# Example Final Conceptual Questions

Topical Questions: (these should be mastered so that you can answer them quickly without thinking)

## Thermodynamics

• What is the difference between a “microstate” and a “macrostate”?
• Explain the concept of entropy
• Explain why is an isothermal expansion of an ideal gas is spontaneous when the change in internal energy is zero
• Can the absolute zero of entropy be determined?
• How does increasing temperature affect the entropy of a system? Why?
• What has the higher entropy, liquid water or water vapor? Why?
• What has the higher entropy, $$NO (g)$$ or $$NO_2 (g)$$?
• What is the second law of thermodynamics?
• What characterizes a spontaneous reaction?
• Why do we introduce Gibbs free energy?
• What is ΔG at equilibrium?
• What drives a chemical reaction?
• Why are some reactions spontaneous only at high temperatures while others are spontaneous only at low temperatures?
• What is the definition of the free energy of formation?
• What does the sign of the standard free energy of a compound tell you?
• What is the standard state?
• How does ΔGº relate to ΔG at nonstandard conditions?
• Explain how the equilibrium constant and Gibbs energy are related
• What is the temperature dependence of K?
• For each of these processes state if they are endothermic or exothermic, cause an increase or a decrease in entropy, and whether the process occurs spontaneously. Not all of them has unique answers, describe what the answers will depend on
• Condensation of a gas to a liquid
• Dissolution of an ionic solid
• Mixing of water and oil
• Mixing of water and ethanol
• Melting of an ice cube
• $$H_2 (g) + O_2 (g) → H_2O (l)$$
• Isothermal expansion of an ideal gas

## Liquids, Solids, and Intermolecular Forces

• Define the following terms
• Cohesive forces (forces between like molecules)
• Adhesive forces (forces between unlike molecules)
• Surface tension
• Viscosity
• Classify IMFs into one of four categories
• Permanent electrostatic-Permanent electrostatic
• Permanent Electrostatic-Induced electrostatic
• Spontaneous Electrostatic-Induced Electrostatic (Dispersion)
• Repulsion
• Define Polarizability
• What IMF category(ies) does polarizability play a strong role in?
• What factors influence the magnitude of polarizability?
• What is the difference between London force and Dispersion force?
• Under what conditions will hydrogen bonding be observed?
• How do the IMFs strengths compare to covalent bonding strengths?
• What is the origin of the repulsion force?
• In what systems are repulsion forces not involved?
• Which IMF has the longest range of interaction (be specific)?
• Which IMF has the shortest range of interaction (be specific)?
• What is a hydrogen bond donor?
• What is a hydrogen bond acceptor?
• Give two examples of phases that do not correspond to differing states.
• What is the definition of vapor pressure?
• How is vapor pressure related to IMFs?
• What is the definition of boiling?
• What are two ways that you can make a liquid boil?
• Define
• critical point
• triple point
• melting
• melting point
• enthalpy of fusion
• heat of fusion
• enthalpy of vaporization
• heat of vaporization
• enthalpy of sublimation
• heat of sublimation
• Supercritical fluids
• Phase transitions
• How is ·Enthalpy of vaporization related to ·Enthalpy of sublimation?
• What is the origin of this relationship?
• What are the differences between the phase diagram for water and for $$\ce{CO2}$$?
• When would you use the Claussius-Clapeyron equation?
• When would you use the Clapeyron equation?
• What is the difference between the two? What is the approximation used to derive the Claussius-Clapeyron equation?

