Question: Write a valid equilibrium constant expression for

> Write an equation to represent the neutralization of an aqueous solution of HCl with an aqueous solution of KOH.

> What are the products of a neutralization reaction?

> Calculate formula mass and the molar mass of each of the following formula units: a. S8 b. (NH4)2SO4 c. CO2

> Two nitrogen atoms in a nitrogen molecule are held together more strongly than the two chlorine atoms in a chlorine molecule. Explain this fact by comparing their respective Lewis structures.

> The hydroxide ion concentration in a sample of urine was determined to be of 1.0 × 10-8 M (at 250C). Calculate the hydronium ion concentration of this aqueous solution.

> In a neutralization reaction, how many mol of NaOH are needed to react with 3 mol of HCl?

> Calculate the pH of a solution that has [OH-] = 6.7 × 10-9 M.

> Calculate the pH of a solution that has [H3O+] = 6.6 × 10-5 M.

> Distinguish between the rate constant and the equilibrium constant for a reaction.

> Can a dilute solution of a strong acid ever have a higher pH than a more concentrated solution of a weak acid? Why or why not?

> Write the rate law for the reaction: H2S(aq) + Cl2 (aq)↽−−−−⇀S(s) + 2HCl(aq) Represent the order as n, n', and so forth.

> A 6.00-mL portion of an 8.00 M stock solution is to be diluted to 0.400 M. What will be the final volume after dilution?

> A 50.0-mL sample of a 0.250 M sucrose solution was diluted to 5.00 × 102 mL. What is the molar concentration of the resulting solution?

> Identify the conjugate acid-base pairs for the reversible reactions in Question 8.4. Question 8.4: Write an equation for the reversible reactions of each of the following with water. a. H3PO4 b. CH3NH2

> Explain the differences between leading strand and lagging strand replication.

> Calculate [H3O+] for a solution of hydrochloric acid for which: a. pH = 2.00 b. pH = 3.00

> How many g of sodium hydroxide are present in 675 mL of a 0.500 M solution?

> Select one enzyme from a later chapter in this book and describe its biochemical importance.

> What is the concentration of hydronium ions in an aqueous solution of acetaminophen if the concentration of hydroxide ions is 2.5 × 10-9 M?

> Calculate the [OH-] of an aqueous solution that is: a. 1.0 × 10-6 M in H3O+ b. 1.0 × 10-8 M in H3O+

> Calculate the [H3O+] of an aqueous solution that is: a. 1.0 × 10-9 M in OH- b. 1.0 × 10-5 M in OH-

> Label each of the following as a strong or weak base: a. KOH b. CN- c. SO42-

> Of the following diagrams, which one represents: a. a concentrated strong acid b. a dilute strong acid c. a concentrated weak acid d. a dilute weak acid H+ x- H+ x- X- H+ x- H+ H+ x- H+ x- X- H+ x- H+ H+x-H+x=x¯H+ H+ H-X X- Н-X H-X H-X Н-X H-X H-X X

> Identify the conjugate acid-base pairs in each of the following chemical equations: a. HCOOH (aq) + NH3(aq) ↽−−−−−−⇀ HCOO-(aq) + NH4+(aq) b. HCl(aq) + OH-(aq) ↽−−−−−−⇀ H2O(l) + Cl-(aq)

> Which is the stronger base, F2 or CH3COO-?

> Distinguish between the term’s formula mass and molar mass.

> Write an equation for the reversible reactions of each of the following with water. a. H3PO4 b. CH3NH2

> Which is the stronger acid, HNO3 or HCN?

> Write the formula of the conjugate acid of F-.

> Write the formula of the conjugate base of HCOOH.

> Write the formula of the conjugate acid of Br-.

> Write an equation for the reaction of each of the following with water: a. HNO3 b. HCOOH c. CH3CH2CH2NH2

> Classify each of the following as either a Brønsted-Lowry acid, base, or as amphiprotic. a. H2SO4 b. HSO4- c. SO42-

> Classify each of the following as either a Brønsted-Lowry acid, base, or as amphiprotic. a. NH4+ b. NH3 c. CH3CH2CH2NH3+

> Why is ammonia described as a Brønsted-Lowry base and not an Arrhenius base?

