HONORS CHEMISTRY Spring 2013 FINAL EXAM study guide
The exam is consists of 100 multiple choice questions (100 points) I will provide a periodic table, a table of E° values, and this information:
The Structure of Matter
1. List the number of protons, neutrons, & electrons in each species: 34S2– 71Ga 56Fe3+
2. Label a blank periodic table with these items:
a. period and group numbers d. alkali metals, alkaline earths, halogens, noble gases
b. the staircase e. trends in atomic size
c. metals, nonmetals, semimetals f. the s, p, and d sublevel regions for e– configurations
(metalloids), and transition metals
3. Draw a diagram of the first 6 energy levels available to the electron in hydrogen, according to Bohr’s model, and use colored pencils to illustrate the transitions associated with the photons we can see in the bright line spectrum. Why can’t we see other transitions?
4. a. What is the difference between an orbit and an orbital?
b. How are a 3s orbital and a 3p orbital alike? How are they different?
c. How are a 2s orbital and a 4s orbital alike? How are they different?
d. There are three orbitals in a p sublevel. How are they alike? How are they different?
e. How many electrons can occupy an orbital? How do these electrons differ from each other?
5. Draw an orbital diagram and write the e– configuration (core notation is fine) for each of these atoms: Si Ca Al S Ni
6. Give name & symbol for each element:
a. period 3, group 14 or 4A g. period 4 element that ends in s2p4
b. period 2, group 16 or 6A h. period 5 element that ends in s2
c. period 4 element that ends in d6 i. period 2 element with 2 unpaired p electrons
d. period 6 alkali metal j. period 3 element with 3 valence electrons
e. period 5 halogen k. period 5 transition metal
f. period 1 noble gas l. period 4 semimetal
7. Predict the Lewis symbols for iodine (I), strontium (Sr), and arsenic (As).
8. a. What is the octet rule?
b. Use octet rule and Lewis dot structures to show the formation of the salts NaBr and MgO.
c. How are the dot structures for salts different from the dot structures for molecules?
9. Write a Lewis dot structure for
CCl4 H2O CO2 NH3 CH2O HCN
The Reactions of Matter
10. Balance each equation in lowest terms & list the coefficients.
a. Cr2O3 (s) + C (s) ® Cr (s) + CO (g)
b. Na (s) + H2O (l) ® NaOH (aq) + H2 (g)
c. C5H12 (l) + O2 (g) ® CO2 (g) + H2O (g)
d. NH3 (g) + O2 (g) ® NO (g) + H2O (g)
11. Estimate the molar mass of (whole numbers, with units): a. As2O5 b. Pb(NO3)2
12. Mole conversions:
a. What is the mass of 2.37 moles NaCl?
b. How many moles are in 14 grams CaF2?
13. Stoichiometry:
a. How many moles of oxygen gas are needed to react with 0.875 moles of NH3?
4 NH3 (g) + 5 O2 (g) ® 4 NO (g) + 6 H2O (l)
b. How many grams of water form when 0.18 moles of NH3 form?
2 NO2 (g) + 7 H2 (g) ® 2 NH3 (g) + 4 H2O (l)
c. How many grams of carbon dioxide gas form when 266 g of octane (C8H18) burns?
2 C8H18 (l) + 25 O2 (g) ® 16 CO2 (g) + 18 H2O (g)
14. Classify each of these reactions as exothermic or endothermic and indicate the sign of ∆H for each. If you touched the side of each reaction vessel, would it feel hot or cold?
a. H2SO4 (aq) + 2 KOH (aq) ® 2 H2O (l) + K2SO4 (aq) + energy
b. energy + NH4NO3 (s) ® NH41+ (aq) + NO31– (aq)
15. Explain these observations about reaction rates in molecular terms:
a. Increasing the concentration of a reactant usually increases reaction rate.
b. Increasing the temperature always increases the reaction rate.
16. Draw energy diagrams for exothermic and endothermic reactions. Label the reactants, products, ∆H, and activation energy. Show how a catalyst affects the reaction rate and the value of ∆H.
17. How do we recognize equilibrium? How do we explain equilibrium?
18. Energy + 2 NH3(g) à N2(g) + 3H2 (g)
a. Predict the effect of the position of equilibrium (shift left, shift right, no effect) of
i. removing NH3 iii. decreasing the container volume
ii. increasing the temperature iv. adding a catalyst
b. Suggest three different ways to increase [NH3] at equilibrium (that is, shift equilibrium to the left).
