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Questions and Answers
Consider the following compounds: methane (CH4), dichloromethane (CH2Cl2), and carbon tetrachloride (CCl4). Which statement correctly describes the trend in their boiling points based on intermolecular forces?
Consider the following compounds: methane (CH4), dichloromethane (CH2Cl2), and carbon tetrachloride (CCl4). Which statement correctly describes the trend in their boiling points based on intermolecular forces?
- Dichloromethane has the highest boiling point due to significant dipole-dipole interactions and London dispersion forces. (correct)
- Methane has the highest boiling point due to its small size and strong London dispersion forces.
- The boiling points will be approximately the same since they all contain carbon.
- Carbon tetrachloride has the highest boiling point due to its tetrahedral symmetry, resulting in the greatest induced dipole interactions.
A chemist dissolves an ionic compound in water and observes a significant decrease in temperature. Which of the following best explains the thermodynamics of this process?
A chemist dissolves an ionic compound in water and observes a significant decrease in temperature. Which of the following best explains the thermodynamics of this process?
- The dissolution is endothermic, with a large positive enthalpy change that is overcome by a significant entropy increase, resulting in a negative Gibbs free energy. (correct)
- The dissolution is exothermic, with a positive enthalpy change and an increase in entropy.
- The dissolution is endothermic, with a positive enthalpy change that outweighs the entropy increase, leading to a negative Gibbs free energy.
- The dissolution is exothermic, with a negative enthalpy change and a decrease in entropy.
Consider a gas mixture in a closed container. If the temperature of the gas is increased while maintaining a constant volume, which of the following statements is most accurate according to the kinetic molecular theory of gases?
Consider a gas mixture in a closed container. If the temperature of the gas is increased while maintaining a constant volume, which of the following statements is most accurate according to the kinetic molecular theory of gases?
- The average kinetic energy of the gas molecules increases, and the frequency of collisions with the container walls decreases.
- The average kinetic energy of the gas molecules remains constant, and the pressure decreases.
- The average kinetic energy of the gas molecules decreases, and the pressure remains constant.
- The average kinetic energy of the gas molecules increases, and the pressure increases due to more frequent and forceful collisions. (correct)
Which of the following scenarios would result in the most significant deviation from ideal gas behavior?
Which of the following scenarios would result in the most significant deviation from ideal gas behavior?
For the reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, $\Delta H < 0$. Which of the following changes will simultaneously increase both the equilibrium constant (K) and the rate of the forward reaction?
For the reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, $\Delta H < 0$. Which of the following changes will simultaneously increase both the equilibrium constant (K) and the rate of the forward reaction?
Consider a solution of $AgCl$ in water at equilibrium. If $NaCl$ is added to this solution, what will happen to the solubility of $AgCl$ and why?
Consider a solution of $AgCl$ in water at equilibrium. If $NaCl$ is added to this solution, what will happen to the solubility of $AgCl$ and why?
Which of the following statements accurately describes the behavior of gases when they deviate significantly from ideal behavior?
Which of the following statements accurately describes the behavior of gases when they deviate significantly from ideal behavior?
In a mixture of ideal gases, how does the partial pressure of each gas relate to the total pressure, and what fundamental principle governs this relationship?
In a mixture of ideal gases, how does the partial pressure of each gas relate to the total pressure, and what fundamental principle governs this relationship?
Consider the dissolution of a solid in a liquid. Under what conditions would the entropy change ($\Delta S$) be most likely to be negative?
Consider the dissolution of a solid in a liquid. Under what conditions would the entropy change ($\Delta S$) be most likely to be negative?
For a certain reaction, $\Delta H$ is positive and $\Delta S$ is negative at standard conditions. Which statement about the spontaneity of the reaction is correct?
For a certain reaction, $\Delta H$ is positive and $\Delta S$ is negative at standard conditions. Which statement about the spontaneity of the reaction is correct?
