In the first-year chemistry course, Chapter 4, titled “Liquids and Solids,” delves into the fascinating world of condensed states of matter. This chapter explores the unique properties and behaviors of liquids and solids, which are essential components of our everyday lives. It begins by discussing the various intermolecular forces that govern the behavior of molecules in the liquid and solid phases, such as van der Waals forces, dipole-dipole interactions, and hydrogen bonding. Students learn how these forces influence properties like boiling points, melting points, and vapor pressures.
The chapter also covers phase diagrams, which help us understand how substances transition between different states under varying temperature and pressure conditions. Moreover, it highlights the concept of crystalline and amorphous solids, discussing their structures and properties. Understanding the principles outlined in Chapter 4 is crucial not only for comprehending the behavior of matter but also for applications in various fields, including materials science and chemical engineering.
Short Questions Chemistry 1st Year Chapter 4 Liquids And Solids
What is the reason for the existence of matter in the form of gases, liquids, and solids in our surroundings?
The existence of matter in our surroundings in the form of gases, liquids, and solids is due to the difference in interacting forces among the constituent particles.
What are the weak forces of attraction that exist between molecules when they are sufficiently close to each other?
The weak forces of attraction that exist between molecules when they are sufficiently close to each other are called van der Waals forces.
- 1st Year Biology Unit No.1 Introduction Notes
- 1st Year Biology Unit No. 2 Biological Molecules Notes
- 1st Year Biology Unit No. 3 Enzymes Notes
- 1st Year Biology Unit No. 4 The Cell Notes
- 1st Year Biology Unit No. 5 Variety of Life Notes
What are the four types of intermolecular forces?
The four types of intermolecular forces mentioned are:
- Dipole-dipole forces
- Ion-dipole forces
- Dipole-induced dipole forces
- Instantaneous dipole-induced dipole forces or London dispersion forces
What are dipole-dipole forces, and how do they arise in polar molecules?
Dipole-dipole forces arise in polar molecules where there is an electronegativity difference between the bonded atoms. The more electronegative atom develops a partial negative charge, while the less electronegative atom develops a partial positive charge. These opposite charges cause molecules to align, leading to electrostatic forces of attraction between them.
How effective are dipole-dipole forces compared to covalent bonds?
Dipole-dipole forces are approximately one percent as effective as a covalent bond.
What factors influence the strength of dipole-dipole forces?
The strength of dipole-dipole forces depends on the electronegativity difference between the bonded atoms and the distance between the molecules.
In which phase (gaseous or liquid) are dipole-dipole forces stronger, and why?
Dipole-dipole forces are stronger in the liquid phase because the distances between molecules in liquids are smaller compared to the gaseous phase, allowing for stronger interactions.
What are dipole-induced dipole forces also known as?
Dipole-induced dipole forces are also known as Debye forces.
How do dipole-induced dipole forces work between polar and non-polar molecules?
The positive end of a polar molecule attracts the mobile electrons of nearby non-polar molecules, inducing polarity in the non-polar molecule and creating dipoles in both molecules.
Who proposed a simple explanation for the weak attractive forces between non-polar molecules?
Fritz London, a German physicist, proposed a simple explanation for these weak attractive forces in non-polar molecules.
What is an “instantaneous dipole” in the context of London dispersion forces?
An instantaneous dipole occurs when the electron density of an atom becomes temporarily asymmetrical due to the repulsion between electrons, leading to a momentary dipole in the atom.
What is the attraction between an instantaneous dipole and an induced dipole called?
The attraction between an instantaneous dipole and an induced dipole is called “instantaneous dipole-induced dipole interaction” or “London force.”
Are London forces present in all types of molecules?
Yes, London forces are present in all types of molecules, whether they are polar or non-polar.
What factors affect the strength of London dispersion forces?
The strength of London forces depends on the size of the electronic cloud, polarizability, and the number of atoms in a non-polar molecule.
What is hydrogen bonding, and why does it occur in water molecules?
Hydrogen bonding is an electrostatic force of attraction between a highly electronegative atom and a partially positively charged hydrogen atom. In water molecules, it occurs because oxygen is more electronegative than hydrogen, creating a polar molecule with dipole-dipole interactions.
How does hydrogen bonding differ from simple dipole-dipole interactions?
Hydrogen bonding is stronger than simple dipole-dipole interactions due to the presence of two lone pairs on the oxygen atom and the sufficient partial positive charge on the hydrogen atom.
Which atoms are responsible for creating hydrogen bonds in molecules?
Electronegative atoms like fluorine, oxygen, nitrogen, and rarely chlorine are responsible for creating hydrogen bonds.
How does hydrogen bonding affect the strength of a covalent bond?
The strength of a hydrogen bond is generally twenty times weaker than that of a covalent bond.
Can hydrogen bonding occur with atoms other than fluorine, oxygen, nitrogen, and chlorine?
Yes, hydrogen bonding can occur with other atoms. For example, in chloroform, the chlorine atoms are responsible for hydrogen bonding with other molecules, like acetone.
Why is hydrogen bonding responsible for the low acidic strength of HF compared to HCl, HBr, and HI?
The strong hydrogen bonding in HF molecules traps the partially positively charged hydrogen atom between two highly electronegative atoms, reducing its acidity compared to the other halogen acids.
What is the influence of hydrogen bonding on the physical properties of covalent hydrides?
Hydrogen bonding affects physical properties like melting and boiling points by creating attractive forces between partial positively charged hydrogen and highly electronegative atoms with partial negative charges.
Why do hydrides of group IV-A have lower boiling points compared to those of group V-A, VI-A, VII-A?
Hydrides of group IV-A have lower boiling points because the elements in this group are less electronegative, and CH4 has the lowest boiling point due to its small size and low polarizability.
What causes NH3, H2O, and HF to have higher boiling points compared to other hydrides?
NH3, H2O, and HF have higher boiling points because of their enhanced electronegativity (N, O, and F), allowing for stronger hydrogen bonding. This results in water being liquid at room temperature while H2S and H2Se are gases.
Why does the boiling point of HF, despite fluorine being more electronegative than oxygen, lower than that of H2O?
The boiling point of HF is lower than that of H2O because fluorine can form only one hydrogen bond with electropositive hydrogen, while water can form two hydrogen bonds per molecule due to its two hydrogen atoms and two lone pairs on the oxygen atom.
Why does HBris have a slightly higher boiling point than HCl?
HBris has a slightly higher boiling point than HCl because chlorine is electronegative enough to form a hydrogen bond, even though HCl is a borderline case of dipole-dipole interaction.
Why do fourth-period hydrides (GeH4, AsH3, H2Se, HBr) have greater boiling points than third-period hydrides?
Fourth-period hydrides have greater boiling points than third-period hydrides due to their larger size and enhanced polarizabilities.
Why can water dissolve ethyl alcohol but not hydrocarbons?
Water can dissolve ethyl alcohol because both can form hydrogen bonds with each other. Hydrocarbons are not soluble in water because they are non-polar compounds and cannot form hydrogen bonds with water molecules.
What causes ice to have a lower density than liquid water, and why does it float on water?
Ice has a lower density than liquid water because its molecules form a regular tetrahedral structure with empty spaces, causing it to occupy more space and have lower density. Ice floats on water because of this lower density, and it insulates the water beneath, allowing aquatic life to survive.
How does the structure of ice compare to that of diamond?
The structure of ice is similar to that of diamond, as both have atoms at the centers of tetrahedrons. In ice, oxygen atoms occupy the center of the tetrahedron, while in diamond, carbon atoms do.
What role does hydrogen bonding in water play in the patterns of life for plants and animals?
Hydrogen bonding in water is essential for life as it affects the properties of water, such as its density and heat insulation abilities. Without hydrogen bonding in water, the patterns of life for plants and animals would have been significantly different.
Why do soaps and detergents perform cleansing action?
Soaps and detergents perform cleansing action because the polar part of their molecules is water-soluble due to hydrogen bonding, while the non-polar parts (alkyl or benzyl portions) remain insoluble in water.
What is the structure of proteins like silk and muscles?
