Answer. Na2O oxide is basic in nature because it reacts with dilute acids to form salt and water. Na2O also reacts with water to form NaOH which is alkaline in nature as it can release OH- in the solution.
On the other hand, P2O5 is acidic in nature because it reacts with water to form H3PO3, which is acidic in nature as it can release H+ in the solution.
Question 2. How acidic basic amphoteric behavior of oxides is explained?
Answer. Generally, acidic or basic or amphoteric behavior of oxides is determined through the electropositive character of the oxide's central atom. If the central atom is more electropositive, the oxide is more basic . On the hand, if the central atom is more electronegative, the oxide is more acidic. Electropositive character decreases from left to right in periodic table and increases down the column. The basic behavior decreases while acidic character increases from left to right , via an amphoteric oxide (aluminum oxide) in the middle. An amphoteric oxide is one that shows both acidic and basic properties.
Question 3. Why the elements of group IA are called alkali metals?
Answer. The group IA of the periodic table contains six elements (i. e. Li, Na, K, Rb, Cs and Fr).These elements are called alkali metals because they form alkalis (i.e. strong bases) when they react with water.
M = Li, Na, K, Rb, Cs and Fr
Question 4. Why all group IA metals have low ionization energies?
Answer. Ionization energy is the minimum energy required to remove valence electron from isolated gaseous atom in its ground state. The magnitude of ionization energy is a measure of how "tightly" the electron is held in the atom. Group IA elements have low ionization energies because of larger atomic sizes and less effective nuclear charge in their respective periods. Furthermore, they have ability to attain a noble gas configuration by loosing just single electron.
Answer. The group IA of the periodic table contains six elements (i. e. Li, Na, K, Rb, Cs and Fr).These elements are called alkali metals because they form alkalis (i.e. strong bases) when they react with water.
M = Li, Na, K, Rb, Cs and Fr
Question 4. Why all group IA metals have low ionization energies?
Answer. Ionization energy is the minimum energy required to remove valence electron from isolated gaseous atom in its ground state. The magnitude of ionization energy is a measure of how "tightly" the electron is held in the atom. Group IA elements have low ionization energies because of larger atomic sizes and less effective nuclear charge in their respective periods. Furthermore, they have ability to attain a noble gas configuration by loosing just single electron.
Question 5. Why do Group IA metals show strong electropositive character?
Answer.
Electropositive character of metals depend upon their tendency to lose electrons. In any period, group IA metals have relatively higher tendency to lose electrons due to
greater atomic sizes and less effective nuclear charge. Therefore, they have low ionization energies and show more electropositive character.
Question 6. Why do group 1st metal show strong reducing properties?
Answer. Reducing properties of metals also depend upon their tendency to lose electrons. In each period, group IA metals have relatively higher tendency to lose electrons due to greater atomic sizes and less effective nuclear charge. Therefore, they have low ionization energies and show more reducing properties.
Answer.
Electropositive character of metals depend upon their tendency to lose electrons. In any period, group IA metals have relatively higher tendency to lose electrons due to
greater atomic sizes and less effective nuclear charge. Therefore, they have low ionization energies and show more electropositive character.
Question 6. Why do group 1st metal show strong reducing properties?
Answer. Reducing properties of metals also depend upon their tendency to lose electrons. In each period, group IA metals have relatively higher tendency to lose electrons due to greater atomic sizes and less effective nuclear charge. Therefore, they have low ionization energies and show more reducing properties.
Question 7:Why different color are imparted by atoms of the group IA first a metals to the flame?
Answer.
The ionization enthalpy of alkali metals is very low. The flame has enough energy to excite outer ns1 electron to higher energy level. When this excited electron is de-excited to ground state, absorbed energy is emitted in the visible region of electromagnetic spectrum. This is why, alkali metals imparts color to flame. Different colors of flame are imparted due to different energies absorbed by different atoms of alkali metals.
Question 8: Why are elements of IIA group called alkaline earth metals?
Answer.
In order to understand term "alkaline earth metals", it can be defined word by word.