Phases of Matter: Gases, Liquids, and Solutions
19. What aspect of molecule behavior causes what we perceive as pressure? What aspect of molecular motion is directly related to temperature?
20. Imagine two identical containers at the same temperature, one filled with He gas and the other with Ar gas. Which sample contains (a) the greater number of atoms? (b) the heavier atoms? (c) the faster atoms? (d) the atoms with the greater average kinetic energy?
21. Gas law problems:
a. A sample of gas occupies 1.50 L at 0.967 atm. What volume will it occupy if the pressure on the sample increases to 1.02 atm? Assume temperature and amount of gas remain constant.
b. A 425 mL balloon at 27 °C is cooled to 4 °C. What is the volume of the cold balloon?
c. An aerosol can contains propellant at room temperature (25 °C) and 1.5 atm pressure . The can will burst if the pressure inside increases to 3.0 atm. At what temperature will the can burst?
d. What is the volume of 0.150 mol of nitrogen at 28 °C and 935 mm Hg?
e. How many moles are in 35.8 mL of hydrogen gas at 21 °C and 733 mm Hg?
22. Moles & stoichiometry
a. Calculate the number of moles of O2 in 35.7 mL of O2 at STP
b. What is the volume of 0.167 mol He at STP?
c. How many grams of KClO3 are needed to produce 41 mL of O2 at STP?
2 KClO3 (s) ® 2 KCl (s) + 3 O2 (g)
d. How many mL of H2 at STP form when 0.054 g Al react?
2 Al (s) + 6 HCl (aq) ® 2 AlCl3 (aq) + 3 H2 (g)
23. If you decrease the temperature of a real gas enough, it will form a liquid. Why?
24. a. Identify the solute, solvent, and solution in salty water.
b. Substances that produce ions when dissolved in water are called _____ . How can you tell whether a substance contains ions?
25. Draw a heating curve for water. Label solid, liquid, gas, melting, freezing, and boiling. Show how the heating curve would be affected if something like salt or sugar were dissolved in the water.
26. Write a balanced chemical equation for the dissociation of each salt when it dissolves in water:
a. KBr b. ZnCl2 c. Na2SO4
27. More moles & stoichiometry
a. How many moles of solute are in 13.4 mL of 0.250 M KI?
b. How many mL of 0.25 M NaCl are needed to provide 0.075 moles NaCl?
c. How many grams of Mg3(PO4)2 form if 25 mL of 0.086 M MgCl2 are consumed in this reaction:
3 MgCl2 (aq) + 2 K3PO4 (aq) ® Mg3(PO4)2 (s) + 6 KCl (aq)
d. If 0.256 g H2C2O4 are neutralized by 15.63 mL of NaOH, what is the concentration of the NaOH solution, in mol/L?
H2C2O4 (s) + 2 NaOH (aq) ® 2 H2O (l) + Na2C2O4 (aq)
Acids & bases
28. a. What ion is produced when an acid dissolves in water? Write a chemical equation, using HF.
b. What ion is produced when a base dissolves in water? Write a chemical equation showing how NaOH dissolves in water to form a basic solution.
c. How can NH3 produce a basic solutions? Write a chemical equation.
29. Complete and balance these neutralization reactions:
a. HC2H3O2 (aq) + NaOH (aq) ® b. HNO3 (aq) + KOH (aq) ®
30. If you had a 0.1 M solution of a strong acid and a 0.1 M solution of a weak acid, what difference(s) would you be able to observe in the lab? How do you account for these differences?
31. Convert these [H3O1+] to pH:
a. 0.0065 M b. 0.125 M c. 3.85 x 10–7 M d. 7.4 x 10–10 M
32. Convert these pH values to [H3O1+]: a. pH 5.35 b. pH 1.4 c. pH 11.7 d. pH 8.26
33. Draw a pH line and mark neutral, acidic, and basic regions. Label the line to show [H3O1+] increasing and [H3O1+] decreasing. Classify these solutions as very basic, somewhat basic, neutral, somewhat acidic, or very acidic:
a. pH 6.2 b. pH 11.6 c. pH 7.0 d. pH 2.1 e. pH 8.5
Electrochemistry
34. Balance each reaction. Label the oxidation and reduction half-reactions and determine the overall E°.
a. Al (s) + Ni2+ (aq) ® Al3+ (aq) + Ni (s)
b. Ag (s) + Sn4+ (aq) ® Ag1+ (aq) + Sn2+ (aq)
c. Fe (s) + MnO41– (aq) ® Fe2+ (aq) + Mn2+ (aq) (in acid solution)
35. Identify the anode and cathode of an electrochemical cell made of Fe (s) in Fe2+ (aq) connected to
Ni (s) in Ni2+ (aq), then draw a diagram of the cell and calculate E° for the cell.
