Types of bonding and Forces in solid Questions in English

(D) $AlCl_3$ is a covalent compound that exists as a dimer $(Al_2Cl_6)$ in the solid state.
Due to weak van der Waals forces between the molecules,it sublimes readily upon heating at $180 \ ^{\circ}C$ under atmospheric pressure.

(B) $NaCl$ is an ionic compound because it has a high electrostatic force of attraction between $Na^+$ and $Cl^-$ ions,resulting in a high melting point and solid state at room temperature.
In contrast,$HCl$ is a covalent molecule with weak intermolecular forces (van der Waals forces),which makes it a gas at room temperature.

(C) . Quartz $(SiO_2)$ is a covalent network solid (also known as a network solid). In this structure,the constituent particles are silicon and oxygen atoms,which are held together by strong covalent bonds throughout the crystal lattice.

(A) The constituent particles of covalent compounds are atoms,which are held together by covalent bonds. Therefore,the structural units in crystals of covalent compounds are atoms.

(C) The radius ratio is calculated as follows:
$\frac{r_{c^+}}{r_{a^-}} = \frac{1.46 \ \mathring{A}}{2.16 \ \mathring{A}} = 0.676$
Since the radius ratio $0.676$ lies in the range $0.414 - 0.732$,it corresponds to an octahedral geometry.
Therefore,the compound exhibits a coordination number of $6$ and a structure of the $NaCl$ type.

(A) $HCl$ is a polar molecule with a permanent dipole moment $(\mu \neq 0)$.
$He$ is a non-polar atom with no permanent dipole moment $(\mu = 0)$.
When a polar molecule approaches a non-polar atom,it induces a dipole in the non-polar atom.
This interaction between a permanent dipole and an induced dipole is known as a dipole-induced dipole interaction.
Therefore,the pair $HCl$ and $He$ atoms exhibits this interaction.

(B) To determine which substances exist as cation-anion pairs in the solid state:
$1$. $I_{2(s)}$: It is a molecular solid held by van der Waals forces. It does not contain ions.
$2$. Dry ice $(CO_{2(s)})$: It is a molecular solid held by van der Waals forces. It does not contain ions.
$3$. $PCl_{5(s)}$: In the solid state,it exists as $[PCl_4]^+ [PCl_6]^-$,which is an ionic lattice.
$4$. $N_2O_{5(s)}$: In the solid state,it exists as $[NO_2]^+ [NO_3]^-$,which is an ionic lattice.
$5$. $PBr_{5(g)}$: In the gaseous state,it exists as discrete $PBr_5$ molecules. It does not contain ions.
Thus,the molecules that have cation-anion pairs are $PCl_{5(s)}$ and $N_2O_{5(s)}$.
The total number is $2$.

(D) In the solid state,ionic substances exist as a lattice of ions.
$PBr_5$ exists as an ionic solid composed of $[PBr_4]^+$ and $[Br]^-$ ions.
$N_2O_5$ exists as an ionic solid composed of $[NO_2]^+$ and $[NO_3]^-$ ions.
$Na_2SO_4$ is a typical ionic salt consisting of $[Na]^+$ and $[SO_4]^{2-}$ ions.
$H_2O$ in the solid state (ice) exists as a molecular solid held together by hydrogen bonding,not as an ionic substance.
Therefore,the correct answer is $H_2O$.

(D) The interaction between an ion $(X^{-})$ and a non-polar molecule $(X_2)$ leads to the formation of an induced dipole in the non-polar molecule.
This interaction is known as an ion-induced dipole attraction.
In the case of $Br^{-} + Br_2 \rightarrow Br_{3}^{-}$,the bromide ion induces a dipole in the bromine molecule,which facilitates the formation of the polyhalide ion $Br_{3}^{-}$.

(D) Concept: For a salt to be soluble in water,its hydration energy must be greater than its lattice energy.
In the case of $Na_2SO_4$,the hydration energy is greater than the lattice energy,making it soluble.
In the case of $BaSO_4$,the lattice energy is significantly higher than the hydration energy,which makes it insoluble.
Therefore,both statements $(A)$ and $(C)$ are correct.