> a. Define a base according to the Arrhenius theory. b. Define a base according to the Brønsted-Lowry theory.

> Describe the reaction catalyzed by each of the enzymes listed in Question 22.37. Question 22.37: Match each of the following enzymes with the class of enzyme to which it belongs. (Hint: An enzyme classification may be used more than once or not at all.

> State the second law of thermodynamics.

> Classify each of the following compounds as a Brønsted-Lowry acid or base. a. PO43- b. CH3NH3- c. HI d. H3PO4

> At a given temperature, the equilibrium constant for a certain reaction is 1 × 10-18. Does this equilibrium favor products or reactants? Why?

> Does the attainment of equilibrium imply that no further change is taking place in the system?

> Use the Henderson-Hasselbalch equation to calculate the pH of a buffer solution in which the acetic acid concentration is 2.0 × 10-1 M and the sodium acetate concentration is 1.0 × 10-1 M. The equilibrium constant, Ka, for acetic acid is 1.8 × 10-5.

> For the buffer system described in Question 8.99, which substance is responsible for buffering capacity against added sodium hydroxide? Explain. Question 8.99: For the equilibrium situation involving acetic acid, CH3COOH (aq) + H2O(l) −↽−−−−−−⇀− CH3COO

> True or false: The position of the equilibrium for an endothermic reaction will shift to the right when the reaction mixture is heated. Explain your reasoning.

> Calculate the [OH-] of a solution of hydrochloric acid with pH = 4.00.

> Sketch the interaction of water with an ethanol, CH3CH2OH, molecule.

> Sketch the “interactive network” of water molecules in the liquid state.

> Predict whether each of the following processes increases or decreases entropy, and explain your reasoning. a. burning a log in a fireplace b. condensing of water vapor on a cold surface

> Using the equilibrium constant expression in Question 7.85, calculate the equilibrium constant if: Question 7.85: Write the equilibrium constant expression for the reaction: N2 (g) + 3H2 (g)↽−−&a

> A change in pressure could have the greatest effect on which type of equilibria: gaseous, liquid, or solid?

> Can a catalyst alter the position of the equilibrium?

> Using the conversion factor in Chapter 1, convert the energy absorbed in Example 7.4 to joules (J). Example 7.4: 7.0 × 102 cal (or 0.70 kcal) of heat energy were absorbed by the dissolution process because the solution lost 7.0 × 102 cal of heat energy

> What is the relationship between the forward and reverse rates for a reaction at equilibrium?

> Calculate the pressure (atm) exerted by 1.00 mol of gas contained in a 7.55-L cylinder at 450C.

> Calculate the density of carbon dioxide at STP.

> Does a large equilibrium constant mean that the reaction must be rapid?

> Explain, in terms of solution properties, why a wilted plant regains its “health” when watered.

> Explain the mechanism by which a fatty acyl group is brought into the mitochondrial matrix

> Calculate the molar volume of O2 gas at STP.

> Predict whether a reaction with positive H and positive S will be spontaneous, nonspontaneous, or temperature dependent. Explain your reasoning.

> Explain how a catalyst can be involved in a chemical reaction without being consumed in the process.

> Distinguish between the terms kinetics and thermodynamics.

> Define and explain the term activation energy as it applies to chemical reactions.

> Provide an example of a reaction that is extremely fast, perhaps quicker than the eye can perceive.

> Which solution is more concentrated: a 20 ppt solution or a 200-ppm solution?

> A solution contains 1.0 mg of Cu2+ per 0.50 kg solution. Calculate the concentration in ppm.

> Explain why the fuel value of foods is an important factor in nutrition science.

> What energy unit is commonly employed in nutrition science?