a. Label the anode, cathode, and salt bridge c. Show the direction of electron flow.
b. Write the ½ -reaction at each electrode d. Show the direction of ion movement.
and identify it as oxidation or reduction e. Describe the function of the salt bridge
Answers to 2012-2013 final exam study guide
The Structure of Matter
1. 34S2– has 16 p+, 18 n0, and 18 e– Ga has 31 p+, 40 n0, and 31 e– 56Fe3+ has 26 p+, 30 n0, and 23 e–
2. See pg 157-171 in text.
3. Diagram is at right; 6®2 is violet, 5®2 is violet, 4®2 is turquoise blue, 3®2 is red
4. a. An orbit is a fixed path that an electron follows; the position and trajectory of the electron are known at every point in an orbit. An orbital is a probability region within which you could expect to find the electron, but its exact location and trajectory within the region cannot be known. Orbital is a more accurate picture of the dual nature (wave/particle) of the electron.
b. 3s and 3p extend the same distance from the nucleus (level 3) but they have different shapes: s is spherical and p is propeller shaped.
c. 2s and 4s are the same shape (spherical) but 4s is larger (extends farther from the nucleus).
d. The 3 orbitals in a p sublevel are the same shape (propeller) and size; they have different orientations in space (one along the x axis, one along y, and one along z).
e. Two electrons of opposite spin can occupy an orbital.
5. The configurations are given below. See me if you need help on the orbital diagrams. Remember that where several orbitals are of equal energy are available, the electrons will spread out, then pair up, so Si, S, and Ni each have two unpaired electrons in their outer sublevel.
Si 1s2 2s2 2p6 3s2 3p2 S 1s2 2s2 2p6 3s2 3p4
Ca 1s2 2s2 2p6 3s2 3p6 4s2 Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8
Al 1s2 2s2 2p6 3s2 3p1
6. a. Si b. O c. Fe d. Cs e. I f. He g. Se h. Sr i. C or O j. Al
k. any from Y through Cd l. Ge or As
7. Iodine should have seven dots, Sr two, As five (one pair and three singles).
8. Octet rule describes bonding behavior in main group elements. Atoms will lose, gain, or share valence electrons to achieve the same number of electrons as the nearest noble gas.
9.
The Reactions of Matter
10. a. 1, 3, 2, 3 b. 2, 2, 2, 1 c. 1, 8, 5, 6 d. 4, 5, 4, 6
11. a. 230 g/mol b. 331 g/mol
12. a. 139 g NaCl b. 0.18 mol CaF2
13. a. 1.09 mol O2 b. 6.5 g H2O c. 820 g CO2
14. a. exothermic, ∆H negative, hot b. endothermic, ∆H positive, cold
15. a. Increasing the concentration of a reactant means more opportunities for effective collisions between reactants, so more product forms per second.
b. Increasing the temperature means more molecules have enough kinetic energy (activation energy) to react when they collide.
16. In every reaction, the reactant molecules must overcome an energy barrier (the activation energy). At higher temperature, more molecules have enough energy to overcome the barrier, so the reaction proceeds faster. A catalyst lowers the activation energy of the reaction (by providing an alternative pathway), allowing the reaction to proceed faster without raising the temperature. A catalyst does not affect the magnitude or sign of ∆H.
EXOTHERMIC ENDOTHERMIC
17. We recognize equilibrium by the presence of constant macroscopic properties such as pressure, concentration, temperature, color intensity, etc. We explain equilibrium as the balance of two opposing processes such as evaporating/condensing, dissolving/crystallizing, or forward reaction/reverse reaction.
18. a. i. left ii. right iii. left iv. no effect
b. add N2 or H2, decrease the temperature, or decrease the volume.
Phases of Matter: Gases, Liquids, and Solutions
19. Gas molecules are in constant motion and collide with the walls of their container. The collective “push” from all these tiny collisions is what we measure as pressure. Absolute (Kelvin) temperature is directly related to molecular kinetic energy. Kinetic energy = ½ mv2 (mass x velocity squared).