to lose electrons easily and form positively charged ions (cations).

As a result, metallic elements tend to be good conductors of electricity and heat, have high malleability and ductility, and exhibit metallic luster.

Non-Metallic Properties:

Non-metallic elements are primarily found on the right side of the periodic table, including the p-block and some in the d-block.

They have relatively high first ionization enthalpies, meaning it requires a significant amount of energy to remove an electron from their outermost shell.

Non-metals prefer to gain electrons and form negatively charged ions (anions) when they react chemically.

Non-metals generally have poor electrical and thermal conductivity and lack the characteristic luster of metals.

They often exist as gases, liquids, or brittle solids and have diverse properties like being insulators or semiconductors.

In summary, the relationship between the first ionization enthalpies and metallic and non-metallic properties is that elements with low ionization energies tend to exhibit metallic properties, while elements with high ionization energies exhibit non-metallic properties. The trend of ionization enthalpies across the periodic table reflects the pattern of metallic and non-metallic character of elements.

Q.21. Differentiate between ionization enthalpy and electron gain enthalpy?

Ans. Ionization enthalpy and electron gain enthalpy are both related to the energy changes that occur during the addition or removal of electrons from atoms. However, they represent different processes and have opposite signs. Let's differentiate between them:

Ionization Enthalpy:

Ionization enthalpy, also known as ionization energy, is the energy required to remove one mole of electrons from one mole of neutral gaseous atoms to form positively charged ions (cations).

It is typically represented by the equation: X(g) → X⁺(g) + e⁻, where X represents the atom, X⁺ represents the cation, and e⁻ is the removed electron.

Ionization enthalpy is always endothermic since energy is required to overcome the attractive forces between the positively charged nucleus and the negatively charged electrons.

The ionization enthalpy generally increases across periods (rows) of the periodic table from left to right due to increased effective nuclear charge, which leads to a stronger hold on the outermost electrons.

It decreases down a group (column) of the periodic table because the outermost electrons are farther from the nucleus, reducing the effective nuclear charge.

Electron Gain Enthalpy:

Electron gain enthalpy, also known as electron affinity, is the energy change that occurs when one mole of electrons is added to one mole of neutral gaseous atoms to form negatively charged ions (anions).

It is typically represented by the equation: X(g) + e⁻ → X⁻(g), where X represents the atom, X⁻ represents the anion, and e⁻ is the added electron.

Electron gain enthalpy can be either exothermic or endothermic, depending on whether energy is released or absorbed during the process.

The electron gain enthalpy generally becomes more exothermic (more negative) across periods of the periodic table from left to right. This is because the effective nuclear charge increases, making it more favorable for atoms to accept an electron and achieve a stable electronic configuration.

It tends to become less exothermic (less negative) down a group because the atomic size increases, leading to a weaker attraction between the nucleus and the incoming electron.

In summary, ionization enthalpy is the energy required to remove electrons from an atom to form cations, while electron gain enthalpy is the energy change when electrons are added to an atom to form anions. Ionization enthalpy is always endothermic, while electron gain enthalpy can be either exothermic or endothermic depending on the atom's properties.

Q.22.What do you mean by valeney? How does it vary along period and group?

Ans. I assume you meant "valeney" instead of "valeney." Valeney refers to the combining capacity or the number of electrons that an atom of an element can gain, lose, or share to achieve a stable electron configuration and form chemical bonds with other atoms. It determines how an element interacts with other elements to form compounds.

The valeney of an element is usually determined by the number of electrons in its outermost energy level, known as the valence electrons. The valence electrons are the electrons in the outermost shell (also called valence shell) of an atom. For most main group elements (s-block and p-block elements), the valency corresponds to the number of electrons needed to achieve a full outer shell (octet) or to have an empty outer shell (duet in the case of hydrogen and helium).

Variation of Valeney Along a Period (Horizontal Row):

As you move from left to right along a period in the periodic table, the number of valence electrons increases by one with each element.

The valeney may vary within a period, but it often starts with one and increases up to a maximum of eight (except for hydrogen and helium, which have valencies of 1 and 2, respectively).

Elements on the left side of the periodic table (Group 1, alkali metals) have a valeney of +1 since they tend to lose one electron to achieve a stable electron configuration.

Elements on the right side of the periodic table (Group 17, halogens) have a valeney of -1 since they tend to gain one electron to achieve a stable electron configuration.

Elements in the middle of the periodic table (transition metals) may exhibit multiple valences because they can lose different numbers of electrons to form different charged ions.

Variation of Valeney along a Group (Vertical Column):

Within a group in the periodic table, the number of valence electrons remains constant as all elements in the group have the same outer electron configuration.

Elements in the same group tend to have similar chemical properties due to the same valeney.

For example, elements in Group 1 (alkali metals) have one valence electron and a valeney of +1, while elements in Group 17 (halogens) have seven valence electrons and a valeney of -1.

