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Questions and Answers
What is the primary factor that leads to the liquefaction of gases?
What happens to the kinetic energy of gas molecules as the temperature decreases?
Which of the following statements about critical temperature is true?
Which of the following factors primarily affects the transition of gas to a liquid state?
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At room temperature, how do intermolecular forces of attraction in gases behave?
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What is a characteristic of amorphous solids?
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Which scenario describes an enantiotropic change?
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Which of the following is an example of a drug that exhibits polymorphism?
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What does the yield value refer to in the context of solids under pressure?
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Which definition best describes monotropic changes in polymorphism?
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Which of the following substances is NOT considered an amorphous solid?
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What type of polymorphism occurs when the transition between forms is irreversible?
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Which property is true regarding solvates?
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What defines pseudoplastic flow in terms of shear rate?
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Which of the following describes thixotropy?
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Which type of flow is characterized by a solid forming a gel readily when shaken?
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What is unique about antithixotropy?
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What aspect does NOT affect the rheological properties of liquids and semisolids?
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Which of the following pairs correctly matches crystal systems with their examples?
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What distinguishes crystalline solids from amorphous solids?
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What is a property of amorphous solids?
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Which of the following is NOT a factor affecting viscosity measurements?
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Which intermolecular forces are stronger in solids compared to liquids or gases?
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What is the effect of a positive heat of solution on solubility?
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What does the variable 'F' represent in Gibbs' phase rule equation?
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Which condition describes the dissolution process of strongly soluble electrolytes in an endothermic reaction?
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What happens to the solubility of sodium chloride when the temperature changes?
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In a dissolution reaction, which component is influenced by the common ion effect?
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Which of the following statements regarding the water-phenol system is true at 50°C?
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What is the significance of the solubility product, Ksp?
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What is the significance of the critical solution temperature?
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How does an increase in temperature affect the solubility of exothermic dissolution reactions?
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What is the outcome when a common ion forms a complex with a salt in a solution?
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What is eutectia?
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Which statement is true regarding the solubility of slightly soluble electrolytes in relation to their solubility product (Ksp)?
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Which of the following mixtures is an example of a eutectic point?
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What does thermochemistry primarily study?
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According to the first law of thermodynamics, what can be derived?
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In Gibbs' phase rule, what does 'C' represent?
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Study Notes
Intermolecular Forces in Gases
- Intermolecular forces of attraction are virtually nonexistent in gases at room temperature.
Liquefaction of Gases
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Factors affecting the transition of gas to liquid state:
- Temperature: Low temperatures reduce the kinetic energy of gas molecules, decreasing their velocity and causing them to transition to the liquid state.
- Critical Temperature: This is the temperature above which a liquid can no longer exist. For water, the critical temperature is 374°C or 647K.
- Pressure: Applying pressure brings gas molecules closer together, increasing the intermolecular forces of attraction and facilitating liquefaction.
Pseudoplastic Flow
- Also known as "shear-thinning systems."
- Viscosity decreases with an increase in shear rate.
- Examples include polymers in solution, such as natural and synthetic gums.
Types of Flow
- Thixotropy: Reversible sol-gel transformation, where viscosity decreases with time under constant stress.
- Antithixotropy: Negative thixotropy, where viscosity increases with time under constant stress.
- Rheopexy: Solid forms a gel readily when gently shaken.
Factors affecting Rheological Properties
- Temperature
- Shear Rate
- Pressure
- Time
- Measuring conditions
- Composition and additives
Solids
- Intermolecular forces of attraction are stronger than in liquids or gases.
- Solids can be crystalline or amorphous:
Crystalline Solids
- Fixed molecular order
- Distinct melting point
- Anisotropic (properties vary with direction)
- Examples:
- Cubic: Sodium chloride
- Rhombic: Iodine
- Tetragonal: Urea
- Hexagonal: Iodoform
- Monoclinic: Sucrose
- Triclinic: Boric acid
Amorphous Solids
- Randomly arranged molecules
- Nondistinct melting point
- Isotropic (properties are the same in all directions)
- Examples: Plastics, glass, pitch
Building Units of Crystals
- Atoms and molecules are held together by various forces of attraction, including:
- Van der Waals forces: Weak attractions between nonpolar molecules.
- Hydrogen bonds: Strong attractions between molecules containing hydrogen bonded to highly electronegative atoms like oxygen, nitrogen, or fluorine.
