Question: Energy is required to break chemical bonds

> 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

> Write a valid equilibrium constant expression for the reaction shown in Question 7.64. Question 7.64: H,S(aq) + Cl2 (aq) =S(s) + 2HCI(aq)

> 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.

> 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?

> What is the major importance of Raoult’s law?

> Describe the transport of lipids digested in the lumen of the intestines to the cells of the body.

> A 300.0-mL portion of H2O is added to 300.0 mL of 0.250 M H2SO4. What is the new molarity?

> After the soda in Question 6.5 is opened, the “fizz” shows a loss of CO2. If the partial pressure of CO2 in the atmosphere is 5.0 × 10-4 atm, calculate the equilibrium CO2 concentration in mol/L in the open bottle of soda. Question 6.5: An unopened bot

> Calculate the volume of a 1.00 × 10-2 M KOH solution containing 3.00 × 10-1 mol of solute.

> Will the volume of gas increase, decrease, or remain the same if the temperature is decreased and the pressure is increased? Explain.

> Write the dilution expression and define each term.

> Calculate the molarity of a solution that contains 1.75 mol of KNO3 dissolved in 3.00 L.

> Calculate the molarity of 2.75 L of solution containing 1.35 × 10-2 mol HCl.

> How many g of solute are needed to prepare each of the following solutions? a. 2.50 × 102 g of 5.00% (m/m) NH4Cl (ammonium chloride) b. 2.50 × 102 g of 3.50% (m/m) Na2CO3

> How many mL of 4.0 mass/volume % Mg (NO3)2 solution would contain 1.2 g of magnesium nitrate?

> Would the soft drink in Question 6.3 go “flat” faster if the bottle warmed to room temperature? Why? Question 6.3: Explain why, over time, a bottle of soft drink goes “flat” after it is opened.

> Use the kinetic molecular theory to explain why aerosol cans carry instructions warning against heating or disposing of the container in a fire.

> Describe the evidence that suggests that mitochondria evolved from free-living bacteria.

> A solution was prepared by dissolving 12.4 g of NaNO3 in sufficient water to produce 95.0 mL of solution. What is the mass/volume % of this solution?

> Calculate the composition of each of the following solutions in mass/mass %: a. 1.00 g KCl in 1.00 × 102 g solution b. 50.0 g KCl in 5.00 × 102 mL solution (d = 1.14 g/mL)

> Calculate the composition of each of the following solutions in mass/volume %: a. 20.0 g benzene dissolved in 1.00 × 102 mL solution b. 20.0 g acetic acid dissolved in 2.50 L solution

> Calculate the composition of each of the following solutions in mass/volume %: a. 0.700 g KCl in 1.00 mL solution b. 95.2 g MgCl2 in 0.250 L solution

> Fish kills (the sudden death of thousands of fish) often occur during periods of prolonged elevated temperatures. Pollution is often, but not always, the cause. Suggest another reason, based on solubility trends.

> Is CH3OH more likely to form a solution in water or benzene (C6H6)? Explain your reasoning.

> An isotope of technetium, mixed with sulfur and colloidally dispersed in water, is frequently used in diagnosing various medical conditions because it is readily taken up by various tissues prior to excretion. Explain why this important mixture is not a

> Compare and contrast the gas, liquid, and solid states with regard to the nature of the interactions among the particles.

> What role does electronegativity play in determining the bonding between atoms in a compound?

> Describe how you would distinguish experimentally between a true solution and a colloidal dispersion.

> In what form are the vitamins riboflavin, thiamine, niacin, and pantothenic acid needed by the pyruvate dehydrogenase complex?