In summary, valeney is the number of electrons that an element can gain, lose, or share to form chemical bonds and achieve a stable electron configuration. It varies along a period as the number of valence electrons increases, and it remains the same within a group as all elements in the group have the same outer electron configuration.

Q.23.How are lithium and magnesium related diagonally?

Ans. Lithium and magnesium are related diagonally in the periodic table as part of a diagonal relationship between certain elements. The diagonal relationship is observed between elements that are found diagonally to each other in the periodic table, crossing between s-block and p-block elements. In this case, lithium (Li) is found diagonally above magnesium (Mg). This relationship results in some similarities in their properties.

The diagonal relationship between lithium and magnesium is primarily due to their similar atomic size and charge-to-size ratio of their ions:

Atomic Size:

Atomic size refers to the size of the atom, specifically the atomic radius.

As you move diagonally from lithium to magnesium, there is a gradual increase in atomic size despite moving from left to right in the periodic table. This is because, as you move down a group, atomic size tends to increase due to the addition of new energy levels (shells).

The atomic size of lithium and magnesium is relatively close, making them similar in this aspect.

Charge-to-Size Ratio of Ions:

Both lithium and magnesium readily form cations (positively charged ions) by losing electrons. Lithium forms Li⁺ ions, and magnesium forms Mg²⁺ ions.

Due to their similar atomic sizes and the removal of one or two electrons, respectively, the charge-to-size ratio of their cations is quite similar.

This charge-to-size ratio is important for determining certain chemical properties, such as their solubilities and behaviors in ionic compounds.

The diagonal relationship between lithium and magnesium leads to some similarities in their chemical properties:

Similar Solubilities:

Both lithium and magnesium salts are relatively soluble in water due to the similarities in their charge-to-size ratio.

For example, lithium carbonate (Li₂CO₃) and magnesium carbonate (MgCO₃) are both soluble in water.

Similar Hydroxides:

Both lithium hydroxide (LiOH) and magnesium hydroxide (Mg(OH)₂) are weak bases and sparingly soluble in water.

Their solubilities are again influenced by the similarities in their ionic sizes and charges.

Formation of Complexes:

Both lithium and magnesium can form coordination complexes due to their relatively high charge-to-size ratio. These complexes are common in various chemical reactions.

It's important to note that the diagonal relationship is not perfect, and there are still distinct differences between lithium and magnesium in many aspects. However, the diagonal relationship helps to explain some of the observed similarities in their properties despite their different positions in the periodic table.

Q.24.What are bridge elements Explain?

Ans. I apologize for the confusion. It seems I provided the same answer as in the previous response. The term "bridge elements" does not have a specific meaning or significance in the context of chemistry. There might be a misunderstanding or miscommunication.

If you are referring to another concept or term related to chemistry, please provide more context or clarify your question, and I'd be happy to help you with the correct information.

Q.25.What are the various defects in the Mendeleev’s periodic table?

Ans. In Mendeleev's original periodic table, which he proposed in 1869, there were certain defects or shortcomings due to the limited knowledge of elements and atomic structure at that time. Some of the main defects in Mendeleev's periodic table are as follows:

Position of Hydrogen:

Mendeleev placed hydrogen in Group 1 (alkali metals) of his periodic table, despite it having properties that are quite different from the alkali metals.

Hydrogen is a unique element that can exhibit both metallic and non-metallic properties, and it can form hydrides similar to the alkali metals but also behaves like a non-metal in other compounds.

The position of hydrogen has been a topic of debate and discussion over the years, and modern periodic tables place it separately above Group 1.

Incomplete Filling of d-Block:

Mendeleev's periodic table did not include the d-block (transition metals) as a separate part of the table.

Some elements in the d-block were placed in inappropriate groups, and there was no clear explanation for the incomplete filling of the d-block elements.

In modern periodic tables, the d-block is recognized as a distinct section, and the transition metals are appropriately placed.

Anomalous Pairs of Elements:

In certain cases, elements with higher atomic masses were placed before elements with lower atomic masses, leading to anomalous pairs in Mendeleev's periodic table.

For example, iodine (I, atomic mass ~ 127) was placed before tellurium (Te, atomic mass ~ 128) and cobalt (Co, atomic mass ~ 58.9) was placed before nickel (Ni, atomic mass ~ 58.7).

These anomalies were later resolved when the concept of atomic number (number of protons) was established as the basis for arranging elements.

Position of Isotopes:

Mendeleev's periodic table did not account for isotopes, which are atoms of the same element with different numbers of neutrons.

For example, isotopes of chlorine (35Cl and 37Cl) were not differentiated, and they were both placed in the same position in the table.

Modern periodic tables account for isotopes and consider them as variations of the same element.

Despite these defects, Mendeleev's periodic table laid the foundation for the modern periodic table. His work provided a systematic approach to organizing elements based on their properties and atomic masses, and it eventually led to the development of the periodic table we use today, which is based on the atomic number and electron configuration of elements.