Elements of group IIA are called alkaline, because they form alkaline solutions (hydroxides) when they react with water; earth, because their oxides are insoluble in water and resistant to heat; metals, because they can conduct electricity and have shiny luster.
The ionization enthalpy of alkali metals is very low. The flame has enough energy to excite outer ns1 electron to higher energy level. When this excited electron is de-excited to ground state, absorbed energy is emitted in the visible region of electromagnetic spectrum. This is why, alkali metals imparts color to flame. Different colors of flame are imparted due to different energies absorbed by different atoms of alkali metals.
Question 8: Why are elements of IIA group called alkaline earth metals?
Answer.
In order to understand term "alkaline earth metals", it can be defined word by word.
Elements of group IIA are called alkaline, because they form alkaline solutions (hydroxides) when they react with water; earth, because their oxides are insoluble in water and resistant to heat; metals, because they can conduct electricity and have shiny luster.
Question 9. Why do alkaline earth metals have high melting and boiling points than alkali metals?
Answer.
Answer.
Alkali metals (Li, Na, K, Rb, Cs and Cs) are soft and have low melting and boiling points because they have only one electron in valence shell (ns1) and their atomic sizes are relatively large in each period. So, binding energy of the atoms in the crystal lattice of the metal is low. Therefore, the metallic bonds in these metals are not very strong.
On the other hand, alkaline earth metals (Be, Mg, Ca, Sr, Ba) are hard and have high melting and boiling points when compared with Alkali metals. This is due to fact that, they have two electrons in valence shell (ns2) and their atomic sizes are relatively small. So, binding energy of the atoms in the crystal lattice of the metal is high Therefore, the metallic bonds in these metals are very strong.
Question 10. How do group I metals resemble with group 2 metals?
Answer
Following points are similar between group 1 metals and Group 2 metals.
i. Metals of groups 1 and 2 are s-block elements. Both have valence electrons in s-orbital.
ii. Metals of both groups don't occur in free state
iii. Metals of both groups are highly electropositive.
iv. Metals of both groups form basic hydroxides.
v. Metals of both groups give characteristics color to flame.
On the other hand, alkaline earth metals (Be, Mg, Ca, Sr, Ba) are hard and have high melting and boiling points when compared with Alkali metals. This is due to fact that, they have two electrons in valence shell (ns2) and their atomic sizes are relatively small. So, binding energy of the atoms in the crystal lattice of the metal is high Therefore, the metallic bonds in these metals are very strong.
Question 10. How do group I metals resemble with group 2 metals?
Answer
Following points are similar between group 1 metals and Group 2 metals.
i. Metals of groups 1 and 2 are s-block elements. Both have valence electrons in s-orbital.
ii. Metals of both groups don't occur in free state
iii. Metals of both groups are highly electropositive.
iv. Metals of both groups form basic hydroxides.
v. Metals of both groups give characteristics color to flame.
Question 11. How do group 1 metals differ from group 2 metals?
Answer.
i. Electronic Configuration
Group 1 metals have one electron in their valence s-orbital while group 2 metals have two electrons in valence s-orbital.
ii. Melting and boiling points
Group 1 metals have low melting and boiling points than group 2 metals due to relative larger atomic size and only one valence electron.
iii. Ionization energy
Group 1 metals have relatively low ionization energies than group 2 metals.
This is because group 1 metals have greater atomic size and more effective nuclear charge than group 2 metals.
iv. Basicity of oxides
Oxides of group 1 metals are more basic than oxides of group 2 metals. This is because group 1 metals have more metallic character (low ionization energies) than group 2 metals.
v. Solubility of hydroxides
Lattice enthalpies of hydroxides group 1 metals are lower than group 2 metals due to greater atomic size and low nuclear charge. Hence, hydroxide of group 1 metals show more solubility in water than group 2 metals.
Answer.
i. Electronic Configuration
Group 1 metals have one electron in their valence s-orbital while group 2 metals have two electrons in valence s-orbital.
ii. Melting and boiling points
Group 1 metals have low melting and boiling points than group 2 metals due to relative larger atomic size and only one valence electron.
iii. Ionization energy
Group 1 metals have relatively low ionization energies than group 2 metals.
