[LEGGETT--AP CHEMISTRY – MINIMAL FINAL REVIEW] / 2015-2016 /

CHAPTER 1.STOICHIOMETRY

  1. Convert 3.01 x 1023 atoms of rubidium to moles (0.500 mol Rb)
  2. The bond energy for HF is 568 KJ/mol. How much energy in Joules is required to break a single HF bond? (9.44 x 10-19 J)
  3. How many moles are present in 50.0 g of calcium chlorate? (0.242 mol)
  4. How many formula units are present in 0.272 g of nickel(II) nitrate? (8.96 x 1020)
  5. The density of liquid water is 0.997 g/mL at 25°C. How many moles of water are in 250.0 mL of water? (13.8 mol H2O)
  6. Determine the molarity of a solution prepared by dissolving 141.6 g of citric acid, C3H5O(COOH)3 , in water and then diluting the resulting solution to 3500.0 mL. (0.2106 M)
  7. What mass of glucose, C6H12O6, would be required to prepare 5.000 x 103 L of a 0.215 M solution? (1.94 x 105 g)
  8. Determine the mole fraction of a solution of 560 g of acetone, CH3COCH3, in 620 g of water. (0.219)
  9. What volume of water would be added to 16.5 mL of a 0.0813 M solution of sodium borate in order to get a 0.0200 M solution? (50.6 mL H2O)
  10. A chemist wants to prepare a stock solution of H2SO4 so that samples of 20.00 mL will produce a solution with a concentration of 0.50 M when added to 100.0 mL of water.
  11. What should the molarity of the stock solution be? (3.0 M)
  12. If the chemist wants to prepare 5.00 L of the stock solution from concentrated H2SO4 , which is 18.0 M, what volume of concentrated acid should be used? (0.83 L)
  13. The density of 18.0 M H2SO4 is 1.84 g/mL. What mass of concentrated H2SO4 should be used to make the stock solution in (b)? (1.5 x 103 g)
  14. Maleic acid, which is used to manufacture artificial resin, has the empirical formula CHO. Its molar mass is 116.1 g/mol. What is its molecular formula? (C4H4O4)
  15. Hydrated salts are very common. If you heat 2.105 g of CoCl2xH2O, and find that 1.149 g of CoCl2 remains, what is the value of x? (6)
  16. Ammonia gas can be prepared by the following reaction: CaO(s)+2NH4Cl(s) 2NH3(g)+ H2O(g)+CaCl2(s). If you mix 112 g of CaO and 224 g of NH4Cl, what is the theoretical yield of NH3? (68.1 g). What mass of excess is remaining?(10.3 g)
  17. Disulfur dichloride can be made by allowing chlorine gas to react with molten sulfur: S8(l) + 4Cl2(g)  4S2Cl2(g) If you begin with 12.0 g of S8 and 13.03 g Cl2 and you isolate only 15.2 g of S2Cl2, what is the percentage yield of S2Cl2? (60.1%)
  18. Styrene, the building block of polystyrene, is a hydrocarbon, a compound consisting only of C and H. If you burn 0.438 g of the compound, and find that it produces 1.481 g of CO2 and 0.303 g of H2O, determine the empirical formula of the compound. (CH)
  19. Aluminum bromide is a valuable laboratory chemical. If you use 25.0 mL of liquid bromine (D = 3.1023 g/mL) and excess aluminum metal, what is the maximum theoretical yield of Al2Br6? (86.3 g)
  20. In the photographic process silver bromide is dissolved by adding sodium thiosulfate, as shown in the balanced equation below. If you want to dissolve 0.250 g of AgBr(molar mass = 187.8 g/mol), how many mL of 0.0138 M Na2S2O3 should you add? (193 mL)

AgBr(s)+2Na2S2O3(aq) Na3Ag(S2O3)2(aq)+NaBr(aq)

  1. A soft drink contains an unknown amount of citric acid, C6H8O7. If 100. mL of the soft drink required 33.51 mL of 0.0102 M NaOH to neutralize completely the citric acid, how many grams of citric acid (MM = 192.13 g/mol) does the soft drink contain per 100 mL? The reaction of citric acid with NaOH is(all aq) shown below. (0.0219 g)

C6H8O7 + 3NaOH  Na3C6H5O7 + 3H2O(l)

