Download 1 2016-17 Honors Chemistry Review for the Final Exam Each unit

2016-17 Honors Chemistry Review for the Final Exam
Each unit has practice problems. Please review VIDEOS, your DILL and your Chemistry
Workbook.
2017 Final Exam Schedule
Fri., May 26 2B,3B
Tues, May 30: 3A
Wed. May 31: 1B, 4B
Thurs, June 1: 1A, 4A
Fri, June 2: 2A
Nomenclature
1.
Potassium chlorate _______________
8. MgS
________________________________
2.
Sulfuric acid
_______________
9. HgO
________________________________
3. Calcium carbonate
_______________
10. Cu2O ________________________________
4. Copper (I) sulfate
_______________
11. N2O3 ________________________________
5. Nitrogen trifluoride _______________
12. Au3PO4_______________________________
6. Tetraphosphorus decoxide
13. HNO3 ________________________________
____________
7. Hydroselenic acid _______________
14. HBr ________________________________
Metrics and Significant Figures
Convert the following to the designated unit while preserving significant figures.
1) 170.4 m3
=______________ cm3 (this is a cubed conversion…careful!)
2) 58 dg
=______________ mg
3) 0.0923 km
=______________ mm
4) 4.51 x 103 L
=______________ mL
5) 45700 cg
=______________pg
6) 82.4 nm
=______________ m
7) 2.00 x 10-5m
=______________km
8) 2.00mL/min
=______________L/hour
Determine the number of significant figures in the following numbers, then convert each to either decimal
or scientific notation.
9) 0.0090
________
___________
10) 3.200 x 104
________
___________
11) 0.3004
________
___________
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12) 4004
________
___________
13) 5.020 x 10 23
________
___________
Perform the following calculations while observing the correct number of significant figures.
14) 6.023•10 23 x 0.052 =
__________
15) 910 + 36 =
__________
16) (2.957 – 2.00)/0.0165 = __________
Chemical Reactions
Predict if the following reactions occur. If they occur, write and balance the formula equation with phase
notation. Then write the net ionic equation.
a. Solutions of lead (II) nitrate and sodium chloride are mixed.
b. Silver is added to a potassium chloride solution.
c. Aluminum is added to hydrochloric acid.
d. Solutions of barium acetate and lithium sulfate are mixed.
e. Chlorine gas is bubbled through aqueous ammonium iodide.
f.
Solutions of ammonium bromide and sodium hydroxide are mixed.
Redox Reactions
For each of the following molecular equations, write the net ionic equation. Then, identify the oxidizing
agent and the reducing agent. Then write both balanced half reactions.
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1.
Cl2 (g) + MgI2 (aq)  MgCl2 (aq) + I2 (s)
2.
Hg (l) + 2AgNO3 (aq)  Hg(NO3)2 (aq) + 2Ag (s)
3.
3Zn (s) + 2Ni(C2H3O2)3 (aq)  3Zn(C2H3O2)2 (aq) +2 Ni (s)
Mole Conversions
1. Determine the number of moles or grams present in each of the following:
(a) How many grams of silicon are present in a sample containing 13.58 moles of silicon?
(b) A sample of calcium has a mass of 10.0 grams. Determine the number of moles of calcium
present.
2. Determine the number of particles or moles in each of the following samples.
(a) How many atoms of neon are in 0.2030 moles of neon gas
(b) How many moles of magnesium chloride are found in a sample containing 4.50 x 1022 formula
units of magnesium chloride?
3. Determine the percent composition of KNO3.
4. Analysis of a compound shows that it contains 32.38% Sodium, 22.65 % sulfur, and 44.97% oxygen.
Find the empirical formula of this compound.
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5. Analysis of a 10.150g sample of a compound to contain only phosphorus and oxygen indicates a
phosphorus content of 4.433g. What is the empirical formula of this compound?
6. The percent composition of styrene glycol is 69.54% carbon, 7.30% hydrogen, and 23.16% oxygen.
What is its empirical formula? If the molecular mass of styrene glycol is 138.2 g/mol, what is its
molecular formula?
7. The mass of a hydrate was measured, heated to drive off the water of hydration, cooled and its final
mass was recorded. Find the formula for the hydrate:
DATA: 237.7 g of NiCl2• nH2O gave a residue of 129.6g after heating
Molarity
1. A 27.8-g sample of ammonium phosphate (NH4)3PO4, is dissolved into enough water to make 1.50L
of solution. Determine the molarity of this solution.
