Download Semester Exam Review

Kinetics
1. Know the difference between a RATE EXPRESSION and a RATE LAW.
2. Know how to determine reaction orders so to obtain the exponents of the rate law.
3. Know the proper integrated rates law and half-life equations, the units of k, and the graph of a zero, first , and
second order process.
4. Understand the significance of activation energy and how a catalyst affects it.
5. Be able to determine a appropriate mechanism for a reaction.
a. Must added up to the overall equation.
b. Must be uni- or bi-molecular.
c. Must be consistent with the experimental rate law.
6. Remember that the rate law for an overall reaction IS NOT stoichiometric, but the rate law for the individual
steps in a mechanism is; always stop at the slow step.
Equilibrium
1. Be able to determine Q from INITIAL CONCENTRATIONS.
2. Determine K from EQUILIBRIUM CONCENTRATIONS.
3. Compare Q vs K when necessary and determine which way the reaction shifts so to know which side to
subtract “x” from.
a. If Q > K, then too many products; shift left and subtract “x” from products.
b. If Q < K, then too many reactants; shift right and subtract “x” from reactants.
c. If Q = K, then at equilibrium.
4. Ignore the “-x” and test with the 5% rule (same a percent dissociation = x / [A] initial x 100). If it fails, use the
quadratic.
5. Be able to convert between Kp and Kc.
6. Know how to do Le Chatelier’s Principle!!!!!!!!!!!!!!!!!!!!!
Acids and Base
1. Know the definitions of acids and bases, especially the Bronsted definition.
2. Ka and Kb problems are the same as regular equilibrium problems, except for the rule of 400 = [A]initial/Ka.
a. If greater than 400, ignore “-x”
b. If less that 400, use quadratic.
3. Ka x Kb = Kw
4. pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14 [H3O+] [OH-] = 1 x 10-14
5. Know the pH scale.
6. Know your strong acids and strong bases!!!!!!
Neutralization
1. Know how to do buffer problems.

When you ADD something to a solution, do a stoichiometry table first using millimoles and total
volumes. Then construct an equilibrium table or use the Henderson Hasselbach Equation.
2. Know how to do a strong acid-strong base titration, a strong acid-weak base titration, and a weak acid-strong
base titration.
3. Ksp Problems.
4. Selective Precipitation Problems.
Thermochemistry
1. 1st law of Thermodynamics – Euniverse = Esystem + Esurroundings = 0.
a. q = mcT.
b. –q lost = + q gained.
c. Hess’s Law of Enthalpy.
d. The Summation of Enthalpy (products – reactants).
e. The Enthalpy of Bond Energies (bonds broken – bonds formed).
Molecular Structure
1. Know Lewis, VSEPR, and hybridization.
2. Know why and when the octet rule can be violated.
3. Know how to determine bond and molecular polarity.
Reaction Writing
1. Look for combustion.
2. Know solubility rules for precipitation reactions.
3. Strong acids and strong bases are written as ions.
4. Only remove or add one proton, except when you have an excess of H+
5. Redox.
a. Acidified.
b. Metal/metal cation or nonmetal/nonmetal anion.
c. Cr type ions.
d. Mn type ions.
e. Must have both oxidation and reduction to be redox.
2. Metal oxide and water = strong base.
3. Nonmetal oxide and water = strong acid.
4. Complex Ions.

NH3, CN-, SCN-, OH-, H2O
5. Hydrides, alkali metals and alkaline earth metals form H2 and a base when placed in water.
6. Know simple organic chem.
Stoichiometry
1. Empirical Formulas and Molecular Formulas.
2. Combustion Analysis.
3. Limiting Reagents.
4. Molar mass = grams per mole
5. Net ionic equations.
6. Balancing regular and redox equations (know oxidation numbers)!
7. Solution Stoichiometry.
Acids and Bases
(a) What is the pH of a 2.0 molar solution of acetic acid. Ka acetic acid = 1.8x10-5. What
is the percent dissociation of the acid?
(b) A buffer solution is prepared by adding 0.10 liter of 2.0 molar acetic acid solution to
0.1 liter of a 1.0 molar sodium hydroxide solution. Compute the hydrogen ion
concentration and the pH of the buffer solution.
(c) Suppose that 0.10 liter of 0.50 molar hydrochloric acid is added to 0.040 liter of the
buffer prepared in (b). Compute the hydrogen ion concentration of the resulting
solution.
Equilibrium and LeChatelier’s Principle
2 H2S(g)  2 H2(g) + S2(g)
When heated, hydrogen sulfide gas decomposes according to the equation above. A 3.40
g sample of H2S(g) is introduced into an evacuated rigid 1.25 L container. The sealed
container is heated to 483 K, and 3.7210–2 mol of S2(g) is present at equilibrium.
(a) Write the expression for the equilibrium constant, Kc, for the decomposition reaction
represented above.
(b) Calculate the equilibrium concentration, in molL-1, of the following gases in the
container at 483 K.