  1. NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of NutraSweet and find the molecular formula. (The molar mass of NutraSweet is 294.30 g/mol). (C14H18N2O5)
  2. Sodium thiosulfate, Na2S2O3, is used as a “fixer” in black and white photography. Assume you have a bottle of sodium thiosulfate and want to determine its purity. You can titrate the thiosulfate ion with I2 according to the equation below. If you use 40.21 mL of 0.246 M I2 in the titration, what is the weight percent of Na2S2O3 (MM = 158.12 g/mol) in a 3.232-g sample of impure Na2S2O3? (96.7 %) 2S2O32-+ I2 S4O62-+ 2I-
  3. Menthol, the substance we can smell in mentholated cough drops, is composed of C, H, and O. A 0.1005 g sample of menthol is combusted, producing 0.2829 g of CO2and 0.1159 g of H2O. What is the empirical formula for menthol? (C10H20O)
  4. You are given a sample of a copper-containing alloy and asked to determine the mass percent of copper. After dissolving the metal in acid, you add an excess of KI, and the Cu2+ and I- ions undergo the reaction:

2Cu2+ + 5I─ 2CuI + I3─

The liberated I3- is titrated with sodium thiosulfate according to the equation:

I3─ + 2S2O32─  S4O62- + 3I─

If 26.32 mL of 0.101 M Na2S2O3 is required for titration to the equivalence point, what is the mass percent of Cu in 0.251 g of the alloy? (67.2%)

CHAPTER 2.ATOM

  1. Conceptual Structure
  2. Explain all of the variables in Coulomb’s law as it applies to the atom.
  3. Differentiate among energy levels, sublevels, and orbitals. Include what each communicates about region is space where an electron is most likely to be found.
  4. Define AufBau principle, Hund’s Rule, and Pauli Exclusion principle. Explain each based upon maximizing attractive forces and minimizing repulsive forces in an atom.
  5. Write the complete electron configurations for the following atoms.
  1. Ruthenium
  2. Radium
  3. Write the complete electron configurations for the following ions.
  1. Bismuth +3 ion
  2. Iron +2 ion
  3. Tellurium -2 ion
  4. Write the noble gas configurations for the following atoms.
  1. Nickel
  2. Tungsten
  3. Identify the element from the following photoelectron spectra. Indicate the sublevel and number of electrons represented in each peak of the spectra. Justify the difference in energy values for the same sublevels in each based upon atomic structure.
  1. PES can be used to identify elements in a mixture. What elements are present in the mixture below?

Potassium / Silicon / Chlorine
Mixture
  1. Conceptual Periodic Trends

INTRO: You need to IDENTIFY and JUSTIFY the trends. You cannot use one trend (like electronegativity) to justify another trend (such as ionization energy). You answer must include a comparison of attractive forces (proton-electron) and repulsive forces (electron-electron). Distance from the nucleus, energy levels, conversion of energy levels, and shielding are critical components of a justification. STUDY MY SUMMARY CHART CAREFULLY!

  1. Which of the following groups of elements is arranged correctly in order of increasing first ionization energy?
  2. Mg < C < N < F
  3. N < Mg < C < F
  4. Mg < N < C < F
  5. F < C < Mg < N
  6. Which of the following groups of elements is arranged correctly in order of decreasing atomic radius?
  7. Mg < S < Al < Cl
  8. Al < Mg < S < Cl
  9. Mg < Al < S < Cl
  10. Cl < S < Mg < Al
  11. Which of the following elements would have the greatest difference between the first and the second ionization energy?
  12. Li B. CC. FD. N
  13. Which of the following groups of isoelectronic species show the elements arranged correctly in order of increasing size?
  14. Na+ < O2- < F-
  15. F- < Na+ < O2-
  16. Na+ < F- < O2-
  17. F- < O2- < Na+
  18. An element having which of the following electronic configurations would have the greatest ionization energy?
  19. [He]2s22p3
  20. [He] 2s22p5
  21. [Ne]3s23p3
  22. [Ne]3s23p5
  23. Periodic trends:
  24. Which should be larger, the oxide ion, O2-, or the oxygen atom?
  25. Which should have the largest difference between the 1st and 2nd ionization energy? O, S, or Se
  26. Which of the following concerning second IE's is true?
  27. That of Al is higher than that of Mg because Mg wants to lose the second electron, so it is easier to take the second electron away.
  28. That of Al is higher than that of Mg because the electrons are taken from the same energy level, but the Al atom has one more proton.
  29. That of Al is lower than that of Mg because Mg wants to lose the second electron, thus the energy change is greater.
  30. That of Al is lower than that of Mg because the second electron taken from Al is in a p orbital, thus it is easier to take.
  31. The second ionization energies are equal for Al and Mg.