Given the following reactions and their corresponding enthalpy changes:
$A \rightarrow B$, $\Delta H_1 = +50 kJ$
$B \rightarrow C$, $\Delta H_2 = -25 kJ$
$C \rightarrow D$, $\Delta H_3 = +10 kJ$
What is the enthalpy change for the reaction $A \rightarrow D$?
Given the following reactions and their corresponding enthalpy changes:
$A \rightarrow B$, $\Delta H_1 = +50 kJ$ $B \rightarrow C$, $\Delta H_2 = -25 kJ$ $C \rightarrow D$, $\Delta H_3 = +10 kJ$
What is the enthalpy change for the reaction $A \rightarrow D$?
Consider a reaction at equilibrium: $aA(g) + bB(g) \rightleftharpoons cC(g) + dD(g)$. If the volume of the container is suddenly decreased, which of the following statements is true regarding the shift in equilibrium?
Consider a reaction at equilibrium: $aA(g) + bB(g) \rightleftharpoons cC(g) + dD(g)$. If the volume of the container is suddenly decreased, which of the following statements is true regarding the shift in equilibrium?
Which of the following statements best describes the behavior of a real gas under conditions of high pressure and low temperature?
Which of the following statements best describes the behavior of a real gas under conditions of high pressure and low temperature?
Which of the following solubility rules contains an exception regarding Group 2 compounds?
Which of the following solubility rules contains an exception regarding Group 2 compounds?
Given the reaction: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$ $\Delta H < 0$. Which action will simultaneously increase the equilibrium constant K and decrease the rate of the reverse reaction?
Given the reaction: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$ $\Delta H < 0$. Which action will simultaneously increase the equilibrium constant K and decrease the rate of the reverse reaction?
For the equilibrium reaction $N_2O_4(g) \rightleftharpoons 2NO_2(g)$, it is observed that the forward reaction is endothermic. Which change will shift the equilibrium such that the ratio of $NO_2$ to $N_2O_4$ decreases?
For the equilibrium reaction $N_2O_4(g) \rightleftharpoons 2NO_2(g)$, it is observed that the forward reaction is endothermic. Which change will shift the equilibrium such that the ratio of $NO_2$ to $N_2O_4$ decreases?
If a soluble salt $AB_2$ is added to water, and the concentration of $A^{2+}$ ions at equilibrium is found to be 's' mol/L, what is the solubility product constant, $K_{sp}$, for the salt?
If a soluble salt $AB_2$ is added to water, and the concentration of $A^{2+}$ ions at equilibrium is found to be 's' mol/L, what is the solubility product constant, $K_{sp}$, for the salt?
Which of the following scenarios would result in the largest increase in entropy?
Which of the following scenarios would result in the largest increase in entropy?
How does the addition of a catalyst affect the thermodynamics and kinetics of a reversible reaction?
How does the addition of a catalyst affect the thermodynamics and kinetics of a reversible reaction?
Consider the reaction $A(g) + B(g) \rightleftharpoons C(g)$. Under what conditions of temperature and pressure would this gas-phase reaction be most likely to deviate from predictions based on the ideal gas law?
Consider the reaction $A(g) + B(g) \rightleftharpoons C(g)$. Under what conditions of temperature and pressure would this gas-phase reaction be most likely to deviate from predictions based on the ideal gas law?
Flashcards
Intermolecular Forces
Intermolecular Forces
Attractive or repulsive forces between molecules, influencing physical properties.
Types of Intermolecular Forces
Types of Intermolecular Forces
Ion/ion, ion/dipole, dipole/dipole (including hydrogen bonding), dipole-induced dipole, and London dispersion forces.
Hydrogen Bonding
Hydrogen Bonding
Strong dipole/dipole interaction with H bonded to F, O, or N.