Proteins found in silk and muscles consist of long chains of amino acids that are coiled about one another into a spiral, known as a helix. This helix can be either right-handed or left-handed and is linked together by hydrogen bonds.
How is DNA structured?
DNA has two spiral chains coiled about each other on a common axis, forming a double helix. These chains are linked together by hydrogen bonding between their subunits.
What role does hydrogen bonding play in food materials like carbohydrates?
Hydrogen bonding is responsible for the properties of food materials like carbohydrates, including glucose, fructose, and sucrose, because they contain -OH groups that can form hydrogen bonds.
Why do paints and dyes exhibit adhesive properties?
Paints and dyes exhibit adhesive properties due to hydrogen bonding, which is one of their most important properties. This hydrogen bonding allows them to adhere effectively.
How does hydrogen bonding contribute to the properties of clothing materials like cotton, silk, and synthetic fibers?
Hydrogen bonding in cotton, silk, and synthetic fibers is crucial for their rigidity and tensile strength, making them suitable for use in clothing materials.
What is evaporation, and what causes it?
Answer: Evaporation is the spontaneous change of a liquid into its vapors. It occurs when high-energy molecules at the liquid’s surface escape the attraction of neighboring molecules and leave the bulk of the liquid. This process continues at all temperatures.
How does evaporation cause cooling?
Evaporation causes cooling because when high-energy molecules leave the liquid, the remaining low-energy molecules experience a decrease in temperature. Heat is transferred from the surroundings to the liquid, resulting in a temperature drop in both the liquid and its surroundings.
What factors control the rate of evaporation?
The rate of evaporation is controlled by several factors:
a. Surface area: Increasing the surface area allows more molecules to escape, leading to quicker evaporation.
b. Temperature: Higher temperatures increase the kinetic energy of molecules, accelerating the rate of evaporation.
c. Strength of intermolecular forces: Weaker intermolecular forces lead to faster evaporation. For example, gasoline, with weaker London forces of attraction, evaporates faster than water.
What is vapor pressure, and how is it related to evaporation?
Vapor pressure is the pressure exerted by the vapors of a liquid in equilibrium with the liquid at a given temperature. It represents the point at which the rate of evaporation equals the rate of condensation. Vapor pressure is not influenced by the amount of liquid, volume of the container, or the liquid’s surface area.
What happens when a liquid reaches a state of dynamic equilibrium?
In a state of dynamic equilibrium, the rate of evaporation is equal to the rate of condensation. This means that the number of molecules leaving the liquid’s surface is balanced by the number of molecules returning to it. It occurs at a constant temperature and represents a stable state for evaporation and condensation processes.
What is the most important parameter that controls the vapor pressure of a liquid?
The most important parameter that controls the vapor pressure of a liquid is its temperature.
How does an increase in temperature affect the vapor pressure of a liquid?
An increase in temperature increases the kinetic energy of molecules, leading to an increase in vapor pressure.
Provide an example of how vapor pressure changes with temperature for water.
Vapor pressure of water increases from 4.579 torr to 9.209 torr when the temperature changes from 0°C to 10°C. It increases from 527.8 torr to 760 torr when the temperature changes from 90°C to 100°C.
How is the strength of intermolecular forces related to the vapor pressure of a liquid at a particular temperature?
The stronger the intermolecular forces, the lower the vapor pressure of the liquid at a particular temperature.
What method is described for measuring the vapor pressure of a liquid accurately?
The manometric method is described as a comparatively accurate method for measuring the vapor pressure of a liquid.
How is the vapor pressure of a liquid determined using the manometric method?
The difference in the heights of mercury columns in a manometer is used to determine the vapor pressure of the liquid. The equation P = Pa + Δh is used, where P is the vapor pressure, Pa is the atmospheric pressure, and Δh is the difference in mercury column heights.
What is the boiling point of a liquid?
The boiling point of a liquid is the temperature at which its vapor pressure becomes equal to the external atmospheric pressure.
What happens at the boiling point of a liquid?
At the boiling point, bubbles of vapor are formed inside the liquid, and they have greater internal pressure than the atmospheric pressure on the liquid’s surface, causing the bubbles to rise and burst.
What is the molar heat of vaporization?
The molar heat of vaporization is the amount of heat required to vaporize one mole of a liquid at its boiling point.
How does the vapor pressure of a liquid change as it approaches its boiling point?
The vapor pressure of a liquid increases very rapidly as it approaches its boiling point.
What causes a liquid to boil?
When the vapor pressure of a liquid equals the external pressure.
How does changing external pressure affect the boiling point of a liquid?
Higher external pressure raises the boiling point, while lower external pressure lowers it.
Why does water boil at a lower temperature in high-altitude locations like Mount Everest?
Answer: Due to the lower external pressure at high altitudes.
How does a pressure cooker increase the boiling temperature of water?
It traps vapor, increasing pressure inside the cooker, which raises the boiling temperature.
What is vacuum distillation, and why is it advantageous?
Vacuum distillation is a process where liquids boil at lower temperatures under reduced pressure, saving time and fuel while preventing decomposition.
What is ΔHf, and what does it represent?
ΔHf is the molar heat of fusion, representing the heat absorbed when one mole of a solid melts into a liquid at constant pressure.
What is ΔHv, and when does it occur?
ΔHv is the molar heat of vaporization, occurring when one mole of a liquid turns into vapor at its boiling point under one atmosphere of pressure.
What is ΔHs, and when does it occur?
ΔHs is the molar heat of sublimation, happening when one mole of a solid sublimes into vapor at a specific temperature and one atmospheric pressure.
Why are ΔHv values typically larger than ΔHf values?
ΔHv represents the energy needed to separate molecules in a liquid, which involves larger changes in intermolecular distances than melting solids.
What does dynamic equilibrium refer to in the context of a change of state?
Dynamic equilibrium is a state where two opposing changes occur at equal rates, such as the equilibrium between solid ice and liquid water at 0°C.
What did Frederick Reinitzer discover in 1888 while studying cholesteryl benzoate?
Frederick Reinitzer discovered that cholesteryl benzoate exhibited a turbid liquid phase at a certain temperature range, which he called a “liquid crystal.”
What distinguishes liquid crystals from regular liquids and crystalline solids?
Liquid crystals have properties intermediate between those of crystalline solids and isotropic liquids. They exhibit the fluidity of liquids and the optical properties of crystals.
How are liquid crystals categorized based on the nature of molecular ordering?
Liquid crystals can be categorized into nematic, smectic, and cholesteric phases depending on the nature of molecular ordering within them.
What are some practical applications of liquid crystals?
Some practical applications of liquid crystals include temperature sensors, detecting potential failures in electrical circuits, locating veins, arteries, infections, and tumors, displaying information in devices like digital watches and laptops, serving as solvents in chromatographic separations, and being used in oscillographic and TV displays.
How are liquid crystals utilized in temperature sensors?
Liquid crystals can diffract light and change color as the temperature changes due to alterations in the distances between layers of liquid crystal molecules. This property allows them to be used as temperature sensors.
What are solids, and what distinguishes them from other substances?
Solids are substances that are rigid, hard, have a definite shape, and a definite volume. They have closely packed atoms, ions, or molecules held together by strong cohesive forces, and they exhibit a well-ordered arrangement.
How can solids be classified based on the arrangement of their constituent particles?
Solids can be classified into two types based on the arrangement of their constituent atoms, ions, or molecules: crystalline solids and amorphous solids.
What are crystalline solids, and what characterizes their structure?
Crystalline solids are solids in which atoms, ions, or molecules are arranged in a definite three-dimensional pattern, exhibiting recurring regular geometrical patterns that extend in three dimensions.
Provide examples of amorphous solids and explain their characteristics.
Amorphous solids, such as glass, plastics, rubber, and glue, do not have a regular, orderly arrangement of their constituent atoms, ions, or molecules. They lack a well-defined crystalline structure.
How can crystalline solids be converted into amorphous solids?
Crystalline solids can be transformed into amorphous solids by melting them and then rapidly cooling the molten mass. This prevents the constituent particles from having time to arrange themselves in an orderly manner.
What are some properties of crystalline solids?