Ionization Energies for element X (kJ mol¯1)
First / Second / Third / Fourth / Five
1086 / 2352 / 4619 / 6221 / 37820
  1. The ionization energies for element X are listed in the table above. On the basis of the data, element X is most likely to be
  2. Li B. BeC. BD. CE. P
  3. Suppose that a stable element with atomic number 119, symbol Q, has been discovered.
  4. Write the ground-state electron configuration for Q, showing only the valence-shell electrons.
  5. Would Q be a metal or a nonmetal?
  6. On the basis of periodic trends, would Q have the largest atomic radius in its group or would it have the smallest? Explain in terms of electronic structure.
  7. What would be the most likely charge of the Q ion in stable ionic compounds?
  8. The correct ranking of alkali metals from most reactive to least reactive is:
  9. Be-Mg-Co-Sr-BaD. I-Br-Cl-F
  10. Cs-Rb-K-Na-LiE. Li-Na-K-Rb-Cs
  11. F-Cl-Br-I12
  1. Mathematical
  1. Calculate the average atomic mass of chromium given the following isotopes with abundance.

Chromium-50 / Chromium-52 / Chromium-53 / Chromium-54
Actual mass / 49.945046 / 51.940519 / 52.940651 / 53.938882
Abundance / 4.35 / 83.79 / 9.50 / 2.36
  1. The mass spectrum below shows the relative % abundance for the isotopes of zinc.
  2. Redraw the spectrum so that the true % abundance is on the y-axis. WATCH Y-AXIS CAREFULLY ! I AM NOT SURE WHETHER AP WILL USE RELATIVE % OR ACTUAL %.
  1. Calculate the average atomic mass of zinc.
  1. Perform the following calculations regarding electromagnetic radiation.
  2. Convert 4.398 x 10-19 J to frequency.
  3. Convert 893 nm to energy
  4. The energy of a photon is 8.22 x 10-18 J. What is the energy in KJ/mol?

CHAPTER 3.BONDING

  1. Explain all of the variables in Coulomb’s law as it applies to bonding.
  2. Ionic
  3. Covalent
  4. Metallic
  5. Differentiate among non-polar covalent, polar covalent, and ionic BONDS.
  6. What is/are the difference(s) between network covalent and molecular covalent bonding/compounds?
  7. Define each of the terms in the acronym VSEPR. How is the model used to predict the structure of molecular covalent substances?
  8. Predict whether the following pairs of atoms are more likely to form an ionic, metallic or covalent bond using only a periodic table. Justify your answer in each case.
  9. Al & Ga
  10. Zn & P
  11. P & Br
  12. Rank the polarity of the following bonds using only a periodic table. Justify your answer. P — N, P—S, P—Br, P—O
  13. Define formal charge. Calculate the formal charge of each element in the molecule below. How is formal charge linked to the selection of a possible structure of a molecule?
  1. Complete (and memorize!) the following summary chart for molecular covalent substances

Bonding atoms (or group of atoms) on a defined central atom / Non-bonded pairs of electrons on a defined central atom / “ABX” / Molecular structure / Bond angle(s) / Hybridization
2 / 0
3 / 0
2 / 1
4 / 0
3 / 1
2 / 2
5 / 0 / N/A
4 / 1 / N/A
3 / 2 / N/A
2 / 3 / N/A
6 / 0 / N/A
5 / 1 / N/A
4 / 2 / N/A
  1. Explain why the bond angle decreases as the number of non-bonded pairs on the central atom increases.
  2. Electromagnetic radiation is commonly used to study atoms and molecules. What type of electromagnetic radiation is used to study the following?
  3. Electronic transitions
  4. Ionization
  5. Molecular Vibrations
  6. Electronic transitions are studied using the UV/Vis range of EMR. Which type of motions, vibrational or electronic, requires the greatest energy? Justify your answer.
  7. The structural formula for the amino acid, glycine is shown. Indicate the structure, hybridization, and bond angles for all elements numbered. Justify your answer in each case.
  8. What is Beer’s law and how is it commonly used to study molecules?
  9. Molecules that have alternating (conjugated) double bonds absorb visible light (you should know this fact for test). What ions are colored and form solutions that absorb visible light? Justify your answer.
  10. Use the data below (excel or calculator) to make a Beer’s Law plot, write down the equation of the line, and calculate the molarity of the unknown solution. (9.01e-5M)