London Dispersion Forces
London Dispersion Forces
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Heat of Vaporization
Heat of Vaporization
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Phase Diagram
Phase Diagram
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Vaporization
Vaporization
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Sublimation
Sublimation
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Deposition
Deposition
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Solvation Interactions
Solvation Interactions
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Entropy (S)
Entropy (S)
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Gibbs Free Energy (ΔG)
Gibbs Free Energy (ΔG)
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Spontaneous Reaction
Spontaneous Reaction
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Equilibrium (ΔG = 0)
Equilibrium (ΔG = 0)
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Non-Spontaneous Reaction
Non-Spontaneous Reaction
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Equilibrium Constant (K)
Equilibrium Constant (K)
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Product-Favored Reaction
Product-Favored Reaction
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Reactant-Favored Reaction
Reactant-Favored Reaction
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Value of K
Value of K
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Le Chatelier's Principle
Le Chatelier's Principle
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Study Notes
Intermolecular Forces
- Basic types include ion/ion, ion/dipole, dipole/dipole (including hydrogen bonding), dipole-induced dipole, and London dispersion forces
- Hydrogen bonds are strong dipole/dipole interactions, ~20 kJ/mole
- Non-hydrogen bond dipole/dipole interactions are ~8-10 kJ/mole
- Typical London dispersion forces between two atoms are ~0.3 kJ/mole
- Covalent bonds are ~300-400 kJ/mole
- All compounds experience London dispersion forces
- Molecular size, shape, and polarity affect intermolecular forces based on Coulomb's law
- Intermolecular forces affect physical properties like boiling point, viscosity, and water solubility
- Solubility requires intermolecular interactions with the dissolving liquid
- Hydrogen bonding in water leads to unique properties: high boiling point, solid phase less dense than liquid, cohesive and adhesive forces, and surface tension
Phases and Phase Changes
- Adding enough heat energy overcomes intermolecular forces, converting a liquid to a gas at its boiling point
- Gases are assumed to have negligible intermolecular interactions
- Gas condensation releases heat as molecules interact
- Liquid-to-solid or gas-to-solid transitions release heat (heat of fusion), dependent on intermolecular force strength
- Thermal energy disperses as vibrational, rotational, and translational motion in atoms, ions, or molecules and correlates to physical phases
- Heating curves estimate relative heat of vaporization/fusion and heat capacities of solid, liquid, and gas phases qualitatively
- Phase diagrams predict a material's phase at various temperatures and pressures
- Phase change terms: Vaporization (boiling) - liquid to gas, condensation - gas to liquid, melting - solid to liquid, fusion (freezing) - liquid to solid, sublimation - solid to gas, deposition – gas to solid
- Phase diagrams include triple points and critical points
Solubility and Solutions
- Dissolving ionic compounds in water involves breaking solute/solute interactions, solvent/solvent interactions (water H-bonds), and forming solute/solvent interactions (ion-dipole)
- ΔG, ΔH, and ΔS signs can be predicted for solution processes
- Solubility rules
- Group 1 cations (Na+, K+, etc.) and ammonium (NH4+) compounds are soluble
- Nitrate (NO3-), chlorate (ClO3-), perchlorate (ClO4-), and acetate (CH3CO2-) compounds are soluble
- Halide (Cl-, Br-, I-) compounds are soluble, EXCEPT those with silver (I) (Ag+), mercury (I) (Hg22+), copper (I) (Cu+), and lead (II) (Pb2+); Fluoride (F-) compounds are EXCEPT group 2 compounds, which are insoluble
- Sulfate (SO42-) compounds are soluble, EXCEPT those with Group 2 cations (Ca2+, Sr2+, Ba2+) and Pb2+
- Most other compounds not explicitly listed are considered insoluble