2. How many grams of potassium sulfate (K2SO4) are found in 4.50-L of a 0.775M solution?
3. A student has 45.0-g of aluminum chloride and wants to make a 2.50 M solution. Determine the
final volume of this solution.
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4. A solution is made by dissolving enough sodium chloride to make a 0.450M solution. A 175-mL
sample of this solution is then added to 250. mL of water for a final volume of 425 mL. Determine the
final concentration of this solution.
5. A 2.50-g sample of solid calcium nitrate is dissolved into enough water to make 25.0 mL of solution.
A 15.0-mL aliquot (sample) of this solution is then added to 75.0 mL of water. Determine the final
concentration of the resulting solution.
Stoichiometry
1) Hydrogen and oxygen gases react under a specific set of conditions to produce liquid water.
Balanced Equation:
a. How many moles of hydrogen would be required to produce 5.0 mol of water?
b. How many moles of oxygen would be required?
2) A 4.50 grams sample of the organic molecule ethane undergoes combustion.
a. How many grams of oxygen are required to consume the ethane?
b. How many grams of each product are formed?
c. Would it require more or less oxygen to combust a 4.50 gram sample of ethene? Explain?
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3) Sodium chloride is produced from its elements via a synthesis reaction. What mass of each reactant
would be required to produce 125 grams of sodium chloride?
4) Methanol (CH3OH) is an important industrial compound that is produced from the following
unbalanced reaction: CO(g) + H2 (g)  CH3OH (g)
a. What mass of each reactant would be needed to produce 100.0 kg of methanol?
b. If you were given 50.0 g of carbon monoxide, how many moles of hydrogen gas would be needed to
completely react it?
5) Sulfuric acid reacts with sodium hydroxide.
a. Predict the products of this reaction and balance the equation.
b. What mass of sulfuric acid would be required to react with 0.75 moles of sodium hydroxide?
c. What mass of each product is formed by this reaction?
6) Given the reactant amounts specified in each balanced chemical equation, determine the limiting
reactant in each case.
a. HCl
+ NaOH
2.0 mole
b. Zn
2.5 mol
6
+

NaCl + H2O

ZnCl2 + H2
2.5 mol
2HCl
6.0 mol
7) For each reaction in #6 above (a and b), determine the moles of excess reactant that remains.
8) A 2.50 gram sample of zinc and 5.50 gram of silver nitrate are available to react by single
displacement reaction.
a. Identify the limiting reactant. Don’t forget to write the balanced chemical equation first.
b. Determine the amount of moles of each product formed.
c. Determine the mass of each product formed.
9) The percent yield for the reaction PCl3 + Cl2  PCl5 is only 83.2 %. What mass of PCl5 is expected
from the reaction of 73.7 g of PCl3 with excess chlorine given the expected yield?
Atomic Structure
C = λν
E = hν
[C = 3.00 x 108 m/s
&
h = 6.626 2 x 10-34 J-s (or J/Hz)]
1. What is the wavelength of a wave having a frequency of 3.76 x 1014 s-1?
2. What is the frequency of a 6.9 x 10-13 m wave?
3. What is the wavelength of a 2.99 Hz wave?
4. What is the energy of a 7.66 x 1014 Hz wave?
5. What is the frequency of a wave carrying 8.35 x 10-18 J of energy?
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Complete the following table.
ELEMENT
LONGHAND
ORBITAL FILLING
QUANTUM
CONFIGURATION
DIAGRAM
NUMBER
Mg
3, 1, 0, +1/2
1s22s22p63s23p5
X-2 (identify)
1s22s22p6
Mo
Periodic Table
1. Identify the names of the following groups:
a. Group I ________________________________________
b. Group II _______________________________________
c. Group 17 ______________________________________
d. Group 18 ______________________________________
2. How do atomic radii change as we move…
(a) …from top to bottom in a group on the periodic table?
(b) …from left to right in the periodic table?
(c) Explain the theory for the size variation from top to bottom and left to right.
3. Arrange the following atoms in order of increasing electron affinity:
(i) B, Al, Si, and Ge.
(ii) Mg, F, P, O, and Ca.
(iii) In, Ar, Rb, Ge, and He.
4. Circle the following elements that have the largest ionization energy.
a. Na or Fr
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b. Na or S
c. Be or Mg
d. Cl or K
5. Which element are the following ions isoelectronic with? Arrange these from smallest to largest
radius.