Molarity / Absorbance
7.817e-5 / 0.1
1.782e-4 / 0.188
2.7817e-4 / 0.272
3.7817e-4 / 0.43
4.7817e-4 / 0.469
5.7818e-4 / 0.566
unknown / 0.111
  1. Calculate the molar absorptivity () of the substance assuming the pathlength of the cuvette was 1 cm.
  1. Which of the following pairs of bonded atoms would be expected to have the longest bond length?
  2. C-N B. C-SC. C-B D. C-F
  3. Which of the descriptions below is the best representation of the energy change involved in the process of breaking bonds in a molecule? (ignore any subsequent bond formation that may occur)
  1. Always exothermic
  2. Always endothermic
  3. Net energy change is zero
  4. Exothermic or endothermic depending on conditions.
  1. How many sigma (σ) and pi(π) electron pairs are there in a carbon dioxide molecule?
  1. Two sigma, zero pi
  2. One sigma, one pi
  3. Two sigma, two pi
  4. Two sigma, one pi
  1. Which of the following elements is most likely to form compound involving an expanded valence shell of electrons?
  1. P B. NaC. OD. N
  1. Which of the following statements best describes the relationship between bond length and bond strength for a series of compounds involving bonds between the same two atoms?
  1. The greater the bond strength, the longer the bond.
  2. The greater the bond strength, the shorter the bond.
  3. Bond length and bond strength are not related
  4. The relationship between bond length and bond strength depends on other factors.
  1. Which of the following combinations of two elements is most likely to produce highly ionic bonds?
  1. Nitrogen and oxygen
  2. Nitrogen and fluorine
  3. Boron and nitrogen
  4. Lithium and fluorine
  1. Which of the following combinations of two elements is most likely to produce covalent bonds?
  1. nitrogen and oxygen
  2. oxygen and calcium
  3. sodium and nitrogen
  4. lithium and fluorine
  1. Which of the following salts is expected to have the highest melting point?
  1. NaF C. NaI
  2. NaClD. NaBr
  1. Predict which compound in each of the following pairs should have the higher melting point.

(i)NaClor RbCl(ii) NaCl or MgCl2

  1. Draw Lewis dot structures for the following:
  1. Ammonia
  2. Hydrogen cyanide
  3. N2O
  4. BrF5
  5. Ca3(PO4)2
  1. Based on the VSEPR theory, what is the molecular shape of PCl5?
  1. Linear
  2. tetrahedral
  3. Trigonal planar
  4. Trigonal bipyramidal
  1. Based on the VSEPR theory, which of the following corresponds most closely to the molecular shape of SCl2?
  1. Linear
  2. “T-shaped”
  3. bent (bond angle 120o)
  4. bent (bond angle 109.5o)
  1. A certain molecule has five structural electron pairs and the molecular structure is linear. How many lone pairs are present in this molecule?
  1. NoneC. Two
  2. One D. Three
  1. A certain molecule has six structural electron pairs and the molecular structure is a square pyramid. How many lone pairs are present in this molecule?
  1. NoneC. two
  2. One D. three
  1. What is the approximate Cl-B-Cl angle in BCl3?
  1. 90o C. 120o
  2. 109.5oD. 180o
  1. What is the approximate I-I-I angle in I3─?
  1. 90o C. 120o
  2. 109.5oD. 180o
  1. Which of the following best describes the variation of electronegativity of the elements with respect to their position on the periodic table?
  1. Increases across a period, increases down a group
  2. Increases across a period, decreases down a group
  3. Decreases across a period, increases down a group
  4. Decreases across a period, decreases down a group
  1. What is the formal charge on the O atoms in SO32─?
  1. 0 B. +1C. -1 D. +2
  1. What is the formal charge of the S atom in SO3?
  1. 0 B. +1C. -1 D. +2
  1. What is the average carbon-oxygen bond order in the formate ion?
  1. One C. Two
  2. 1 ½ D. 2 ½
  1. What is the average sulfur-oxygen bond order in SO3?
  1. One C. Two
  2. ½ D. 1 1/3
  1. In which species is the carbon-oxygen bond longer?
  1. B.
  1. Given the bond dissociation energies below, calculate the standard molar enthalpy of formation of NF3 (in KJ/mol).

½ N2(g) + 3/2 F2(g)  NH3(g)

Bond / Dissociation Energy (KJ/mol)
N≡N / 946
F─F / 159
N─F / 272
  1. 833 C. -104
  2. 440.D. -578
  1. Which of the bonds below is least polar?
  1. C─OC. C─N
  2. C─F D. C─B
  1. Which of the following molecules is polar?
  1. NCl3C. SF6
  2. O2 D. CS2
  1. Which of the following molecules is most likely to have a dipole moment?
  1. CH4C. SF6
  2. BeF2D. NF3
  1. Cyanic Acid has the electron dot structure below (you must add non-bonding pairs of electrons):

H─O─C≡N

  1. How many sigma (σ) bonds are there?
  2. How many pi(π) bond are there?
  3. What is the value of the C─O─H angle?
  4. What is the value of the N─C─O angle?
  1. Which of the following elements is most likely to display sp3d hybridization?
  1. OxygenC. Phosphorus
  2. NitrogenD. Carbon
  1. What type of hybrid orbital set is used by the nitrogen atom in the molecule NH3?
  1. sp B. sp2 C. sp3 D. sp3d2
  1. What type of hybrid orbital set is used by the xenon atom in the compound XeF4?
  1. sp B. sp2 C. sp3 D. sp3d2
  1. What hybrid orbital set is used by the nitrogen atom in the following molecule? You must add non-bonded pairs of electrons.