(D) Potash alum is a double salt with the formula $K_2SO_4 \cdot Al_2(SO_4)_3 \cdot 24H_2O$.
Double salts are compounds that dissociate completely into their constituent ions when dissolved in water.
Therefore,in an aqueous solution,potash alum dissociates to give $K^+$,$Al^{3+}$,and $SO_4^{2-}$ ions.
Since all these ions are present in the solution,it exhibits the characteristics of all of them.

(A) The bond length is inversely proportional to the bond order.
In $Diamond$, the $C-C$ bond is a pure single bond (bond order = $1$, bond length = $154 \text{ pm}$).
In $Graphite$, the $C-C$ bond has partial double bond character (bond order = $1.33$, bond length = $141.5 \text{ pm}$).
In $Benzene$, the $C-C$ bond has partial double bond character (bond order = $1.5$, bond length = $139 \text{ pm}$).
In $Ethene$ $(CH_2=CH_2)$, the $C-C$ bond is a double bond (bond order = $2$, bond length = $134 \text{ pm}$).
Therefore, the bond length is maximum in $Diamond$.

(C) dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a non-polar molecule by disturbing the arrangement of electrons in the non-polar species.

(A) In the solid state,$PCl_5$ exists as an ionic compound composed of $[PCl_4]^+$ and $[PCl_6]^-$ ions.
This occurs because the ionic lattice is more stable than the molecular form in the solid state.
The $[PCl_4]^+$ cation has a tetrahedral geometry,while the $[PCl_6]^-$ anion has an octahedral geometry.

(A) network solid (or covalent network solid) consists of a continuous 3D network of covalent bonds.
$SiO_2$ (silica) forms a giant covalent structure where each silicon atom is bonded to four oxygen atoms in a tetrahedral arrangement,creating a rigid network.
In contrast,$NO_2$,$SO_2$,and $CO_2$ are molecular solids held together by weak van der Waals forces.

(B) The potential energy of interaction for dispersion forces (London forces) is proportional to $1/r^6$,where $r$ is the distance between particles.
Because of this high power dependence on distance,dispersion forces are the most sensitive to changes in distance compared to other intermolecular forces.

(B) Water in the solid phase (ice) is best described as a molecular solid.
In ice,$H_2O$ molecules are held together by hydrogen bonding,which is a type of intermolecular force characteristic of molecular solids.

(A) In the solid state,$BeCl_2$ exists as a polymeric chain structure where each $Be$ atom is surrounded by four $Cl$ atoms.
In the vapour phase,at high temperatures (approx. $1200 \ K$),$BeCl_2$ exists as a dimeric structure $(Be_2Cl_4)$,which dissociates into a linear monomeric structure at even higher temperatures (approx. $1800 \ K$).

(D) The correct order of molecular force of attraction is $H_2O > OBr_2 > O(CH_3)_2$.
$H_2O$ exhibits strong hydrogen bonding.
For $OBr_2$ and $O(CH_3)_2$,the forces are dominated by dipole-dipole interactions,which are generally proportional to molecular mass.
Since the molar mass of $OBr_2$ $(183.8 \ g/mol)$ is significantly higher than that of $O(CH_3)_2$ $(60.1 \ g/mol)$,the force of attraction in $OBr_2$ is greater than in $O(CH_3)_2$.
Therefore,the order given in option $D$ is incorrect.

(C) London dispersion forces are directly proportional to the molecular mass and the size of the electron cloud of a molecule.
As we move down Group $14$ in the periodic table,the atomic size and molecular mass of the hydrides increase in the order: $CH_4 < SiH_4 < GeH_4 < SnH_4$.
Therefore,$SnH_4$ has the largest molecular mass and the most polarizable electron cloud,resulting in the strongest London dispersion forces.

(D) The melting point depends on the strength of intermolecular forces and the type of lattice structure:
$1$. $C_2H_6$ (Ethane) is a non-polar molecule with weak van der Waals forces (molecular lattice),resulting in the lowest melting point.
$2$. $CH_3OH$ (Methanol) is a polar molecule with hydrogen bonding (molecular lattice),which is stronger than the forces in $C_2H_6$.
$3$. $KCl$ (Potassium chloride) is an ionic compound with strong electrostatic forces (ionic lattice),requiring more energy to break.
$4$. $Si$ (Silicon) forms a giant covalent network (covalent lattice),which has the highest melting point among these substances.
Therefore,the correct order of increasing melting point is: $C_2H_6 < CH_3OH < KCl < Si$.