(i) H2(g)
(ii) H2S(g)
(c) Calculate the value of the equilibrium constant, Kc, for the decomposition reaction at
483 K.
(d) Calculate the partial pressure of S2(g) in the container at equilibrium at 483 K.
(e) For the reaction H2(g) +
equilibrium constant, Kc.
1
2
S2(g)  H2S(g) at 483 K, calculate the value of the
For the system 2 SO2(g) + O2(g)  2 SO3(g) , H is negative for the production of SO3.
Assume that one has an equilibrium mixture of these substances. Predict the effect of
each of the following changes on the value of the equilibrium constant and on the number
of moles of SO3 present in the mixture at equilibrium. Briefly account for each of your
predictions. (Assume that in each case all other factors remain constant.)
(a) Decreasing the volume of the system.
(b) Adding oxygen to the equilibrium mixture.
(c)
Raising the temperature of the system
Kinetics
For a hypothetical chemical reaction that has the stoichiometry 2 X + Y  Z, the
following initial rate data were obtained. All measurements were made at the same
temperature.
Initial Rate of
Formation of Z, Initial [X]o, Initial [Y]o,
(mol.L-1)
(mol.L-1)
(mol.L-1.sec-1)
7.010-4
0.20
0.10
1.410-3
0.40
0.20
-3
2.810
0.40
0.40
4.210-3
0.60
0.60
(a) Give the rate law for this reaction from the data above.
(b) Calculate the specific rate constant for this reaction and specify its units.
(c) How long must the reaction proceed to produce a concentration of Z equal to 0.20
molar, if the initial reaction concentrations are [X]o = 0.80 molar, [Y]o = 0.60 molar
and [Z]0 = 0 molar?
(d) Select from the mechanisms below the one most consistent with the observed data,
and explain your choice. In these mechanisms M and N are reaction intermediates.
(1) X + Y  M
(slow)
X+MZ
(fast)
(2) X + X  M
(fast)
Y+MZ
(slow)
(3) Y  M
(slow)
M+XN
(fast)
N+XZ
(fast)
2 HI(g)  H2(g) + I2(g)
The following data give the value of the rate constant at various temperatures for the
gas phase reaction above.
T (K)
k (litre/mol sec)
647
8.5810-5
666
2.1910-4
683
5.1110-4
700
1.1710-3
716
2.5010-3
Describe, without doing any calculations, how a graphical method can be used to obtain
the activation energy for this reaction.
(e)
(f)
H2(g) + I2(g)  2 HI(g)
For the exothermic reaction represented above, carried out at 298K, the rate law is as
follows.
Rate = k[H2][I2]
Predict the effect of each of the following changes on the initial rate of the reaction and
explain your prediction.
(i) Addition of hydrogen gas at constant temperature and volume
(ii) Increase in volume of the reaction vessel at constant temperature
(iii) Addition of catalyst. In your explanation, include a diagram of potential energy
versus reaction coordinate.
Ksp
The solubility of iron(II) hydroxide, Fe(OH)2, is 1.4310-3 gram per litre at 25C.
(a) Write a balanced equation for the solubility equilibrium.
(b) Write the expression for the solubility product constant, Ksp, and calculate its value.
(c) Calculate the pH of a saturated solution of Fe(OH)2 at 25C.
(d) A 50.0 millilitre sample of 3.0010-3 molar FeSO4 solution is added to 50.0 millilitres
of 4.0010-6 molar NaOH solution. Does a precipitate of Fe(OH)2 form? Explain and
show calculations to support your answer.
Neutralization
An approximately 0.1-molar solution of NaOH is to be standardized by titration. Assume
that the following materials are available.
• Clean, dry 50 mL buret
• 250 mL Erlenmeyer flask
• Wash bottle filled with distilled water
• Analytical balance
• Phenolphthalein indicator solution
• Potassium hydrogen phthalate, KHP, a pure solid monoprotic acid (to be used as
the primary standard)
(a) Briefly describe the steps you would take, using the materials listed above, to
standardize the NaOH solution.
(b) Describe (i.e., set up) the calculations necessary to determine the concentration of
the NaOH solution.
(c) After the NaOH solution has been standardized, it is used to titrate a weak
monoprotic acid, HX. The equivalence point is reached when 25.0 mL of NaOH
solution has been added. In the space provided at the right, sketch the titration curve,
showing the pH changes that occur as the volume of NaOH solution added increases
from 0 to 35.0 mL. Clearly label the equivalence point on the curve.
(d) Describe how the value of the acid-dissociation constant, Ka, for the weak acid HX
could be determined from the titration curve in part (c).
(e) The graph below shows the results obtained by titrating a different weak acid, H2Y,
with the standardized NaOH solution. Identify the negative ion that is present in the
highest concentration at the point in the titration represented by the letter A on the
curve.
Thermochemistry
Propane, C3H8, is a hydrocarbon that is commonly used as fuel for cooking.
(a) Write a balanced equation for the complete combustion of propane gas, which yields
CO2(g) and H2O(l).