Crystalline solids have distinct properties, including:
a. Geometrical shape
b. Sharp melting points
c. Cleavage planes
d. Anisotropic properties
e. Symmetry
f. Habit of a crystal
g. Isomorphism
What is meant by “cleavage planes” in crystalline solids?
Cleavage planes are specific planes along which crystalline solids tend to break when fractured. These planes have characteristic angles for each crystalline solid.
What is “anisotropy,” and how does it relate to crystalline solids?
Anisotropy refers to the variation in physical properties of crystalline solids depending on the direction. Some crystals exhibit anisotropic properties because of the different arrangements of particles in different directions within the crystal.
Explain the concept of “isomorphism” in crystalline solids.
Isomorphism is the phenomenon in which two different substances exist in the same crystalline form, with identical or similar arrangements of atoms, ions, or molecules. Isomorphic substances can crystallize together in various proportions in homogeneous mixtures.
What is polymorphism?
Polymorphism is a phenomenon in which a compound exists in more than one crystalline form.
What are the different forms of a compound in polymorphism called?
The different forms of a compound in polymorphism are called polymorphs.
Do polymorphs have the same chemical properties?
Yes, polymorphs have the same chemical properties.
What causes the differences in physical properties among polymorphs?
The differences in physical properties among polymorphs are due to different structural arrangements of their particles.
Give examples of compounds that exhibit polymorphism.
Examples of compounds that exhibit polymorphism include AgNO3, CaCO3, and others.
What is allotropy?
Allotropy is the existence of an element in more than one crystalline form.
What are the different forms of an element in allotropy called?
The different forms of an element in allotropy are called allotropes or allotropic forms.
Provide examples of elements that exhibit allotropy.
Examples of elements that exhibit allotropy include sulfur (S), carbon (C), and tin (Sn).
What is a transition temperature?
A transition temperature is the temperature at which two crystalline forms of the same substance can coexist in equilibrium with each other.
How do the transition temperatures of allotropic forms compare to their melting points?
The transition temperature of allotropic forms is always less than their melting points.
What are lattice points in a crystal lattice?
Lattice points in a crystal lattice are the positions where atoms, ions, or molecules are located in a crystalline solid.
What is a unit cell in a crystal lattice?
A unit cell is the smallest part of a crystal lattice that has all the characteristic features of the entire crystal and can be used to build up the crystal in three dimensions.
How is the crystal lattice’s structure determined from the unit cell?
The complete information about the crystalline structure is present within a unit cell, so knowing the arrangement of atoms in a unit cell provides information about their arrangement in the whole crystal.
What are the six parameters used to describe a unit cell’s size and shape?
The six parameters are unit cell lengths a, b, c, and unit cell angles α, β, γ.
How many crystal systems are there, and can you name them?
There are seven crystal systems: cubic, tetragonal, orthorhombic, monoclinic, hexagonal, rhombohedral (trigonal), and triclinic.
Provide an example of a substance for each crystal system.
Examples include cubic (Fe), tetragonal (SnO2), orthorhombic (Iodine), monoclinic (Sugar), hexagonal (Graphite), rhombohedral (Bi), and triclinic (H3BO3).
How are crystalline solids classified based on the dimensions of unit cells?
Crystalline solids are classified into seven systems based on the dimensions of unit cells.
What are the four types of crystalline solids based on the type of bonds present in them?
The four types of crystalline solids based on the type of bonds are:
(i) Ionic solids
(ii) Covalent solids
(iii) Metallic solids
(iv) Molecular solids
What are ionic solids, and how are they held together?
Ionic solids are crystalline solids where the particles are positively and negatively charged ions, held together by strong electrostatic forces of attraction, also known as ionic bonds.
Give examples of ionic solids.
Examples of ionic solids include NaCl (sodium chloride) and KBr (potassium bromide).
What are some properties of ionic solids?
Ionic solids have properties such as high stability, high melting and boiling points, low volatility, non-conductivity in the solid state (but conductivity in solution or when molten), brittleness, and high density.
Why do ionic solids not exist as individual neutral independent molecules?
Ionic solids do not exist as individual neutral independent molecules because they consist of cations and anions held together by electrostatic forces, and these forces are non-directional. They form a crystal lattice structure.
How does the structure of ionic crystals depend on the radius ratio of cations and anions?
The structure of ionic crystals depends on the radius ratio of cations and anions. For example, NaCl and CsF have the same geometry because they have the same radius ratio.
When do ionic crystals conduct electricity?
Ionic crystals conduct electricity when they are in solution or in the molten state, as in these conditions, the ions become free to move.
Why are ionic solids highly brittle?
Ionic solids are highly brittle because they consist of parallel layers of cations and anions in alternate positions. When an external force is applied, the layers slide, causing like ions to repel each other, leading to brittleness.
What is the formula unit used to describe ionic solids?
The formula unit is used to describe ionic solids instead of the molecular mass, as they do not exist in the form of molecules.
What is lattice energy in the context of ionic solids?
Lattice energy is the energy released when one mole of an ionic crystal is formed from gaseous ions or the energy required to break one mole of a solid into isolated ions in the gas phase. It is expressed in kJ mole^-1.
How does the size of cations and anions affect lattice energy?
Lattice energy decreases with the increase in the size of cations while keeping the anions the same size. It also decreases with the increase in the size of anions. This is because larger ions result in less tightly packed oppositely charged ions in the crystal lattice.
What are the properties of covalent crystals?
Covalent crystals have an open structure, are very hard, have high melting points, and low volatility. They are poor conductors of electricity due to the absence of free electrons and ions, except for materials like graphite that have a layered structure with delocalized electrons.
Which type of solvents are covalent crystalline solids typically soluble in?
Covalent crystalline solids are typically insoluble in polar solvents like water but readily soluble in non-polar solvents like benzene and carbon tetrachloride. However, covalent crystals with giant molecules like diamond and silicon carbide are insoluble in all solvents due to their large size.
What is the structure of diamond?
Diamond has a tetrahedral structure where each carbon atom forms four covalent bonds with other carbon atoms. The bonds run through the crystal in three dimensions, with bond angles of 109.5° and bond lengths of 154 pm. Diamond is a continuous, three-dimensional carbon lattice and is often described as a giant macro-molecule.
What are molecular solids, and what kinds of particles can form their crystals?
Molecular solids are solid substances where the particles forming the crystals are polar or non-polar molecules or atoms.
What are the two types of intermolecular forces that hold molecular solids together?
(i) Dipole-dipole interactions.
(ii) van der Waals forces.
How do the intermolecular forces in molecular solids compare to the forces of attraction in ionic and covalent crystals?
The intermolecular forces in molecular solids are much weaker than the forces of attraction between cations and anions in ionic crystals and between atoms in covalent crystals.
Give examples of molecular crystals with polar molecules and non-polar molecules.
Polar molecular solids: Ice and sugar.
Non-polar molecular solids: Iodine, sulfur, phosphorus, and carbon dioxide.
Why do polar molecular solids generally have higher melting and boiling points compared to non-polar molecular solids?
Polar molecular solids have stronger intermolecular forces due to dipole-dipole interactions, which require more energy to break, leading to higher melting and boiling points.
What is X-ray analysis used for in the study of molecular solids?
X-ray analysis helps determine the regular arrangements of atoms in constituent molecules of molecular solids and provides the exact positions of all the atoms.
What are the characteristics of molecular solids in terms of softness, compressibility, volatility, melting and boiling points, electrical conductivity, density, and solubility?
- Molecular solids are soft and easily compressible.
- They are mostly volatile with low melting and boiling points.
- They are bad conductors of electricity.
- They have low densities.
- Polar molecular crystals are mostly soluble in polar solvents, while non-polar molecular crystals are usually soluble in non-polar solvents.
Describe the structure of solid iodine and its electrical conductivity.
In the solid state, iodine molecules align in a layer lattice structure.
I-I bond distance is 271.5 pm, longer than in gaseous iodine (266.6 pm).
Iodine is a poor conductor of electricity.
What are the theories proposed to explain metallic bonding in metallic solids?
Electron pool or electron gas theory.
Valence bond theory.
Molecular orbital theory (band theory).
Why are metals good conductors of electricity?