This is because group 1 metals have greater atomic size and more effective nuclear charge than group 2 metals.
iv. Basicity of oxides
Oxides of group 1 metals are more basic than oxides of group 2 metals. This is because group 1 metals have more metallic character (low ionization energies) than group 2 metals.
v. Solubility of hydroxides
Lattice enthalpies of hydroxides group 1 metals are lower than group 2 metals due to greater atomic size and low nuclear charge. Hence, hydroxide of group 1 metals show more solubility in water than group 2 metals.
Question 12. Discuss the metallic and non-metallic character of group four elements.
Answer
The metallic character of elements depend upon their ability to lose their valence electron. Metallic character increases as you move down the group in periodic table because of easier to lose valence electron due to larger atomic size.
Carbon, Silicon and grey Tin
Carbon, silicon and grey tin have four valence electrons and make four covalent bonds. These elements do not lose valence electrons to form cation and don't show metallic character.
Germanium, White Tin, and lead
Germanium, White Tin, and lead make metallic bonds. They can lose their valence electrons to form metal ions and show metallic character.
Answer
The metallic character of elements depend upon their ability to lose their valence electron. Metallic character increases as you move down the group in periodic table because of easier to lose valence electron due to larger atomic size.
Carbon, Silicon and grey Tin
Carbon, silicon and grey tin have four valence electrons and make four covalent bonds. These elements do not lose valence electrons to form cation and don't show metallic character.
Germanium, White Tin, and lead
Germanium, White Tin, and lead make metallic bonds. They can lose their valence electrons to form metal ions and show metallic character.
Question 13. Discuss the general group Trends of group 7 elements.
Answer.
i. Physical Properties
Melting points and boiling points increase down the group due to increase in atomic size and shielding effect and decrease in effective nuclear charge.
ii. Electron affinity and Electronegativity
Electron affinity and electronegativity
gradually decreases down the group due to gradual increase in atomic size of elements.
iii. Bond energy
Bond energies of Cl-Cl, Br-Br and I-I bond decrease down the group due to increase in atomic size except F-F bond.
IV. Oxidizing power
Oxidizing power decreases decreases down the group due decrease in electron affinity.
Bond energies of Cl-Cl, Br-Br and I-I bond decrease down the group due to increase in atomic size except F-F bond.
IV. Oxidizing power
Oxidizing power decreases decreases down the group due decrease in electron affinity.
Next
Question 14. Why the term halogen is used for group 7 elements?
Answer. The term "halogens" means "salt-producing". As group 7 elements form salts when react with different metals, so they are called halogens. They can form various salts including CaCl2, MgCl2, NaCl (Table salt) NaF etc
Question. 15: why does fluorine differ from other members of its group?
Answer.
Fluorine different from other members of its group due to following reasons.
1. Fluorine is more electronegative because its atomic size is very small and effective nuclear charge is higher as compared to other halogens.
2. Fluorine has higher oxidizing power as compared to other halogens. This is because of
i. Lower bond enthalpy of F2 as compared to Cl2 and Br2 due to greater electron-electron repulsion of lone pairs in small sized F2 molecule.
ii. Higher heat of hydration of F2 due to its higher charge density.
iii. Although, fluorine has less electron affinity value as compared to chlorine, but due to above two factor, still F2 has higher oxidizing power.
Answer. The term "halogens" means "salt-producing". As group 7 elements form salts when react with different metals, so they are called halogens. They can form various salts including CaCl2, MgCl2, NaCl (Table salt) NaF etc
Question. 15: why does fluorine differ from other members of its group?
Answer.
Fluorine different from other members of its group due to following reasons.
1. Fluorine is more electronegative because its atomic size is very small and effective nuclear charge is higher as compared to other halogens.
2. Fluorine has higher oxidizing power as compared to other halogens. This is because of
i. Lower bond enthalpy of F2 as compared to Cl2 and Br2 due to greater electron-electron repulsion of lone pairs in small sized F2 molecule.
ii. Higher heat of hydration of F2 due to its higher charge density.
iii. Although, fluorine has less electron affinity value as compared to chlorine, but due to above two factor, still F2 has higher oxidizing power.