(C) $CO$,$CO_2$,and $P_2O_5$ are covalent compounds that exist as molecular lattices with weak intermolecular forces.
$SiO_2$ is a covalent compound that possesses a $3$-dimensional network structure (giant covalent lattice).
Due to the strong covalent bonds extending throughout the entire structure,$SiO_2$ has a significantly higher melting point compared to the others.

(C) The lattice energy $(LE)$ is directly proportional to the product of the charges of the ions and inversely proportional to the sum of their ionic radii: $LE \propto \frac{q^+ \cdot q^-}{r^+ + r^-}$.
$A) AlF_3 > MgF_2$: $Al^{3+}$ has a higher charge than $Mg^{2+}$,so $AlF_3$ has higher lattice energy.
$B) Li_3N > Li_2O$: $N^{3-}$ has a higher charge than $O^{2-}$,so $Li_3N$ has higher lattice energy.
$C) NaCl > LiF$: The ionic radii of $Li^+$ and $F^-$ are smaller than those of $Na^+$ and $Cl^-$. Therefore,$LiF$ should have higher lattice energy than $NaCl$. The given order $NaCl > LiF$ is incorrect.
$D) TiC > ScN$: $Ti^{4+}C^{4-}$ has higher charges than $Sc^{3+}N^{3-}$,resulting in higher lattice energy for $TiC$.
Thus,the incorrect order is $NaCl > LiF$.

(A) $I_2$ is a non-polar molecule.
In the solid state,$I_2$ molecules are held together by weak van der Waals forces,specifically London dispersion forces.
Therefore,the correct option is $(A)$.

(D) $CO_2$ exists as discrete molecules held together by weak van der Waals' forces,making it a gas at room temperature.
In contrast,$Si$ has a larger atomic size and cannot form stable $p\pi-p\pi$ multiple bonds with $O$ atoms.
Therefore,$Si$ forms four single covalent bonds with $O$ atoms,resulting in a giant $3D$ covalent network structure,which makes $SiO_2$ a solid.

(D) network solid,also known as a covalent network solid,is a material in which atoms are linked by a continuous network of covalent bonds extending throughout the entire crystal structure.
In such solids,there are no discrete individual molecules; instead,the entire crystal acts as a single macromolecule.
Diamond consists of a three-dimensional network of carbon atoms,and $SiO_2$ (quartz) consists of a three-dimensional network of silicon and oxygen atoms.
Conversely,$P_4O_{10}$,$P_4O_6$,and $SO_3$ are molecular solids consisting of discrete molecules held together by weak van der Waals forces.
Therefore,the pair consisting of only network solids is Diamond and $SiO_2$.
Hence,option $D$ is correct.

(D) Noble gases are monoatomic and non-polar in nature.
Since they lack permanent dipoles,the only intermolecular forces acting between their atoms are weak,temporary induced dipole-induced dipole attractions,which are known as $London$ dispersion forces or van der Waals forces.
These forces increase with the size and polarizability of the atoms,which explains the trend in their boiling points.

(B) The stability of an ionic crystal is primarily determined by its lattice energy.
Lattice energy is the energy released when gaseous ions combine to form one mole of an ionic solid.
$A$ higher magnitude of lattice energy indicates a more stable crystal structure,as it represents the net attractive interactions within the crystal lattice.

(B) $AlF_3$ is an ionic compound with a high lattice energy due to its three-dimensional giant ionic structure,resulting in a high melting point $(1040 \, ^{\circ}C)$.
$SiF_4$ is a covalent molecule with weak van der Waals forces of attraction between the molecules,leading to a very low melting point ($-77 \, ^{\circ}C$ at which it sublimes).