N-3 ; O-2 ; F-1 ; Na+1 ; Mg+2 ; Al+3
6. List the oxidation state(s) for the following elements: Cl, K, Ar, Mg, Ga, Mn, Fe, O, Al, Hg, Sn, Si.
7. Based on their position in the periodic table, select the atom with the larger first ionization energy
from each of the following pairs and justify your answers:
(a) B, F (b) N, P (c) Hf, Cs, (d) O, P (e) Ga, Ge
8. Identify the following elements based on the description below.
a. The noble gas isoelectronic with the alkaline-earth metal ion from period 5
b. A radioactive element from the d block whose energy level is less than 5
c. The metal in Period 6 that is a liquid at room temperature
d. The halogen with the highest ionization energy
e. The actinide which has a downspin electron in the orbital designated a zero
Chemical Bonding and Intermolecular Forces
1) On another sheet of paper, draw the Lewis structure and identify the Shape, Polarity, and the
Bond angle for the following: NI3, SiO4-4, F2CO, and SeF6. This will help you review intermolecular
forces.
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2) Why is fluorine a gas, but iodine a solid when they both are at room temperature? Explain by
referencing IMF’s.
3) Identify and describe the bonds within the particle and the forces between particles for:
a) two molecules of SO2
b) two molecules of CF4
c) two molecules of H2O
d) two formula units of NaCl.
4) Draw the Lewis Dot structures for the following molecules. Make sure you show the dots representing
bonds directly between the bonding atoms and show all electrons paired. Then, describe the 3-D
shape of each and note any dipole moments as well as the most important IMF.
a) H2S
b) PCl3
c) NCl3
d) OF2
e) CH4
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f) OH-
g) CO2
h) NO3-
i) CN-
j) CO3 2-
k) PO4 3-
5. Rank the following compounds from lowest to highest boiling point: calcium carbonate, methane,
methanol (CH3OH), dimethyl ether (CH3OCH3). Justify your answer.
6. A solid is heated to a temperature of 105.3 °C and remains solid; a liquid is heated to 35°C and
remains liquid; and a gas is at room temperature (22°C). Which one has the most kinetic energy?
7.
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Which of the following molecules would have the highest vapor pressure? Justify your answer.
Cl2, NH3, or H2S
8. You have 2 unknown liquids. When you put 2 drops of each liquid on your lab table, unknown
liquid A evaporates very quickly while unknown liquid B does not. Based on this information,
answer the following questions.
a. Which liquid would have the highest boiling point? _________________
b. Which liquid would have the lowest vapor pressure? _______________
c. Which liquid would have the highest freezing point? ________________
d. Which liquid would have the highest viscosity? __________________
Phase Changes
1.
Determine the amount of heat needed to raise the temperature of 20.0 grams of ice at 0.0oC to
liquid water at 80.0oC.
2.
Find the heat lost by 20.0 grams of water initially at 40.0oC to form ice at - 20.0oC.
3.
How much heat energy is needed to raise the temperature of 50.0 grams of ice at - 10.0oC to form
steam at 120.0oC?
4.
Xylene has the following constants: M.P. = 25oC, B.P. = 80oC, Hf = 120 cal/g, Hvap = 360 cal/g,
Specific Heat(solid) = 0.25 cal/goC, Specific Heat(liq) = 0.30 cal/goC, Specific Heat(gas) = 0.20 cal/goC.
How much heat is given off (in calories) when 40.0 grams of xylene has its temperature dropped
from 100.0oC, to 0.00oC?
Organic Chemistry – Nomenclature
1) For the following names, draw the molecule. You do not have to include the hydrogen atoms.
a. Butane
b. 2,2-dimethyl hexane
c. Hexene
d. 2-heptyne
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2) What is the difference between unsaturated and saturated hydrocarbons?
Gases
1. A gas is contained in a cylinder with a volume of 600.mL at a temperature of 40.0 C and a pressure
of 740.mmHg. The gas is then compressed to a volume of 13.0 mL, and the temperature is lowered
to 18.0 C. What is the pressure of the gas in atm?
2. A balloon is filled to a volume of 300.mL at a temperature of 18.0 C. The balloon is cooled at a
constant pressure to a temperature of 150.0 K. What is the final volume of the balloon in milliliters?
3. You have a tank of argon gas at 23.5 atm pressure. The volume of argon in the tank is 150.0 L.
What would be the volume of this gas (in liters) if you allowed it to expand to the pressure of the
surrounding air (0.974 atm)? Assume temperature remains constant.