Question 16: What are structures of CO2 and SiO2 and when they differ?
AnswerCarbon dioxide is a simple molecule while silicon dioxide is a giant molecule.
In carbon dioxide, there exist weak Ven der wall forces or London dispersion forces between molecules. These Van der Waal forces can be broken with little energy. That is why, CO2 melts just above - 56.6 oC. Thus, CO2 exists as a gas at room temperature.
In silicon dioxide, each Si is bonded with four oxygen atoms through covalent bonds in tetrahedral shape and each oxygen is bonded with two silicon atoms in bent shape. This forms a a giant (macromolecular) structure, which makes it incredibly strong. Thus, SiO2 is a hard sold at room temperature.
Question 17. Explain why nitrates and carbonates of lithium are not stable?
Answer. The nitrates and carbonates of alkali metals are stable except lithium.
Lithium ion has small size and hence, high charge density. So, it has tendency to attract the negative charge of carbonate or nitrate ions towards itself and increase covalent character. Due to covalent character of Li2CO3 and LiNO3 are unstable.
Furthermore, due to close packing of small sized Li+ and O2-, lattice energy of Li2O is significant, so it is highly stable as compared to Li2CO3 and LiNO3. Thus it is readily formed.
Question 17. Explain why nitrates and carbonates of lithium are not stable?
Answer. The nitrates and carbonates of alkali metals are stable except lithium.
Lithium ion has small size and hence, high charge density. So, it has tendency to attract the negative charge of carbonate or nitrate ions towards itself and increase covalent character. Due to covalent character of Li2CO3 and LiNO3 are unstable.
Furthermore, due to close packing of small sized Li+ and O2-, lattice energy of Li2O is significant, so it is highly stable as compared to Li2CO3 and LiNO3. Thus it is readily formed.
Question 18. Differentiate the behavior of Li and Na with atmospheric oxygen.
Answer. Although both lithium and sodium are present in the same group of periodic table, but lithium forms normal oxide which contains O2- and Na forms superoxide which contains complicated [O-O]2-.
Answer. Although both lithium and sodium are present in the same group of periodic table, but lithium forms normal oxide which contains O2- and Na forms superoxide which contains complicated [O-O]2-.
Na ion with large atomic size and small charge density forms more complicated oxide. However, at the top of group, the small sized Li+ has high charge density and has greater tendency to polarize the more complicated ions to form normal oxide.
Question 19. Alkali metal carbonates are more soluble than alkaline earth metal carbonate. why?
Answer.
Alkali metal carbonates are more soluble because of lesser nuclear charge and larger size and hence more ionic character.
On the other hand, alkaline earth metals are less soluble because of increased nuclear charge and smaller size which increase the covalent character and decrease ionic character.
Both lattice enthalpy and hydration energy are inversely related to the size of ions.
For alkaline metal carbonates, solubility decreases down the group as decrease in hydration energy is more than decrease in lattice enthalpy.
Question 21. Oxidizing power of F2 is greater than oxidizing power of I2. why?
Answer.
Fluorine F2 has higher oxidizing power than I2 because it can be reduced easily as shown in three steps.
i. First step involves breaking of F-F bond to form F(g) that requires energy equivalent to its bond enthalpy. Bond enthalpy of F2 is lesser as compared to Cl2 and Br2 due to greater electron-electron repulsion of lone pairs in small sized F2 molecule.
ii. Next, F(g) is converted to F(g)-. During this process energy released is also known as enthalpy of electron affinity (E. A.). E. A. of F is higher due to its higher charge density.
iii. Now, F- will be hydrated to form F-(aq.) is and energy is released known hydration enthalpy. Hydration enthalpy is also very high due to small size of F- and high charge density.
Although, fluorine has less bond enthalpy as compared to Iodine, but due to high E. A. and hydration enthalpy, still F2 has higher oxidizing power.