Metals are good conductors of electricity because mobile electrons within the metal lattice can move freely in response to an applied electric field, allowing them to carry electrical current.
How does an increase in temperature affect the electrical conductivity of metals?
Increasing temperature can decrease the electrical conductivity of metals because as the temperature rises, the positive metal ions in the lattice begin to oscillate, hindering the free movement of mobile electrons and reducing electrical conductivity.
What is the relationship between thermal conductivity and metals?
Metals have high thermal conductivity because when a metal is heated at one end, mobile electrons absorb heat energy and rapidly transfer it through the metallic lattice to the cooler end by colliding with adjacent electrons.
Why do freshly cut metals often have a metallic luster?
Freshly cut metals have a metallic luster because when light falls on their surfaces, it excites the mobile electrons. These excited electrons, when de-excited, emit energy in the form of light, giving the metal a shiny appearance.
What happens to the structure of metals when stress is applied to them?
When stress is applied to metals, their layers can slip past each other, and the metal changes its shape without fracturing. This property makes metals malleable and ductile.
How are metal atoms arranged in a crystal lattice?
Metal atoms are arranged in a definite pattern in a crystal lattice, with free electrons roaming within this lattice. Metals can be thought of as assemblies of positively charged spheres (atoms) of identical radii closely packed together.
What is the concept of closed packing of atoms in metal structures?
The concept of closed packing involves arranging hard spherical balls (representing metal atoms) in a way that maximizes their proximity and stability. When these balls are shaken and rearranged, they tend to form closely packed arrangements.
How are various unit cells of the crystal lattice in metals developed?
Various unit cells of the crystal lattice in metals are developed by arranging metal atoms in layers. For example, in a tetrahedral structure, three balls join together in one plane, and the fourth ball is inserted in the space created by the other three as a second layer. This process continues to build the crystal lattice.
What is the pattern of arrangement when atoms of the third layer fit into interstices marked ‘b’ in CCP?
The pattern of arrangement is called ABC ABC or 123 123, also known as face-centered cubic arrangement.
In CCP, where are the balls of the fourth, seventh, and tenth layers located in relation to each other?
The balls of the fourth, seventh, and tenth layers are in front of each other.
How are the atoms of the third layer arranged in HCP?
In HCP, the atoms of the third layer occupy the depressions created by the second layer, forming an ABAB pattern.
Where are the balls of the third, fifth, and seventh layers located in relation to each other in HCP?
The balls of the third, fifth, and seventh layers are in front of each other in HCP.
What type of particles are present in metallic crystals, and what are their typical properties?
Metallic crystals consist of metallic cations and delocalized electrons. They have varying hardness, high melting points, lustrous appearance, ductility, malleability, and excellent conductivity of heat and electricity.
What are the structural particles and intermolecular forces in ionic crystals?
Ionic crystals consist of cations and anions with electrostatic attractions between them. They are hard with moderate to very high melting points and are nonconductors of electricity in the solid state.
What are molecular crystals composed of, and what are their typical properties?
Molecular crystals are composed of molecules held together by London dispersion forces, dipole-dipole forces, and hydrogen bonds. They are soft, have low melting points, do not conduct heat or electricity, and can sublime easily.
What type of bonds exist in network covalent crystals, and what are their properties?
Network covalent crystals consist of atoms bonded by covalent bonds. They are very hard, have very high melting points, and do not conduct electricity.
How can Avogadro’s number be determined, and what is it based on?
Avogadro’s number can be determined by studying crystalline solids. It is based on knowing the volume of one gram-mole of a crystalline solid and the distance between its atoms or ions in the crystal lattice.
In the example involving LiF, how is the Avogadro’s number calculated?
Avogadro’s number is calculated using the density of LiF, its molar mass, and the distance between Li+ and F- ions in the crystal lattice. The calculated value is approximately 6.02×10^23.
Long Questions Chemistry 1st Year Chapter 4 Liquids And Solids
Question: Explain the following with reasons.
(i) In the hydrogen bonded structure of HF, which is the stronger bond: the shorter covalent bond or the longer hydrogen bond between different molecules.
Answer: In the hydrogen-bonded structure of HF, the covalent bond between the hydrogen (H) and fluorine (F) within the same molecule is stronger than the hydrogen bond between different molecules. The covalent bond is a result of the sharing of electrons between the H and F atoms, which creates a strong, localized bond. In contrast, a hydrogen bond is an intermolecular force between the partially positive hydrogen atom of one molecule and the partially negative atom (usually nitrogen, oxygen, or fluorine) of another molecule. Hydrogen bonds are weaker than covalent bonds.
(ii) In a very cold winter the fish in garden ponds owe their lives to hydrogen bonding?
In very cold winters, the survival of fish in garden ponds can be attributed to the unique properties of water and its ability to form hydrogen bonds. When water freezes, it forms a crystalline structure where water molecules arrange themselves in a hexagonal lattice due to hydrogen bonding. This lattice structure makes ice less dense than liquid water, causing it to float.
Because ice is less dense, it forms an insulating layer on the surface of ponds, which helps in maintaining a relatively stable temperature beneath the ice. This insulation prevents the entire body of water from freezing solid, allowing fish and other aquatic organisms to survive in the slightly warmer liquid water beneath the ice. Without the expansion and buoyancy effects of hydrogen bonding in water, ponds would freeze from the bottom up, making it extremely challenging for aquatic life to survive.
(iii) Water and ethanol can mix easily and in all proportions.
Water and ethanol can mix easily and in all proportions because they are both polar molecules that can form hydrogen bonds with each other. Water molecules have a strong tendency to form hydrogen bonds due to the presence of polar O-H bonds. Ethanol (C2H5OH) also contains an -OH group, making it capable of hydrogen bonding.
When water and ethanol are mixed, the hydrogen bonds between water molecules and ethanol molecules can overcome the intermolecular forces within each liquid. This leads to the formation of a homogeneous solution, with the water and ethanol molecules interspersed throughout. The ability of these two liquids to mix easily and in all proportions is due to the favorable interactions between their polar molecules.
(iv) The origin of the intermolecular forces in water.
The intermolecular forces in water primarily arise from three sources:
Hydrogen Bonding: The most significant intermolecular force in water is hydrogen bonding. In a water molecule (H2O), the oxygen atom is more electronegative than the hydrogen atoms, leading to a partial negative charge on the oxygen atom and partial positive charges on the hydrogen atoms. This results in the formation of hydrogen bonds between water molecules. The hydrogen bond is an electrostatic attraction between the partially positive hydrogen atom of one water molecule and the partially negative oxygen atom of another water molecule.
Dipole-Dipole Interactions: Water is a polar molecule, and polar molecules have permanent dipole moments. The partial positive and negative charges on the hydrogen and oxygen atoms in water lead to dipole-dipole interactions between neighboring water molecules.
London Dispersion Forces: Although water is a polar molecule, it also experiences London dispersion forces (van der Waals forces) due to the constant motion of electrons. These temporary fluctuations in electron distribution can result in short-lived, instantaneous dipoles, leading to weak attractions between water molecules.
Question: Briefly consider some of the efects on our lives if water has only a very weak hydrogen bonding present among its molecules.
Answer: No Liquid Water at Room Temperature: Water would not exist in its liquid form at room temperature and pressure. Instead, it would remain as a gas (vapor) until much lower temperatures were reached. This would make it extremely challenging for life as we know it to exist, as liquid water is essential for various biological processes and the habitability of our planet.
Altered Climate and Weather Patterns: The absence of liquid water would drastically alter climate and weather patterns on Earth. Water vapor plays a crucial role in the Earth’s climate system, and without it, temperature regulation and precipitation cycles would be disrupted.
Lack of Hydrogen Bonding Properties: Many unique properties of water, such as its high heat capacity, surface tension, and solvent capabilities, are a result of strong hydrogen bonding. Weak hydrogen bonding would lead to the loss of these essential properties, affecting various industrial, biological, and environmental processes.
Impact on Biological Systems: Weak hydrogen bonding in water would impact biological systems, as hydrogen bonding is essential for the structure and function of biomolecules like DNA, proteins, and enzymes. Life as we know it relies on the specific properties of water’s hydrogen bonding.