- Ionic bonds: Electrostatic attraction between positively and negatively charged ions.
- Covalent bonds: Sharing of electrons between atoms.
Solvates
- Crystals containing solvent molecules, sometimes called pseudopolymorphs.
Polymorphic Forms
- Different crystalline forms of the same substance with different physical properties.
- Enantiotropic: Change from one form to another is reversible.
- Monotropic: Transition only occurs in one direction, usually from metastable to stable form.
Amorphous Solids
- Molecules are not packed in a repeating long-range ordered fashion.
- No definite melting point.
- Tend to flow under sufficient pressure.
- Exhibit a yield value: The force below which the body shows elastic properties.
Examples of Drugs Exhibiting Polymorphism
- Theobroma oil
- Estrogens (form crystalline solvates)
- Tamoxifen
- Carbamazepine
Gibbs Phase Rule
- F = C - P + 2
- F: Number of degrees of freedom (temperature, pressure, concentration)
- C: Number of components (minimum number of constituents to express the composition of each phase)
- P: Number of phases present (homogeneous, physically distinct portions of a system separated by bounding surfaces)
Condensed Systems
- Systems where the vapor phase is ignored, considering only solid and/or liquid phases.
Water-Phenol System
- Partially miscible system.
- Two liquid phases exist at 11% phenol in water to 63% phenol in water at 50°C.
- One liquid phase exists when the concentration exceeds 63% at 50°C.
- All combinations of phenol-water above 66.8°C are completely miscible, yielding one-phase liquid.
- Critical solution temperature (Upper consolute temperature): The maximum temperature at which the two-phase region exists.
Eutexia
- Liquefaction resulting from combining two solids due to lowering of their melting points.
- Examples:
- Salol-Thymol: Eutectic point at 34% thymol in salol, 13°C.
- Salol-Camphor: Eutectic point at 56% by weight of salol in camphor, 6°C.
- Camphor-Thymol-Menthol
- Prilocaine-Lidocaine: 1:1 mixture at 18°C.
Thermodynamics
- Deals with the quantitative relationship between heat and other forms of energy.
Thermochemistry
- The study of heat changes in chemical reactions and physical processes.
- 1st Law of Thermodynamics: Derivation of enthalpy (heat content) as HS = ΔH, where ΔH is the heat absorbed when the solid melts.
Non-Ideal Solutions
- The heat of solution can be positive (energy absorbed) or negative (energy liberated).
- A negative heat of solution favors solubility, while a positive heat works against it.
Solubility of Strong Electrolytes
- Dissolution is generally accompanied by:
- Absorption of Heat (Endothermic): Solution becomes cold; an increase in temperature increases solubility (e.g., dissolution of Na2SO4.H2O).
- Evolution of Heat (Exothermic): Solution is hot; solubility decreases with an increase in temperature (e.g., dissolution of anhydrous Na2SO4).
- Sodium chloride does not evolve or absorb heat when dissolving in water, so its solubility is largely unaffected by temperature change (heat of solution is approximately zero).
Effect of Temperature on Solubility
- Decrease in solubility with temperature (exothermic).
- Increase in solubility with temperature (endothermic).
Solubility of Slightly Soluble Electrolytes
- Solubility Product (Ksp): Describes the solubility of sparingly soluble electrolytes to form a saturated solution. Ksp = [A+]a[B-]b.
Common Ion Effect
- Reduces the solubility of slightly soluble electrolytes.
- When the common ion forms a complex with the salt, the net solubility may increase.
- Example: Adding Ag+ or Cl- ions to AgCl will alter the equilibrium, reducing its solubility: AgCl + NaCl → AgCl(↓).
Commonly Used Concentration Expressions for Solutions
- Molarity (M): Moles of solute per liter of solution.
- Molality (m): Moles of solute per kilogram of solvent.
- Normality (N): Gram equivalent weight of solute per liter of solution.
- Weight/volume %: Weight of solute per 100mL of solution.
- Volume/volume %: Volume of solute per 100mL of solution.
- Parts Per Million (ppm): mg of solute per liter of solution.
- Parts Per Billion (ppb): µg of solute per liter of solution.
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Description
Explore the concepts of intermolecular forces in gases and the conditions under which gases transition to liquids. Learn about pseudoplastic flow and other types of flow, including thixotropy, and how they relate to viscosity changes. This quiz covers essential principles in fluid dynamics and physical chemistry.