> Use electronegativity values to classify the bonds in each of the following compounds as ionic, polar covalent, or nonpolar covalent. a. CaCl2 b. CO c. ICl d. H2

> Which substance has the greatest entropy, H2O(l) or H2O(g)? Explain your reasoning.

> Construct a diagram of a coffee-cup calorimeter.

> Explain what is meant by the term specific heat.

> Predict whether a reaction with a negative H and a negative S will be spontaneous, nonspontaneous, or temperature dependent. Explain your reasoning.

> Energy is required to break chemical bonds during the course of a reaction. When is energy released?

> The Henry’s law constant, k, for N2 in aqueous solution is 6.1 × 10-4 mol/ (L · atm) at 250C. When a diver is at a depth of 240 m, the pressure is approximately 25 atm. Calculate the equilibrium solubility of N2 at this depth (250C) in units of mol/L.

> Explain what is meant by the term entropy.

> Write the expression for free energy. When G is a negative value, what does it indicate about the spontaneity of the reaction?

> Provide an explanation for the fact that most decomposition reactions are endothermic but most combination reactions are exothermic.

> Describe what is meant by an endothermic reaction.

> What type of bond is the peptide bond? Explain why the peptide bond is rigid.

> Are the following processes exothermic or endothermic? a. S(s) + O2(g) −−−−→ SO2(g), H = -71 kcal b. N2(g) + 2O2(g) + 16.2 kcal −−−−→ 2NO2(g)

> A certain change in reaction conditions for an equilibrium process was found to increase the rate of the forward reaction much more than that of the reverse reaction. Did the amount of product increase, decrease, or remain the same? Explain your reasonin

> Bacterial growth decreases markedly in a refrigerator. Why?

> Describe conditions that can lead to dangerously low concentrations of potassium in the blood.

> Suggest a change in experimental conditions that would increase the yield of H2 in the reaction in Question 7.104. Question 7.104: C(s) + H2O(g)↽−−−−⇀CO(g) + H2 (g)

> a. Write the equilibrium constant expression for the reaction: C(s) + H2O(g)↽−−−−⇀CO(g) + H2 (g) b. Calculate the equilibrium constant if [H2O] = 0.40 M [CO] = 0.40 M [H2] = 0.20 M

> Carbonated beverages quickly go flat (lose CO2) when heated. Explain, using LeChatelier’s principle.

> Distinguish between the terms net energy and activation energy.

> How can the octet rule be used to determine the number of electrons gained or lost by an atom as it becomes an ion?

> Name the two most important cations in biological fluids.

> Explain how the pH of blood would change under each of the conditions described in Questions 8.15 and 8.16. Questions 8.15: Explain how the molar concentration of H2CO3 in the blood would change if the partial pressure of CO2 in the lungs were to incre

> label solution as isotonic, hypotonic, or hypertonic in comparison to 0.9% (m/V) NaCl (0.15 M NaCl). 3% (m/V) NaCl

> label solution as isotonic, hypotonic, or hypertonic in comparison to 0.9% (m/V) NaCl (0.15 M NaCl). 0.35 M glucose

> A gas mixture has three components, N2, F2, and He. The partial pressure of N2 is 0.35 atm and F2 is 0.45 atm. If the total pressure is 1.20 atm, what is the partial pressure of helium?

> Determine the osmolarity of 2.5 × 10-4 M C6H12O6 (nonelectrolyte).

> Two solutions, A and B, are separated by a semipermeable membrane. For each case, predict whether there will be a net flow of water in one direction and, if so, which direction. A is 0.10 M NaCl and B is 0.20 M glucose.

> Two solutions, A and B, are separated by a semipermeable membrane. For each case, predict whether there will be a net flow of water in one direction and, if so, which direction. A is 0.10 M glucose and B is 0.10 M KCl.

> Calculate the number of g of silver nitrate required to prepare 2.00 L of 0.500 M AgNO3.

> A sample of O2 gas occupies 257 mL at 208C and 1.20 atm. What is the volume of this O2 gas sample at STP?

> Using salt to try to melt ice on a day when the temperature is -200C will be unsuccessful. Why?