1) In an experiment to determine the molecular weight and the ionization constant for ascorbic acid (vitamin C), a student dissolved 1.3717 grams of the acid in water to make 50.00 milliliters of solution. The entire solution was titrated with a 0.2211-molar NaOH solution. The pH was monitored throughout the titration. The equivalence point was reached when 35.23 milliliters of the base had been added. Under the conditions of this experiment, ascorbic acid acts as a monoprotic acid that can be represented as HA.
(a) From the information above, calculate the molecular weight of ascorbic acid.
(b) When 20.00 milliliters of NaOH had been added during the titration, the pH of the solution was 4.23. Calculate the acid ionization constant for ascorbic acid.
(c) Calculate the equilibrium constant for the reaction of the ascorbate ion, A¯, with water.
(d) Calculate the pH of the solution at the equivalence point of the titration.
2) The electrolysis of an aqueous solution of potassium iodide, KI, results in the formation of hydrogen gas at the cathode and iodine at the anode. A sample of 80.00 milliliters of a 0.150-molar solution of KI was electrolyzed for 3.00 minutes, using a constant current. At the end of this time, the I2 produced was titrated against a 0.225-molar solution of sodium thiosulfate, which reacts with iodine according to the equation below. The end point of the titration was reached when 37.2 milliliters of the Na2S2O3 solution had been added.
I2 + 2 S2O32¯ ---> 2 I¯ + S4O62¯
(a) How many moles of I2 was produced during the electrolysis?
(b) The hydrogen gas produced at the cathode during the electrolysis was collected over water at 25 °C at a total pressure of 752 millimeters of mercury. Determine the volume of hydrogen collected. (The vapor pressure of water at 25 °C is 24 millimeters of mercury.)
(c) Write the equations for the half reaction that occurs at the anode during the electrolysis.
(d) Calculate the current used during the electrolysis.
3) Br2 (l) ---> Br2 (g)
At 25 °C the equilibrium constant, Kp, for the reaction above is 0.281 atmosphere.
(a) What is DG°298 for this reaction?
(b) It takes 193 joules to vaporize 1.00 gram of Br2(l) at 25 °C and 1.00 atmosphere pressure. What are the values of DH°298 and of DS°298 for this reaction?
(c) Calculate the normal boiling point of bromine. Assume that DH° and DS° remain constant as the temperature is changed.
(d) What is the equilibrium vapor pressure of bromine at 25 °C ?
4) Give the formulas to show the reactants and the products for FIVE of the following chemical reactions. Each of the reactions occurs in aqueous solution unless otherwise indicated. Represent substances in solutions as ions if the substance is extensively ionized. Omit formulas for any ions or molecules that are unchanged by the reaction. In all cases a reaction occurs. You need not balance. Example: A strip of magnesium is added to a solution of silver nitrate.
(a) Solutions of zinc sulfate and sodium phosphate are mixed.
(b) Solutions of silver nitrate and lithium bromide are mixed.
(c) A stream of chlorine gas is passed through a solution of cold, dilute sodium hydroxide.
(d) Excess hydrochloric acid solution is added to a solution of potassium sulfite.
(e) A solution of tin(II) chloride is added to an acidified solution of potassium permanganate.
(f) A solution of ammonium thiocyanate is added to a solution of iron(III) chloride.
(g) Samples of boron trichloride gas and ammonia gas are mixed.
(h) Carbon disulfide vapor is burned in excess oxygen.
| CF4 | XeF4 | ClF3 |
(a) Draw a Lewis electron-dot structure for each of the molecules above and identify the shape of each.
(b) Use the valence shell electron-pair repulsion (VSEPR) model to explain the geometry of each of these molecules.
6) The melting points of the alkali metals decrease from Li to Cs. In contrast, the melting of the halogens increase from F2 to I2.
(a) Using bonding principles, account for the decrease in the melting point of the alkali metals.
(b) Using bonding principles, account for the increase in the melting points of the halogens.
(c) What is the expected trend in the melting points of the compounds LiF, NaCl, KBr, and CsI? Explain this trend using bonding principles.
7) Consider three unlabeled bottles, each containing small pieces of one of the following metals.
- Magnesium
- Sodium
- Silver
The following reagents are used for identifying the metals.
- Pure water
- A solution of 1.0-molar HCl
- A solution of concentrated HNO3
(a) Which metal can be easily identified because it is much softer than the other two? Describe a chemical test that distinguishes this metal from the other two, using only one of the reagents above. Write a balanced chemical equation for the reaction that occurs.
(b) One of the other two metals reacts readily with the HCl solution. Identify the metal and write the balanced chemical equation for the reaction that occurs when this metal is added to the HCl solution. Use a table of standard reduction potentials to account for the fact that this metal reacts with HCl while the other does not.
(c) The one remaining metal reacts with the concentrated HNO3 solution. Write a balanced chemical equation for the reaction that occurs.
(d) The solution obtained in (c) is diluted and a few drops of 1 M HCl is added. Describe what would be observed. Write a balanced chemical equation for the reaction that occurs.
8) C2H4 (g) + H2 (g) ---> C2H6 (g)
For the above reaction, DH° = -137 kJ
Account for the following observations regarding the exothermic reaction represented by the equation above.
(a) An increase in the pressure of the reactants causes an increase rate.
(b) A small increase in temperature causes a large increase in the reaction rate.
(c) The presence of metallic nickel causes an increase in reaction rate.
(d) The presence of powdered nickel causes a larger increase in reaction rate than does the presence of a single pieces of nickel of the same mass.
9) The carbon isotope of mass 12 is stable. The carbon isotopes of mass 11 and mass 14 are unstable. However, the type of radioactive decay is different for the these two isotopes. Carbon-12 is not produced in either case.
(a) Identify a type of decay expected for carbon-11 and write the balanced nuclear reaction for that decay process.
(b) Identify the type of decay expected for carbon-14 and write the balanced nuclear reaction for that decay process.
(c) Gamma rays are observed during the radioactive decay of carbon-11. Why is it unnecessary to include the gamma rays in the radioactive decay equation of (a) ?
(d) Explain how the amount of carbon-14 in a piece of wood can be used to determine when the tree died.
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