Question 19. Alkali metal carbonates are more soluble than alkaline earth metal carbonate. why?
Answer.
Alkali metal carbonates are more soluble because of lesser nuclear charge and larger size and hence more ionic character.
On the other hand, alkaline earth metals are less soluble because of increased nuclear charge and smaller size which increase the covalent character and decrease ionic character.
Therefore, due to less ionic and more covalent nature (Fajan'rule of covalency) alkaline metal carbonates are less soluble in water.
Question 20. Explain why solubility of alkaline earth metal carbonates decreases down the group?
Answer. Solubility depends upon lattice enthalpy and hydration energy. A compound can only be soluble when hydration energy is equal or greater than lattice enthalpy.Both lattice enthalpy and hydration energy are inversely related to the size of ions.
For alkaline metal carbonates, solubility decreases down the group as decrease in hydration energy is more than decrease in lattice enthalpy.
Question 21. Oxidizing power of F2 is greater than oxidizing power of I2. why?
Answer.
Fluorine F2 has higher oxidizing power than I2 because it can be reduced easily as shown in three steps.
i. First step involves breaking of F-F bond to form F(g) that requires energy equivalent to its bond enthalpy. Bond enthalpy of F2 is lesser as compared to Cl2 and Br2 due to greater electron-electron repulsion of lone pairs in small sized F2 molecule.
ii. Next, F(g) is converted to F(g)-. During this process energy released is also known as enthalpy of electron affinity (E. A.). E. A. of F is higher due to its higher charge density.
iii. Now, F- will be hydrated to form F-(aq.) is and energy is released known hydration enthalpy. Hydration enthalpy is also very high due to small size of F- and high charge density.
Although, fluorine has less bond enthalpy as compared to Iodine, but due to high E. A. and hydration enthalpy, still F2 has higher oxidizing power.
Question 22. HF is a weaker acid than HI. Why?
Answer.
HF is a weaker acid than HI. This is due to following reasons.
1. Presence of strong hydrogen bonding in HF which decreases its tendency to donate proton, while in HI, there is no hydrogen bonding. .
2 Bond enthalpy of H-F is considerably higher due to the highest electronegativity of F atom, while bond enthalpy of HI is lower due to lower electronegativity of I atom.
3. Conjugate base of HF is highly unstable and readily combine with proton. On the other hand, conjugate base of HI is highly stable.
Question 23. On what factors does the oxidizing power of halogens depend?
Answer.
Oxidizing power of halogens depends upon following factors.
1. Bond enthalpy of X-X
2. Electron affinity of X
3. Hydration energy of halide ion (X-).
1. Bond enthalpy of X-X
Greater the X-X bond enthalpy, lower will be oxidizing power and vice versa.
2. Electron Affinity
Greater the electron affinity of halogen, higher will be oxidizing power and vice versa.
3. Hydration enthalpy
Greater the hydration enthalpy of halide ion, higher will be oxidising of halogen and vice versa.
HF is a weaker acid than HI. This is due to following reasons.
1. Presence of strong hydrogen bonding in HF which decreases its tendency to donate proton, while in HI, there is no hydrogen bonding. .
2 Bond enthalpy of H-F is considerably higher due to the highest electronegativity of F atom, while bond enthalpy of HI is lower due to lower electronegativity of I atom.
3. Conjugate base of HF is highly unstable and readily combine with proton. On the other hand, conjugate base of HI is highly stable.
Question 23. On what factors does the oxidizing power of halogens depend?
Answer.
Oxidizing power of halogens depends upon following factors.
1. Bond enthalpy of X-X
2. Electron affinity of X
3. Hydration energy of halide ion (X-).
1. Bond enthalpy of X-X
Greater the X-X bond enthalpy, lower will be oxidizing power and vice versa.
2. Electron Affinity
Greater the electron affinity of halogen, higher will be oxidizing power and vice versa.
3. Hydration enthalpy
Greater the hydration enthalpy of halide ion, higher will be oxidising of halogen and vice versa.
Very simple notes,
ReplyDeleteExcellent work
A nice work
ReplyDelete