Environmental Consequences: Natural habitats and ecosystems that depend on the availability of liquid water would be severely affected. Aquatic life, including fish and other aquatic organisms, would face challenges, and wetlands, rivers, and lakes might not exist in the same form.
Question: All gases have a characteristic critical temperature. Above the critical temperature it is impossible to liquefy a gas. The critical temperatures of carbon dioxide and methane are
31.14 0C and -81.9 0C, respectively. Which gas has the stronger intermolecular forces? Briely
explain your choice?
Answer: Comparing the Intermolecular Forces in Carbon Dioxide and Methane:
The critical temperature of a gas is a measure of the strength of its intermolecular forces. The higher the critical temperature, the stronger the intermolecular forces.
Carbon Dioxide (CO2)
Critical Temperature: 31.14°C
Intermolecular Forces: Carbon dioxide molecules are held together by dispersion forces (van der Waals forces) and dipole-dipole interactions. While it doesn’t form hydrogen bonds, CO2 is a polar molecule due to the electronegativity difference between carbon and oxygen.
Methane (CH4)
Critical Temperature: -81.9°C
Intermolecular Forces: Methane molecules are primarily held together by dispersion forces. It is a nonpolar molecule since the carbon-hydrogen bonds are relatively nonpolar and there are no permanent dipoles.
Explanation:
Carbon dioxide (CO2) has a significantly higher critical temperature (31.14°C) compared to methane (-81.9°C), indicating stronger intermolecular forces. Although both gases rely on dispersion forces, CO2’s polarity due to its asymmetric molecular structure contributes to stronger dipole-dipole interactions compared to the nonpolar methane. However, neither of these gases forms hydrogen bonds like water does.
Question: Explain the term saturated vapour pressure. Arrange in order of increasing vapour pressure: 1dm3 water, 1 dm3 ethanol, 50 cm3 water, 50 cm3 ethanol and 50 cm3 of ether.
Answer: Saturated Vapor Pressure
Saturated vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid phase at a specific temperature. It represents the maximum pressure that a vapor can exert in a closed system at that temperature before condensation (liquid to vapor) or vaporization (vapor to liquid) reaches an equilibrium. It is a characteristic property of a substance at a given temperature and is independent of the volume of the container.
Now, let’s arrange the given volumes of substances in order of increasing vapor pressure:
50 cm^3 of water: This is a small volume of water, and the vapor pressure will be relatively low.
50 cm^3 of ethanol: Ethanol has a higher vapor pressure compared to water due to weaker intermolecular forces.
50 cm^3 of ether: Ether has even weaker intermolecular forces compared to ethanol, so it will have a higher vapor pressure than ethanol at the same temperature.
1 dm^3 of ethanol: This is a larger volume of ethanol compared to the previous cases, but the vapor pressure will still be lower than the smaller volumes of ethanol due to the same temperature.
1 dm^3 of water: This is a larger volume of water compared to the previous cases, but the vapor pressure will still be lower than the smaller volumes of water due to the same temperature.
So, in increasing order of vapor pressure, it would be:
50 cm^3 of water < 1 dm^3 of water < 1 dm^3 of ethanol < 50 cm^3 of ethanol < 50 cm^3 of ether.
Question: While a volatile liquid standing in a breaker evaporates, the temperature of the liquid remains the same as that of its surrounding. If the same liquid is allowed to vapourize into atmosphere in an insulated vessel, its temperature falls below that of its surrounding. Explain the difference in behaviour.
Answer: The difference in behavior between a volatile liquid evaporating in an open container and vaporizing into an insulated vessel lies in the process of energy exchange and the nature of the surroundings. This behavior can be explained by the principles of thermodynamics.
- Evaporation in an Open Container (Non-Insulated)
When a volatile liquid evaporates in an open container that is not insulated, the liquid’s temperature remains the same as that of its surroundings. Here’s why:
As the volatile liquid molecules gain enough kinetic energy from their surroundings, some of them at the liquid’s surface break the attractive forces (intermolecular forces) holding them in the liquid phase and become vapor molecules.
This process requires energy in the form of heat. The energy is supplied by the surroundings as some of the molecules in the surroundings collide with the liquid’s surface molecules and transfer energy to them.
As a result, the temperature of the liquid remains constant because the heat energy input from the surroundings compensates for the energy lost during vaporization. Essentially, the heat energy absorbed by the liquid during vaporization is taken from the surroundings, so the liquid doesn’t experience a change in temperature.
- Vaporization into an Insulated Vessel
When the same volatile liquid is allowed to vaporize into an insulated vessel, the temperature of the liquid falls below that of its surroundings. Here’s why:
In an insulated vessel, there is minimal heat exchange with the surroundings. The insulation prevents heat from flowing into or out of the vessel.
As the volatile liquid molecules vaporize and escape into the vessel’s interior, they are essentially removing energy (in the form of heat) from the liquid and its surroundings. This is because vaporization requires energy, and in this case, there is no external source of heat to replenish it.
As more and more molecules vaporize, the liquid loses energy, causing its temperature to drop. This process continues until the vaporization reaches an equilibrium, and the temperature inside the insulated vessel becomes lower than the temperature of the surroundings.
Question: How does hydrogen bonding explain the following indicated properties of the substances?
(i) Structure of DNA
Explanation: Hydrogen bonding is essential in explaining the structure of DNA (deoxyribonucleic acid). DNA is composed of two long chains of nucleotides that form a double helix structure. The hydrogen bonds occur between complementary nitrogenous bases: adenine (A) forms hydrogen bonds with thymine (T), and guanine (G) forms hydrogen bonds with cytosine (C). These hydrogen bonds between base pairs help hold the two DNA strands together in a stable and specific manner. This complementarity and hydrogen bonding between the base pairs are crucial for DNA replication, transcription, and genetic information storage.
(ii) Hydrogen bonding in proteins
Explanation: Proteins are composed of amino acids, and their structure and function are highly dependent on hydrogen bonding. In proteins, hydrogen bonds form between different amino acids and within the backbone of the protein molecule. These hydrogen bonds help stabilize the secondary and tertiary structures of proteins.
In protein secondary structures like alpha helices and beta sheets, hydrogen bonds form between the amino acid residues in the polypeptide chain. These interactions contribute to the folded and compact structure of proteins.
In protein folding and three-dimensional structure (tertiary structure), hydrogen bonds between amino acid side chains and polar groups help maintain the specific conformation of the protein, which is essential for its function.
(iii) Formation of ice and its lesser density than liquid water
Explanation: Hydrogen bonding also plays a role in the formation of ice and its lower density compared to liquid water. In the liquid state, water molecules are in constant motion, forming and breaking hydrogen bonds. However, as water cools and freezes into ice, the hydrogen bonds become more stable and organized.
In the solid state (ice), water molecules arrange themselves in a hexagonal lattice structure due to hydrogen bonding. This structure results in open spaces between water molecules, making ice less dense than liquid water.
Hydrogen bonding in ice holds water molecules farther apart in a regular pattern, which leads to a lower density. In contrast, in liquid water, the molecules are closely packed due to the constant breaking and reforming of hydrogen bonds, resulting in higher density.
(iv) Solubilities of compounds
Explanation: Hydrogen bonding is a critical factor in the solubility of various compounds. It explains why some compounds are more soluble in water than others.
Hydrogen bonding in water molecules allows water to effectively solvate polar and ionic compounds. Water’s ability to form hydrogen bonds with solute molecules can disrupt the attractive forces within the solute and facilitate its dissolution.
Substances with polar or ionic groups that can engage in hydrogen bonding, such as alcohols, organic acids, and some salts, tend to be highly soluble in water because they can form favorable interactions with water molecules.
In contrast, nonpolar compounds, such as hydrocarbons, do not readily engage in hydrogen bonding with water and are typically insoluble or only sparingly soluble in water.
Question: What are liquid crystals? Give their uses in daily life.