- Solutions with concentrations greater than 0.05 M are considered soluble
- Predict if an ionic compound is soluble and if a precipitate will form when mixing two ionic solutions, based on solubility rules
- Spectator ions in reactions, used to write net ionic equations for precipitation reactions
- Calculate concentration (mol/l or molarity) and number of moles of a substance in a given volume of solution: (mol/l) x L = mol
- Write chemical equations for ionic substances dissolving in water and the expression for the Ksp
- The general format is MaXb(s) + H2O(l) ⇌ a M+(aq) + bX-(aq)
- Ksp (K) values indicate solubility to make calculations
Solution Dilution
- Solutions can be diluted by adding more solvent
- Use M1 x V1 = M2 x V2 to calculate the new molarity or volume required for dilution (M = molarity, V = volume in same units)
Gases and Kinetic Molecular Theory
- Ideal gas law: PV = nRT
- Use P1V1/n1T1 = R = P2V2/n2T2 for gases changing conditions
- Basic assumptions of kinetic molecular theory
- Individual gas particle volume is negligible compared to container volume
- No interactions exist between particles
- All collisions are elastic
- Gas particles are in constant random motion
- Temperature relates to average kinetic energy (½ mvave2)
- Gas molecules possess translational energy, but has some rotational and vibrational
- Total pressure of a gas mixture is the sum of individual partial pressures, proportional to the number of moles of each gas (Dalton’s Law)
- Real gases deviate from ideal behavior at high pressures and low temperatures due to intermolecular forces
Thermochemistry, Thermodynamics and Equilibrium
- Thermal energy disperses as translational, rotational, and vibrational motion in atoms, ions, or molecules, correlated to physical phases
- ΔH and ΔE mean and how they are related to each other
- ΔH is defined as ΔE + PΔV at constant T and P
- Estimate ΔH (ΔHrxn) for a reaction using differences between the bond enthalpies or energies of the products and reactants
- Hess’s Law involves combining different ΔH values for a series of reactions to calculate the ΔH of a reaction (ΔHrxn)
- Use ΔH°f values to calculate the ΔH° of a reaction (ΔH°rxn)
- Calculate ΔH for a reaction from experimental calorimetry data, using specific or molar capacity, amount, and temperature rise
Characterizing Chemical Reactions: Energy, Entropy, ΔH, ΔG, Equilibrium and K
- Determine direction, forward or reverse, of a chemical reaction is favored
- Entropy is not conserved and measures energy dispersion; energy moves from less to more dispersed conditions
- In the general statement ΔSuniv = ΔSsurd + ΔS, the value of ΔSuniv > 0 for spontaneous processes
- ΔSsurd and ΔS must be considered separately, ΔSsurd relates directly with the ΔH of the reaction
- ΔSsurd for a reaction is given as – ΔH/T
- ΔS (system) depends on the system's entropy change affected by the phases of products vs reactants and bonds in products vs reactants
- In that the substitution of – ΔH/T for ΔSsurd is made in the ΔSuniv = ΔSsurd + ΔS equation and each side of the equation is multiplied by –T, –TΔSuniv = ΔH – TΔS with ΔG defined as –TΔSuniv
- A reaction is endothermic or exothermic shown by signs
- Sign of ΔG determines if reaction is product- or reactant-preferred
Sign of ΔG
- Negative ΔG means the reaction is spontaneous or exergonic, therefore preferred in the forward direction
- ΔG = 0 means the reaction is at equilibrium, without net change
- Positive ΔG means the reaction is non-spontaneous or endergonic, therefore preferred in the reverse direction
Characterizing Chemical Reactions at Equilibrium
- Reactions proceed to equilibrium, reaching a ΔG value of 0
- At equilibrium, there is no net change in reactant or product amounts
- An equilibrium constant, Keq denoted as ([C]^c [D]^d)/ ([A]^a [B]^b) using concentrations (M)
- For reactions with reactants and products that are solids or liquids, then the [ ] term in the K expression has a value of 1
Using K Values
- Large K means product to reactant is large at equilibrium, which means the reaction is product-favored as the product/reactant ratio at equilibrium is > 1
- Small K means the product to reactant ratio is small at equilibrium, which means the reaction is reactant-favored as the product/reactant ratio at equilibrium is 1
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