4. The maximum safe pressure that a certain 13.0L vessel can hold is 15.3 atm. If the vessel contains
1.65 moles of gas, what is the maximum temperature to which this vessel can be subjected?
5. A 0.109 gram sample of a pure gaseous compound occupies 112 mL at 100.°C and 750 torr. What is
the molecular mass of the compound?
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6. You are conducting an experiment much like the one you did in the lab where you collected the gas
from the lighter by water displacement. This time, you collect 5.65 g of a halogen gas by displacing
liquid carbon tetrachloride from a flask (instead of water). The temperature of the liquid and gas is
23.0 °C and the pressure in the lab is 1140 mmHg. The amount of carbon tetrachloride displaced is
1.43 L. The vapor pressure of the carbon tetrachloride at this temperature is 114 mmHg. Determine
the molar mass and the identity of the halogen.
Bond Energy
Bond type
Bond Energy, kJ/mol
Bond type
Bond Energy, kJ/mol
H-H
436
C-O
358
O-O
146
C=O
799
O=O
498
C-H
413
O- H
463
1. Using bond energies, calculate the ΔHrxn for each of the following:
a) 2 H2 (g) + O2 (g) -------> 2 H2O (l)
b) CH4 (g) +
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2 O2 (g) ------->
CO2 (g) + 2 H2O (l)
Hess’s Law
2. Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values:
2CO2(g) + H2O(g) → C 2H2(g) + 5/2O2(g)
C2H2(g) + 2H2(g) → C2H6(g)
H2O(g) → H2(g) + 1/2O2 (g)
C2H6(g) + 7/2O2(g) → 2CO2(g) + 3H2O(g)
ΔH =-94.5 kJ
ΔH =71.2 kJ
ΔH =-283 kJ
AMHS Honors Chemistry Super Problem:
I recommend this to all students that will be taking AP Chemistry next year! (Or those who actually enjoy a
challenge!)
A student is synthesizing a compound of osmium metal with oxygen gas which is accomplished by heating the metal in
the presence of oxygen gas at temperatures near 400˚C. An 85.57g sample of osmium metal is heated in a crucible
(crucible mass = 254.48g and lid mass = 66.79g). The final product and the crucible w/lid have a total mass of 435.67g.
The osmium/oxygen product has a known or accepted molar mass of 254.23 g/mol. The melting and boiling points of
the product are determined to be 40.25˚C and 129.7˚C respectively.
a) Determine both the empirical and molecular formulas for this product.
b) This is a product made from a metal and a nonmetal. However, the bonds have a more covalent character to them
than ionic. Name this product using both the ionic and covalent nomenclature systems.
c)
What does the information in “b” above indicate to you about the elements respective electronegativity in
comparison to each other?
d) Oxygen has an electronegativity of 3.5 which is higher than that of osmium metal. What is the minimum E.N.
value for osmium if the bond can still be considered covalent and not ionic?
e) Write the full electron configuration of osmium. Identify the core and valence electrons in this configuration.
What would its oxidation number be if the product was truly an ionic compound?
f)
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The shape of the “molecule” is found to be tetrahedral. Assuming this is true, what are the bond angles in this
molecule? Why aren’t the bond angles 90˚ if the central atom has four atoms around it and no lone pairs? Explain.
g) The molecule is known to be 518 times more soluble in carbon tetrachloride than water. Does this support or
conflict with the assumption about the shape of the molecule from “f” above? Explain your statement by clearly
discussing the involved chemicals. What does this indicate about the type of intermolecular force available to this
molecule? How could the molecule be soluble at all in water if it is so much more soluble in carbon tetrachloride?
What could be added to the water to increase the solubility of the osmium/oxygen species? Explain.
h) Consider the melting and boiling points of the product. Are the intermolecular forces less than, equal to, or
greater than that of the hydrogen bonded water molecules? Explain.
The osmium/oxygen compound synthesized above occupies a volume of 1.46 x 10 16 µm3 at a temperature of 138.7˚C and a
pressure of 110. kPa.
i)
Determine the density of this compound in g/cm3 and then g/L.
j)
Confirm the molar mass stated above by calculating the molar mass using the “Dirt Pile” gas law equation.
Calculate the percent error assuming the given molar mass is the accepted value.
k) Determine the volume that 0.588g of this gas would occupy if it were collected over water under the conditions of
755 torr, 22.5˚C. The vapor pressure of water at this temperature is 20.5 mmHg. Determine the volume of this
sample under STP conditions.
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