(D) Lattice energy is directly proportional to the product of the charges of the ions and inversely proportional to the sum of the ionic radii $(U \propto \frac{q_1 q_2}{r_+ + r_-})$.
$(i)$ Between $KCl$ $(K^+, Cl^-)$ and $MgO$ $(Mg^{2+}, O^{2-})$,$MgO$ has higher charges $(+2, -2)$ and smaller ionic radii,so $MgO$ has higher lattice energy.
$(ii)$ Between $LiF$ $(Li^+, F^-)$ and $LiBr$ $(Li^+, Br^-)$,the charges are the same. Since $F^-$ is smaller than $Br^-$,$LiF$ has higher lattice energy.
$(iii)$ Between $Mg_3N_2$ $(Mg^{2+}, N^{3-})$ and $NaCl$ $(Na^+, Cl^-)$,$Mg_3N_2$ has much higher ionic charges,resulting in significantly higher lattice energy.
Therefore,the set with higher lattice energy is $MgO, \, LiF, \, Mg_3N_2$.

(B) The solubility of an ionic compound in water depends on the balance between its lattice energy and hydration energy.
For $Al_2O_3$,the lattice energy is extremely high due to the high charge density of $Al^{3+}$ and $O^{2-}$ ions.
When $Al_2O_3$ is placed in water,the energy released during the hydration of these ions (hydration energy) is not sufficient to overcome the very high lattice energy required to break the crystal structure.
Therefore,$Al_2O_3$ remains insoluble in water.

(C) Graphite is used as a lubricant due to its slippery nature.
Graphite consists of layers of carbon atoms.
These layers are held together by weak van der Waals forces,which allow the layers to slide over one another easily,making it an effective lubricant.

(D) $Si(IV)$ oxide is a macromolecular solid with a three-dimensional non-layer structure,similar to the structure of diamond,where each silicon atom is tetrahedrally bonded to four oxygen atoms.

(C) The boiling point of hydrogen halides depends on the magnitude of van der Waals' forces of attraction.
As the size of the halogen atom increases from $Cl$ to $I$,the molecular mass and the surface area of the molecule increase.
This leads to an increase in the magnitude of van der Waals' forces of attraction.
Therefore,the boiling point increases in the order $HCl < HBr < HI$.

(A) $S_8$ is a molecular solid,where $S_8$ molecules are held together by weak van der Waals forces,not a covalent lattice.
$P_4$ has a tetrahedral structure.
$S_4^{2-}$ (polysulfide ion) exhibits a zig-zag chain structure.
$SiO_2$ is a network solid with a covalent lattice structure where each $Si$ atom is bonded to four $O$ atoms and each $O$ atom is bonded to two $Si$ atoms.
Therefore,the incorrectly matched characteristic is $S_8$ : Covalent lattice.

(D) $SiH_4$ has the lowest boiling point of $161 \ K$.
$HCl$ has a boiling point of $188 \ K$.
$H_2S$ has a boiling point of $216 \ K$.
$PH_3$ has a boiling point of $185 \ K$.
Therefore,$SiH_4$ has the lowest boiling point among the given substances.

(C) Noble gases are monoatomic and non-polar. The only attractive forces present between noble gas atoms are weak $Van \ der \ Waals$ forces,specifically known as $London \ dispersion \ forces$. Since $London \ dispersion \ forces$ are a type of $Van \ der \ Waals$ force,and the question asks for the effective force,$London \ dispersion \ forces$ is the most specific and accurate description.

(D) In the solid state,phosphorus pentachloride $(PCl_5)$ exists as an ionic solid.
It undergoes auto-ionization to form a mixture of tetrahedral $PCl_4^+$ and octahedral $PCl_6^-$ ions.
Therefore,the correct representation is $[PCl_4]^+ [PCl_6]^-$.
Hence,the correct option is $D$.