H3C─N═C═O

  1. sp B. sp2 C. sp3 D. sp3d2

CHAPTER 4.BONDING FUNCTION PURE

  1. Explain all of the variables in Coulomb’s law as it applies to intermolecular forces.
  2. What is the difference between a “bond” and an “intermolecular force”? Do you think an AP reader would give you credit if you used these words interchangeably? Why or why not?
  3. Identify the force of attraction that must be broken for the following substances. If the force of attraction is an intermolecular force, list all IMF’s involved. Please do not use “LD”, “HB” etc – spell it out! AP won’t accept acronyms unless they are defined first.

Substance / Covalent, Ionic, Metallic, or IMF? / If IMF, which one(s)
Brass
NH3
C6H12O6
HCl
SiO2
Si
CaCl2
PCl3

(an amino acid)
Na
NaCl
C3H6
  1. Answer the questions below regarding hydrogen iodide and iodine.
  2. Identify the intermolecular forces present in each of the pure substances.
  3. The melting points are -50.80oC for HI and 113.7oC for I2. Explain the large difference using bonding and intermolecular forces.
  4. Below is a table of the vapor pressure at 20oC for a variety of substances.

Pvapor at 20oC (KPa)
Pentane / 57.90
2-pentanone / 3.6
1-pentanol / 0.200
butane / 203
  1. Use intermolecular forces to explain the trend observed among the three 5-carbon substances. (You may want to look up their structural formulas online)
  2. Use intermolecular forces to explain the difference between butane and pentane.
  1. Use molecular structures to model both hydrogen-bonding possibilities in water. Label each atom with partial charges.
  2. Predict which substance is likely to have the highest boiling point in each of the following pairs. Justify your answer using Coulomb’s law.
  3. Na2O or CaOc. KF or KCl
  4. SrO or CaOd. RbCl or SrBr2

CHAPTER 5.BONDING FUNCTION MIXTURE

  1. Dissolving and mixing always involves an increase in entropy. Despite this favorable change, not all substances mix or dissolve due to unfavorable energetics. For each of the following pairs, indicate the forces of attraction that must be broken and formed in order for the two to mix.

PAIR / BROKEN / FORMED / NET ENERGY PREDICTION
Water and ethanol
Water and benzene
Water and copper(II) sulfate
Benzene and cyclohexane
  1. Use coulombs law to explain why the aluminum ion, Al3+, has a greater attraction to water than the gallium ion, Ga3+.
  2. Use structural formulas to show ethanol hydrogen bonding with water molecules. Make sure to show all possibilities. Include partial charges on each atom.
  3. Differentiate between P-type doping and N-type doping of silicon. Provide an example of each.
  4. Differentiate between interstitial and substitutional alloys. Look up an example of each and list how properties changed in the alloy compared to the pure metal.
  5. How does the radius of a metal atom solute affect the hardness of the metal solvent in alloys?

CHAPTER 6.REACTIONS

  1. What is characteristic of reactions that can be described by net ionic equations?
  2. What classifications (decomposition, synthesis, combustion, double replacement, single replacement, neutralization, redox) are often described by net ionic reactions?
  3. When potassium chlorate is decomposed, a gas forms that is able to re-ignite a glowing splint. Write the balanced reaction for the decomposition.
  4. Write the balanced chemical reactions (including units) for the following:
  5. Solutions of iron(III) perchlorate + calcium hydroxide.
  6. A piece of copper is added to a solution of lead(II) nitrate.
  7. Nitrous acid + aqueous potassium hydroxide
  8. Write net ionic equations for the reactions in #3.
  9. Write the balanced reaction equation for the combustion of the following substances
  10. C3H6
  11. C6H6
  12. C5H9OH
  13. Draw a particle diagram showing the synthesis reaction between nitrogen monoxide gas and oxygen gas to form nitrogen dioxide. Use 5 molecules of oxygen in your drawing.

CHAPTER 7.KINETICS