## Solutions

• What are the definitions of the following concentration units:
• Mass percent?
• ppm?
• Mole fraction?
• Molarity?
• Molality?
• Which concentration units are temperature dependent?
• Which concentration units are temperature independent?
• What determines if the enthalpy of solution is positive or negative?
• What is an ideal solution?
• What drives the formation of an ideal solution?
• Crystal lattice energies are typically very strong. Give two reasons why many salts nevertheless are easily soluble in water
• How does the solubility of gases depend of the partial pressure of the gas?
• What is the definition of a lattice energy?
• What is the definition of a Solvation energy?
• Definition between real and ideal solutions
• Definition of Activities
• How does the vapor pressure of an ideal solution depend on composition?
• Is the composition of the vapor phase and liquid phase for a solution different? Why?
• What is fractional distillation? How does it work?
• What is an azeotrope?
• What is osmotic pressure?
• What is reversed osmosis?
• What are colligative effects? Which four did we discuss in class?
• Using a phase diagram, explain how freezing point depression and boiling point elevation are consequences of Raoult’s law
• What is the van ’t Hoff factor?
• What is Henry's law?
• Is it applicable for ideal solutions?
• What is Raoult's law?
• Is it applicable for ideal solutions
• Is Raoult's law applicable for volatile or non-volatile solutions? Or both or neither?
• What is the definition of a saturated solution?
• What does supersaturated mean?
• Does a solid have a vapor pressure?
• Is vaporization an endothermic or an exothermic process?

## Equilibria

• What is a dynamic equilibrium? How does it compare to static equilibria?
• What is the law of mass action?
• What is the definition of $$K_c$$?
• What are the units of $$K_c$$?
• How does the equilibrium constant depend on the initial conditions?
• What is the activity?
• If a reaction can be written as the sum of two reactions with known equilibrium constants, how can you determine the equilibrium constant for the overall reaction?
• What is $$K_p$$?
• What is the difference between $$K_p$$ and $$K_c$$?
• How do pure liquids and pure solids enter the equilibrium constant? Why?
• What can you tell from the size of the equilibrium constant?
• What is the difference between K and Q?
• How can you predict the direction of a reaction knowing K and Q?
• If you add more reactant to a reaction that is initially at equilibrium, what will happen?
• If you increase the partial pressures equally of all gases involved in a reaction that is initially at equilibrium, what will happen?
• According to Le Chatelier’s principle, what happens to an exothermic reaction if the temperature is increased? Does K change?
• If Q >K reaction, does the reaction (as written) go to the left, right or not change?
• If Q <K reaction, does the reaction (as written) go to the left, right or not change?
• If Q =K reaction, does the reaction (as written) go to the left, right or not change?

## Acids and Bases

• Is an aqueous solution of NaCl acidic or basic? Why?
• Is an aqueous solution of NH4Cl acidic or basic? Why?
• What is the pH of a solution containing 0.1 M HCl and 0.1 M acetic acid? What is the concentration of acetic acid and acetate (the conjugate base) in this solution?
• What is a buffer solution?
• How do you make a buffer solution?
• What is the pH of a buffer solution?
• What is the buffer capacity? What does it depend on?
• What characterizes an acid-base indicator ?
• What is the equivalence point in a titration?
• What is pH halfway to the equivalence point during a titration of a weak base with a strong acid?
• How are pH and pOH related?
• How is the pKa of an acid related to the pKb of the conjugate base?
• What is pH for a 0.1M solution of HNO3?
• What is pH for a 10-8 M solution of HNO3?
• What defines a strong acid?
• Name three strong acids
• Which one is the stronger acid: Acetic acid (pKa= 4.74) or formic acid (pKa= 3.74)?
• Which one is the stronger base: the acetate ion ($$C_2H_3O_2^−$$) or formate ion ($$CHO_2^−$$)?
• When can you simplify the solution of an equilibrium calculation (by assuming that the change x is small)?

## Solubility

• What is the definition of the solubility product?
• How does the solubility product relate to the solubility defined in chapter 13?
• How do you determine if precipitation occurs in a solution?
• What is the common-ion effect?
• Explain why the solubility of a sparingly soluble salt increases when a highly soluble salt of spectator ions is added to the solution?
• What is “salting out”?
• Give one explanation for why “salting out” occurs
• What is the activity coefficient?
• Why is Ksp not constant?
• What is fractional precipitation?
• For what types of salts is the solubility affected by pH?