(b) Calculate the volume of air at 30C and 1.00 atmosphere that is needed to burn
completely 10.0 grams of propane. Assume that air is 21.0 percent O2 by volume.
(c) The heat of combustion of propane is -2,220.1 kJ/mol. Calculate the heat of
formation, Hf, of propane given that Hf of H2O(l) = -285.3 kJ/mol and Hf of
CO2(g) = -393.5 kJ/mol.
(d) Assuming that all of the heat evolved in burning 30.0 grams of propane is transferred
to 8.00 kilograms of water (specific heat = 4.18 J/g.K), calculate the increase in
temperature of water.
Molecular Structure
Answer the following questions about the element selenium, Se (atomic number 34).
(a) Samples of natural selenium contain six stable isotopes. In terms of atomic structure,
explain what these isotopes have in common, and how they differ.
(b) Write the complete electron configuration (e.g., 1s2 2s2... etc.) for a selenium atom in
the ground state. Indicate the number of unpaired electrons in the ground-state atom,
and explain your reasoning.
(c) In terms of atomic structure, explain why the first ionization energy of selenium is
(i) less than that of bromine (atomic number 35), and
(ii) greater than that of tellurium (atomic number 52).
(d) Selenium reacts with fluorine to form SeF4. Draw the complete Lewis electron-dot
structure for SeF4 and sketch the molecular structure. Indicate whether the molecule
is polar or nonpolar, and justify your answer.
Consider the molecules PF3 and PF5.
(a) Draw the Lewis electron-dot structures for PF3 and PF5 and predict the molecular
geometry of each.
(b) Is the PF3 molecule polar, or is it nonpolar? Explain.
(c) On the basis of bonding principles, predict whether each of the following compounds
exists. In each case, explain your prediction.
(i) NF5
(ii) AsF5
Stoichiometry
Water is added to 4.267 grams of UF6. The only products are 3.730 grams of a solid
containing only uranium, oxygen and fluorine and 0.970 gram of a gas. The gas is 95.0%
fluorine, and the remainder is hydrogen.
(a) From these data, determine the empirical formula of the gas.
(b) What fraction of the fluorine of the original compound is in the solid and what
fraction in the gas after the reaction?
(c) What is the formula of the solid product?
(d) Write a balanced equation for the reaction between UF6 and H2O. Assume that the
empirical formula of the gas is the true formula.
Answer the following questions about BeC2O4(s) and its hydrate.
(a) Calculate the mass percent of carbon in the hydrated form of the solid that has the
formula BeC2O4•3H2O.
(b) When heated to 220.C, BeC2O4•3H2O(s) dehydrates completely as represented
below.
BeC2O4•3H2O(s)  BeC2O4(s) + 3 H2O(g)
If 3.21 g of BeC2O4•3H2O(s) is heated to 220.C calculate
(i) the mass of BeC2O4(s) formed, and,
(ii) the volume of the H2O(g) released, measured at 220.C and 735 mm Hg.
(c) A 0.345 g sample of anhydrous BeC2O4, which contains an inert impurity, was
dissolved in sufficient water to produce 100. mL of solution. A 20.0 mL portion of
the solution was titrated with KMnO4(aq). The balanced equation for the reaction that
occurred is as follows.
16 H+(aq) + 2 MnO4-(aq) + 5 C2O42-(aq)  2 Mn2+(aq) + 10 CO2(g) + 8 H2O(l).
The volume of 0.0150 M KMnO4(aq) required to reach the equivalence point was
17.80 mL.
(i) Identify the reducing agent in the titration reaction.
(ii) For the titration at the equivalence point, calculate the number of moles of each
of the following that reacted.
• MnO4-(aq)
• C2O42-(aq)
(iii) Calculate the total number of moles of C2O42-(aq) that were present in the 100.
mL of prepared solution.
(iv) Calculate the mass percent of BeC2O4(s) in the impure 0.345 g sample.
Reactions
a) magnesium metal is burned in nitrogen gas.
b) sulfur dioxide gas is passed over solid calcium oxide.
c) lead foil is immersed in silver nitrate solution.
d) a solution of ammonium sulfate is added to a saturated solution of barium hydroxide.
e) acetic acid solution is added to a solution of sodium hydrogen carbonate.
f) solid sodium dichromate is added to an acidified solution of sodium iodide.
g) a drop of potassium thiocyanate is added to a solution of iron(III) chloride.
a)
b)
c)
d)
solid calcium is added to warm water.
powdered magnesium oxide is added to a container of carbon dioxide gas.
gaseous hydrogen sulfide is bubbled through a solution of nickel(II) nitrate.
excess concentrated sodium hydroxide solution is added to solid aluminum
hydroxide.
e) solid silver is added to a dilute nitric acid (6M) solution.
f) excess potassium hydroxide solution is added to a solution of potassium dihydrogen
phosphate.
g) hydrogen peroxide solution is added to a solution of iron(II) sulfate.
h) propanol is burned completely in air.