Answer: Liquid crystals are a distinct phase of matter that exhibits properties of both liquids and crystalline solids. They have unique structural characteristics and can flow like liquids while maintaining some degree of order and alignment of their constituent molecules like crystalline solids. Liquid crystals find various applications in daily life due to their optical, electrical, and structural properties. Here are some uses of liquid crystals in daily life:
Liquid Crystal Displays (LCDs): LCDs are perhaps the most well-known application of liquid crystals. They are used in devices like televisions, computer monitors, smartphones, tablets, digital watches, and calculators. Liquid crystals are sandwiched between two glass plates and manipulated using electrical currents to control the passage of light, resulting in the display of images and information.
Thermometers: Liquid crystal thermometers are used in medical applications and everyday household thermometers. The change in the alignment of liquid crystal molecules in response to temperature variations is used to indicate the temperature on a scale.
Mood Rings: Mood rings contain liquid crystals that change color based on the wearer’s body temperature. These rings were popular in the 1970s and are often used for entertainment and fashion accessories.
Room Thermometers and Displays: Some room temperature displays use liquid crystals to provide information about the temperature of a space. These can be useful for maintaining indoor comfort levels.
Wristwatches: Liquid crystal displays are used in digital wristwatches to show the time and other information. These watches are common and convenient for everyday use.
Home and Car Thermometers: Liquid crystal thermometers are often used in homes and vehicles to monitor indoor and outdoor temperatures.
Liquid Crystal Light Modulators: Liquid crystals are used in various optical devices, including light modulators for adjusting the intensity of laser beams. These are essential in scientific and industrial applications.
Polarizing Filters: Liquid crystals are used in polarizing filters for sunglasses, camera lenses, and LCD screens to reduce glare and enhance visibility in bright conditions.
Liquid Crystal Temperature Labels: These labels are used on food packaging to indicate whether a product has been subjected to excessive heat during transport or storage. They change color irreversibly if the temperature exceeds a certain threshold.
Liquid Crystal Privacy Filters: Some laptop screens and smartphone screen protectors use liquid crystal layers to provide privacy by restricting the viewing angle to the user directly in front of the screen.
Liquid Crystal Light Valves: These devices are used in optical systems to control the intensity of light, making them valuable in scientific research, laser applications, and photography.
Alignment Layers in Optics: Liquid crystals can serve as alignment layers in optical devices to control the orientation of other materials like liquid crystal polymers, affecting the passage of light.
Question: Explain the following with reasons.
(i) Evaporation causes cooling.
Explanation: Evaporation causes cooling because it is an endothermic process. During evaporation, molecules with higher kinetic energy at the liquid’s surface escape into the vapor phase, leaving behind molecules with lower kinetic energy. This process requires energy to overcome intermolecular forces and change the phase from liquid to vapor. The energy is taken from the surroundings, leading to a decrease in the average kinetic energy (and thus temperature) of the remaining liquid molecules, causing cooling.
(ii) Evaporation takes place at all temperatures.
Explanation: Evaporation can occur at all temperatures, not just at the boiling point. It happens because the molecules in a liquid possess a range of kinetic energies, and some of them have enough energy to overcome intermolecular forces and transition to the vapor phase. Even at temperatures below the boiling point, molecules with sufficient energy can evaporate, albeit at a slower rate than at higher temperatures.
(iii) Boiling needs a constant supply of heat.
Explanation: Boiling requires a constant supply of heat because it is an endothermic phase transition. To convert a liquid into a gas (vapor) during boiling, the molecules within the liquid must overcome not only intermolecular forces but also increase their kinetic energy significantly. This extra energy is continuously supplied in the form of heat to maintain the boiling process.
(iv) Earthenware vessels keep water cool.
Explanation: Earthenware vessels keep water cool due to a process called evaporative cooling. The porous nature of earthenware allows water to slowly seep through the vessel’s walls and evaporate from the outer surface. This evaporation process removes heat from the water inside the vessel, leading to a cooling effect. It’s similar to how sweating cools our bodies by evaporating moisture from our skin.
(v) One feels sense of cooling under the fan after bath.
Explanation: One feels a sense of cooling under a fan after a bath because the moving air from the fan enhances the evaporation of moisture on the skin. As the moisture evaporates, it absorbs heat from the skin, creating a cooling effect and making us feel cooler and more comfortable.
(vi) Dynamic equilibrium is established during evaporation of a liquid in a closed vessel at constant temperature.
Explanation: When a liquid undergoes evaporation in a closed vessel at a constant temperature, a dynamic equilibrium is established. At this equilibrium, the rate of condensation of vapor molecules back into the liquid phase equals the rate of evaporation. This dynamic balance occurs because, despite being in a closed system, there are always some molecules with enough kinetic energy to evaporate, and others return to the liquid phase due to collisions with the liquid’s surface. The temperature remains constant because the heat absorbed during evaporation is released when the vapor condenses.
(vii) The boiling point of water is different at Murree hills and at Mount Everest.
Explanation: The boiling point of water is lower at higher altitudes like Mount Everest compared to lower altitudes like Murree hills. This difference is due to the decrease in atmospheric pressure at higher altitudes. Lower atmospheric pressure reduces the pressure exerted on the liquid, which means the liquid’s vapor pressure can match atmospheric pressure at a lower temperature, resulting in a lower boiling point.
(viii) Vacuum distillation can be used to avoid decomposition of a sensitive liquid.
Explanation: Vacuum distillation is a technique used to distill heat-sensitive or volatile liquids at lower temperatures. By reducing the pressure in the distillation apparatus, the liquid can boil at a lower temperature than its normal boiling point. This helps avoid the decomposition or overheating of sensitive compounds that would occur under normal atmospheric pressure and higher temperatures.
(ix) Heat of sublimation of a substance is greater than its heat of vaporization.
Explanation: The heat of sublimation of a substance (the energy required to change a solid directly into a vapor) is generally greater than its heat of vaporization (the energy required to change a liquid into a vapor). This is because sublimation involves breaking the solid’s intermolecular forces and converting it into vapor, whereas vaporization only involves breaking the intermolecular forces in the liquid phase.
(x) Heat of sublimation of iodine is very high.
Explanation: The heat of sublimation of iodine is relatively high because iodine molecules are held together by strong van der Waals forces (London dispersion forces) in the solid phase. Subliming iodine involves breaking these strong intermolecular forces and converting the solid iodine directly into gaseous iodine molecules. As a result, a significant amount of energy is required for sublimation, leading to a high heat of sublimation for iodine.
Question: Explain the following properties of crystalline solids. Give three examples in each case.
(i) Anisotropy
Explanation: Anisotropy refers to the property of crystalline solids having different physical properties in different crystallographic directions. This occurs because the arrangement of atoms or molecules in crystals varies along different crystal axes. Examples:
Quartz (SiO2): Quartz is an anisotropic mineral used in electronics. It exhibits variations in electrical conductivity, thermal expansion, and refractive index depending on the direction of measurement.
Calcite (CaCO3): Calcite is an anisotropic mineral that shows double refraction, meaning it splits incident light into two different rays with different velocities and directions.
Mica (e.g., muscovite): Mica minerals have perfect basal cleavage along specific crystallographic planes, making them anisotropic in terms of their cleavage properties.
(ii) Cleavage
Explanation: Cleavage refers to the tendency of crystalline solids to break along specific planes or directions, resulting in smooth, flat surfaces. Cleavage is a characteristic property of minerals. Examples:
Diamond (C): Diamond exhibits cleavage along four octahedral planes, giving it a characteristic octahedral shape when broken.
Halite (NaCl): Halite, or rock salt, exhibits perfect cubic cleavage, breaking into cubes.
Muscovite (KAl2(AlSi3O10)(OH)2): Muscovite mica shows perfect basal cleavage along the (001) plane, resulting in thin, flexible sheets.
(iii) Habit of a crystal
Explanation: The habit of a crystal refers to its characteristic external shape or form, which is determined by its internal atomic arrangement. Examples:
Pyrite (FeS2): Pyrite crystals often have a cubic habit with distinct faces and sharp edges.
Bismuth (Bi): Bismuth crystals have a stair-stepped, hopper-like habit with distinctive step-like terraces.
Garnet (e.g., almandine): Garnet crystals may exhibit dodecahedral or trapezohedral habits, with 12 or 24 faces, respectively.