(D) The lattice energy $(U)$ of an ionic crystal is directly proportional to the product of the charges $(q_1, q_2)$ and inversely proportional to the sum of the ionic radii $(r_+ + r_-)$.
$U \propto \frac{|q_1 \times q_2|}{r_+ + r_-}$
For $JX$: $q_1=1, q_2=1, r_+ + r_- = 0.14 + 0.14 = 0.28 \ nm$. So,$U \propto \frac{1}{0.28} \approx 3.57$.
For $LY$: $q_1=1, q_2=1, r_+ + r_- = 0.18 + 0.18 = 0.36 \ nm$. So,$U \propto \frac{1}{0.36} \approx 2.78$.
For $MZ$: $q_1=2, q_2=2, r_+ + r_- = 0.15 + 0.15 = 0.30 \ nm$. So,$U \propto \frac{4}{0.30} \approx 13.33$.
Comparing the values: $13.33 > 3.57 > 2.78$,which corresponds to $MZ > JX > LY$.

(C) Electrical conductivity in ionic compounds requires the presence of free mobile ions.
In molten $NaOH$,molten $KOH$,and aqueous $NaCl$,the ions are free to move,allowing them to conduct electricity.
In solid $NaCl$,the ions are held in a rigid crystal lattice structure by strong electrostatic forces and are not free to move.
Therefore,solid $NaCl$ does not conduct electricity.

(D) Covalent compounds consist of discrete molecules.
These molecules are held together by weak intermolecular forces known as van der Waals forces.
Because these forces are weak,they require very little energy to overcome,resulting in low melting and boiling points for covalent compounds.

(A) $KCl$ is an ionic compound.
In water,it dissociates into $K^{+}$ and $Cl^{-}$ ions.
The interaction between $KCl$ and $H_2O$ is an $\text{ion-dipole}$ interaction,not a $\text{dipole-dipole}$ interaction.
Therefore,option $A$ is incorrectly matched.

(C) Instantaneous dipole-induced dipole attractions,also known as London dispersion forces,occur between non-polar molecules or atoms.
$CO_2$ is a non-polar molecule due to its linear geometry where the dipole moments of the two $C=O$ bonds cancel each other out.
Therefore,the attraction between two $CO_2$ molecules is an example of instantaneous dipole-induced dipole interaction.

(B) The formation of an ionic lattice from gaseous ions,$Na_{(g)}^{+} + Cl_{(g)}^{-} \to NaCl_{(s)}$,releases a large amount of energy known as lattice energy,which is an exothermic process.
In contrast,adding an electron to an anion (like $S_{(g)}^{-}$) or removing an electron from a cation (like $Al_{(g)}^{+2}$) generally requires energy (endothermic) due to electrostatic repulsion or high ionization energy.

(A) $SiO_2$ (silica) is a network covalent solid with a giant $3-D$ tetrahedral structure.
Due to the presence of strong covalent bonds throughout the entire crystal lattice,it requires a very high amount of energy to break these bonds,resulting in a very high melting point (approximately $1973 \ K$).
In contrast,$NaCl$,$Na_2CO_3$,and $NaF$ are ionic compounds which,while having high melting points,are generally lower than that of the network covalent solid $SiO_2$.

(B) covalent solid (or network solid) consists of atoms held together by a continuous network of covalent bonds.
$Diamond$ $(C)$,$Graphite$ $(C)$,and $Quartz$ $(SiO_2)$ are classic examples of covalent network solids.
$Phosphorus$ $(P_4)$ is a molecular solid,where individual $P_4$ molecules are held together by weak van der Waals forces.
Therefore,$Phosphorus$ is not a covalent solid.

(D) Metallic luster is primarily due to the presence of $free \ electrons$ in the metallic lattice.
These $free \ electrons$ absorb incident light and re-emit it,which gives the metal its characteristic shiny appearance.

(A) $CO_2$ exists as discrete,linear molecules held together by weak van der Waals forces,making it a gas at room temperature.
In contrast,$SiO_2$ (silica) forms a giant three-dimensional covalent network structure where each silicon atom is bonded to four oxygen atoms in a tetrahedral geometry.
This continuous network requires a large amount of energy to break,resulting in a high melting point and a solid state at room temperature.

(D) Graphite has a layered structure where each carbon atom is $sp^2$ hybridized and covalently bonded to three other carbon atoms in the same layer,forming hexagonal rings.
These layers are held together by weak van der Waals forces,which allows the layers to slide over each other,making graphite a soft solid lubricant.
The strong covalent bonds within the layers require a very high temperature to break,making it extremely difficult to melt.