(iv) Isomorphism
Explanation: Isomorphism is a phenomenon where different crystalline solids have similar crystal structures and chemical compositions, allowing the substitution of one element by another within the crystal lattice. Examples:
Olivine group minerals (e.g., forsterite and fayalite): These minerals form a solid solution series, where magnesium (Mg) can substitute for iron (Fe) in the crystal structure.
Feldspar group minerals (e.g., albite and anorthite): Different feldspars within this group can have varying ratios of sodium (Na) and calcium (Ca) ions within their crystal structures.
Garnet group minerals (e.g., almandine and pyrope): Garnets can contain different combinations of aluminum (Al), iron (Fe), and magnesium (Mg) ions in their structures.
(v) Polymorphism
Explanation: Polymorphism is the property of a substance to exist in multiple crystalline forms (polymorphs) with distinct crystal structures but the same chemical composition. Examples:
Carbon (C): Carbon exhibits polymorphism, with diamond and graphite being two well-known polymorphs. Diamond has a tetrahedral crystal structure, while graphite has a layered hexagonal lattice.
Sulfur (S): Sulfur has several polymorphs, including rhombic sulfur and monoclinic sulfur, each with a different arrangement of sulfur atoms.
Silicon dioxide (SiO2): Silicon dioxide can exist as quartz, cristobalite, tridymite, and coesite, each having a different crystal structure.
Transition Temperature
Explanation: Transition temperature refers to the temperature at which a crystalline solid undergoes a phase transition, changing its physical properties. Examples:
Graphite to diamond transition in carbon: Occurs at extremely high pressures and temperatures.
Ferromagnetic to paramagnetic transition in iron: Occurs at the Curie temperature (about 770°C for iron).
Quartz to beta-quartz transition: Quartz undergoes a reversible phase transition at high temperatures (around 573°C) to form beta-quartz with different crystal properties.
(vii) Symmetry
Explanation: Symmetry in crystalline solids refers to the repetitive and orderly arrangement of atoms or molecules within the crystal lattice. Crystals can exhibit various symmetry elements, including rotational axes, mirror planes, and inversion centers. Examples:
Sodium chloride (NaCl): Common table salt has a cubic crystal lattice with three mutually perpendicular mirror planes, demonstrating cubic symmetry.
Calcite (CaCO3): Calcite crystals often show rhombohedral symmetry with a threefold rotational axis and a single mirror plane.
Quartz (SiO2): Quartz crystals exhibit hexagonal symmetry, featuring sixfold rotational symmetry.
(viii) Growing of a Crystal
Explanation: The growth of a crystal involves the gradual addition of molecules or ions to the crystal’s surface, allowing it to expand in a controlled manner. Examples:
Growing sugar crystals: Sugar can be dissolved in water and allowed to crystallize as it cools, forming edible sugar crystals.
Growing alum crystals: Potassium aluminum sulfate (alum) can be dissolved in water, and as the solution cools, alum crystals can be grown in a laboratory setting.
Growing protein crystals: In biotechnology and structural biology, proteins can be crystallized to study their atomic structures. The process involves carefully controlling the conditions to promote crystal growth.
(b) How polymorphism and allotropy are related to each other? Give examples.
Polymorphism and allotropy are related terms that both refer to the existence of multiple crystallographic forms of a substance. While the terms are often used interchangeably, they are typically associated with different types of materials: polymorphism is more commonly used for compounds, while allotropy is used for elements. Nevertheless, the underlying concept of different crystal structures for the same substance applies to both.
Polymorphism
Polymorphism refers to the property of a substance to exist in multiple crystalline forms (polymorphs) with the same chemical composition but different crystal structures. Polymorphism is commonly observed in chemical compounds. Examples include:
Carbon Dioxide (CO2): Carbon dioxide exhibits polymorphism. At standard conditions, it exists as a gas. However, at high pressures, it can form solid polymorphs with distinct crystal structures, such as dry ice (solid CO2) and a high-pressure phase known as Phase II.
Sulfur (S): Sulfur is known to exhibit polymorphism. It has multiple solid allotropes with varying structures and properties. Common forms include rhombic sulfur and monoclinic sulfur.
Phenanthrene: This organic compound can exist in several polymorphic forms, including alpha, beta, and gamma phenanthrene, each with different crystal structures.
Allotropy
Allotropy is a term primarily used for elements and refers to the property of an element to exist in multiple structural forms in the same physical state (solid, liquid, or gas) under different conditions of temperature and pressure. Examples include:
Carbon (C): Carbon is a classic example of an element that exhibits allotropy. Depending on the crystal structure, carbon can exist as diamond (a hard, transparent crystal lattice), graphite (a layered hexagonal lattice), or in various other forms like amorphous carbon.
Oxygen (O2): Oxygen exists as dioxygen (O2) in the gas phase and as ozone (O3) in the gas phase, each having a distinct molecular arrangement.
Phosphorus (P): White phosphorus and red phosphorus are well-known allotropes of phosphorus. White phosphorus has a tetrahedral arrangement of P4 molecules, while red phosphorus has a polymeric structure.
Question: (a) Deine unit cell. What are unit cell dimensions? How the idea of crystal lattice is developed from the concept of unit cell?
Answer: A unit cell is the smallest repeating structural unit in a crystalline solid that, when repeated in three dimensions, generates the entire crystal lattice. It is a fundamental concept in the study of crystallography and solid-state physics. The unit cell represents the arrangement of atoms, ions, or molecules within a crystal.
Unit Cell Dimensions
Unit cell dimensions are the parameters that define the size and shape of a unit cell. These dimensions are typically described using three edge lengths (a, b, and c) and three interaxial angles (α, β, and γ). The choice of unit cell and its dimensions depends on the specific crystal system to which the crystal belongs.
Development of Crystal Lattice from Unit Cell
The concept of a crystal lattice is developed from the idea of a unit cell by imagining that this smallest repeating unit is extended in all three dimensions throughout space to fill the entire crystal structure. When you replicate the unit cell in three directions (along the crystallographic axes), it generates a three-dimensional network of points that represent the positions of atoms, ions, or molecules in the crystal lattice. This lattice of points forms a repeating pattern, which is characteristic of the crystal structure.
In summary, the crystal lattice is a three-dimensional arrangement of points representing the positions of constituent particles in the crystal, and this lattice is constructed by repeating the unit cell in all three dimensions.
(b) Explain seven crystal systems and draw the shapes of their unit cells.
Seven Crystal Systems and Unit Cell Shapes
There are seven crystal systems, each characterized by its own unit cell shape and symmetry. Here are the seven crystal systems along with the shapes of their unit cells:
Cubic (Isometric) System:
Unit Cell Shape: Cube (all edges are equal in length, and all angles are 90 degrees).
Examples: Sodium chloride (NaCl), diamond (C), and body-centered cubic (BCC) iron (Fe).
Tetragonal System:
Unit Cell Shape: Rectangular prism (two edges are equal in length, and all angles are 90 degrees).
Examples: Zirconium dioxide (ZrO2), tin (Sn), and rutile (TiO2).
Orthorhombic System:
Unit Cell Shape: Rectangular prism (all edges have different lengths, and all angles are 90 degrees).
Examples: Barite (BaSO4), olivine ((Mg,Fe)2SiO4), and sulfur (S).
Hexagonal System:
Unit Cell Shape: Hexagonal prism (all edges have two equal lengths, and angles are 90 degrees except for the vertical axis, which is 120 degrees).
Examples: Graphite (C), beryl (Be3Al2(SiO3)6), and quartz (SiO2).
Rhombohedral (Trigonal) System:
Unit Cell Shape: Rhombohedron (all edges have equal lengths, and all angles are not 90 degrees; they are typically 60 degrees).
Examples: Calcite (CaCO3), rhodium (Rh), and cinnabar (HgS).
Monoclinic System:
Unit Cell Shape: Parallelepiped (all edges have different lengths, and only one angle is 90 degrees).
Examples: Gypsum (CaSO4•2H2O), stibnite (Sb2S3), and orthoclase (KAlSi3O8).
Triclinic System:
Unit Cell Shape: Parallelepiped (all edges have different lengths, and none of the angles are 90 degrees).
Examples: Microcline (KAlSi3O8), plagioclase feldspar ((Ca,Na)Al1-2Si3-2O8), and turquoise (CuAl6(PO4)4(OH)8•4H2O).
Question: (a) What are ionic solids? Give their properties. Explain the structure of NaCl. Sketch a model to justify that unit cell of NaCl has four formula units in it.
Answer: Ionic solids are a type of crystalline solid composed of positively charged ions (cations) and negatively charged ions (anions) held together by strong electrostatic forces (ionic bonds). These solids are typically formed by the combination of metal cations and non-metal anions. Some common properties of ionic solids are:
High melting and boiling points: Ionic solids have high melting and boiling points due to the strong electrostatic forces between ions, which require a significant amount of energy to break.
Brittle: Ionic solids are often brittle and tend to shatter when subjected to stress because the layers of ions can slide past each other only to a limited extent before the repulsive forces between like charges cause the crystal to break.
Solubility in water: Many ionic solids are soluble in water because water molecules can surround and separate the ions, allowing them to dissolve and form aqueous solutions.
Conductivity: Ionic solids are typically poor conductors of electricity in the solid state because the ions are fixed in a rigid lattice. However, they become good conductors when melted or dissolved in a solution.
Structure of NaCl (Sodium Chloride)
Sodium chloride (NaCl), or common table salt, is a well-known example of an ionic solid. Its crystal structure is a face-centered cubic (FCC) arrangement, where sodium ions (Na+) occupy the face-centered positions of the unit cell, and chloride ions (Cl-) occupy the octahedral positions between the sodium ions. Each sodium ion is surrounded by six chloride ions, and vice versa.
To justify that the unit cell of NaCl has four formula units in it, consider the arrangement:
Each sodium ion (Na+) at a corner of the unit cell is shared by eight neighboring unit cells.
Each chloride ion (Cl-) at the center of an edge of the unit cell is also shared by four neighboring unit cells.
By considering the contributions from corners and edges, you can determine that there are four formula units of NaCl within a single unit cell of the FCC structure.
(b) What are covalent solids? Give their properties. Explain the structure of diamond.
Covalent solids, also known as network solids, are crystalline solids composed of atoms connected by covalent bonds, forming an extensive three-dimensional network. These solids have unique properties, including:
High melting and boiling points: Covalent solids typically have high melting and boiling points because breaking the covalent bonds requires a substantial amount of energy.
Brittle: Similar to ionic solids, covalent solids are often brittle because the breaking of covalent bonds along specific planes or directions leads to fracture.
Poor electrical conductivity: Most covalent solids do not conduct electricity because there are no free ions or electrons to carry an electric current.
Structure of Diamond
Diamond is a well-known example of a covalent solid. In diamond, each carbon atom is bonded covalently to four neighboring carbon atoms, forming a tetrahedral network structure. This continuous network of covalent bonds extends throughout the crystal lattice, resulting in a three-dimensional, interconnected structure.
The strong covalent bonds between carbon atoms in diamond give it exceptional hardness, making it one of the hardest natural materials known. Its transparent and brilliant appearance is also due to its crystal structure, which allows light to be refracted and dispersed, resulting in its characteristic sparkle.
(c) What are molecular crystals? Give their properties. Justify that molecular crystals are softer than ionic crystals.
Molecular crystals are crystalline solids composed of discrete molecules held together by weak van der Waals forces (London dispersion forces), dipole-dipole interactions, or hydrogen bonding. Some properties of molecular crystals include:
Low melting and boiling points: Molecular crystals generally have lower melting and boiling points compared to ionic and covalent solids because the intermolecular forces holding the molecules together are weaker.
Softness: Molecular crystals are often soft and have lower hardness compared to ionic and covalent solids due to the weaker forces between molecules.
Solubility: Many molecular crystals are soluble in nonpolar or polar solvents, depending on the nature of the intermolecular forces present.
Electrical conductivity: Molecular crystals are typically poor conductors of electricity because the molecules are electrically neutral and not free to move charge.
The softness of molecular crystals is justified by the weak intermolecular forces between the molecules, which allow the crystal lattice to deform and shift more easily when subjected to stress compared to the stronger bonds in ionic and covalent solids.
Question: Give different theories of a metallic bond. How does electron sea theory justify the electrical conductivity, thermal conductivity and shining surfaces of metals?
Answer: here are several theories that attempt to explain the nature of metallic bonds in metals. One of the most widely accepted theories is the “Electron Sea Model” or “Electron Sea Theory.” This theory, along with some aspects of band theory, helps justify the electrical conductivity, thermal conductivity, and shining surfaces of metals.
Electron Sea Theory (Electron Sea Model)
The Electron Sea Theory describes the metallic bond as a sea of free electrons that move throughout the crystal lattice of metal cations. This model proposes the following features of metallic bonds:
Metal Cations: In a metallic crystal lattice, the metal atoms lose valence electrons to become positively charged metal cations. These cations are arranged in a regular, three-dimensional array.
Delocalized Electrons: The valence electrons from the metal atoms are released into a “sea” of delocalized electrons that are free to move throughout the entire crystal lattice. These electrons are not bound to any particular metal cation and are collectively shared among all the cations.
Electrical Conductivity
The Electron Sea Theory justifies the electrical conductivity of metals because of the presence of delocalized electrons. When a voltage is applied across a metal, these free electrons can drift in response to the electric field, creating an electric current. Since the electrons are abundant and mobile, metals are excellent conductors of electricity.
Thermal Conductivity
Thermal conductivity in metals can also be explained by the Electron Sea Theory. When heat is applied to a metal, the free electrons gain kinetic energy and move more rapidly. These energetic electrons collide with neighboring atoms and transfer thermal energy through the lattice. The high thermal conductivity of metals is attributed to this efficient transfer of kinetic energy by the free electrons.
Shining Surfaces (Luster)
The shiny, reflective surfaces of metals, known as their luster, can be explained by the Electron Sea Theory. When light strikes a metal surface, it encounters the sea of free electrons. These electrons interact with the incoming photons of light and quickly re-radiate them in various directions. This scattering of light by the free electrons creates a shiny appearance and accounts for the metallic luster.
In summary, the Electron Sea Theory provides a useful framework for understanding the electrical conductivity, thermal conductivity, and luster of metals. The presence of a sea of delocalized electrons allows metals to conduct electricity efficiently, transfer heat effectively, and reflect light, giving them their characteristic properties.
Question: Crystals of salts fracture easily but metals are deformed under stress without fracturing. Explain the difference.
Answer: The difference in the mechanical behavior of crystals of salts (ionic solids) and metals can be explained by their distinct atomic/molecular structures and the nature of their chemical bonds.
- Crystal Structure
Ionic Solids (Salts): Ionic solids are composed of positively charged metal cations and negatively charged non-metal anions, held together by strong electrostatic forces (ionic bonds). These ions are arranged in a repeating three-dimensional lattice. When stress is applied, the planes of ions can easily slide past each other along specific crystallographic planes. This limited ability to slip results in brittle behavior because when the planes of ions reach a certain point, the repulsive forces between like charges cause the crystal to break along planes of weakness, leading to fracture.
Metals: Metals have a metallic crystal structure, where metal atoms are arranged in a closely packed, regular array. The metallic bond that holds the atoms together is described by the “Electron Sea Model,” where valence electrons are free to move throughout the crystal lattice. When stress is applied to a metal, the atomic layers can slide past each other easily due to the mobility of these free electrons. This ability to deform plastically without breaking is known as ductility, and it allows metals to be malleable and capable of being shaped without fracturing.
- Bonding Nature
Ionic Bonds (Salts): Ionic bonds are strong electrostatic attractions between oppositely charged ions. These bonds are typically rigid, and when stress is applied, the crystal lattice tends to break along planes of weakness due to the inability of ions to move freely within the lattice.
Metallic Bonds (Metals): Metallic bonds are characterized by a “sea” of delocalized electrons shared among all metal atoms. This electron mobility allows the layers of metal atoms to slip past each other without breaking the bonds, even when subjected to stress. The metallic bond’s malleability and ductility properties enable metals to deform under stress without fracturing.