Download Chemistry EOC Review

name: _____________________
date: __________
Chemistry EOC Review
1) Express the following numbers in scientific notation:
a. 60200 L
b. 4520 J
c. 0.00600 mg
d. 0.01023 km
e. 80.70 mL
2) How many significant figures are in each of the numbers from #1?
3) Convert the following (remember K H D – D C M).
a. 682 kg  cg
b. 5.0 m  mm
c. 1250 cL  kL
d. 54.8 mm  dm
4) Record the volume of water in the graduated cylinder shown to the right,
to the appropriate number of significant figures.
5) List four evidences of a chemical change.
6) You have an unknown silver-colored metal in front of you on the lab
table. Using a balance, you find its weight to be 24.6 g. You use a
graduated cylinder and water and find the volume of this metal to be 4.28
ml. Determine the density of this metal, and using your EOC reference
table, identify the unknown metal.
7) In front of you is an unknown white solid. You discover that it melts at about 267C. Using
your EOC reference table, identify this unknown white solid. (Note units for temperature!)
8) List three examples of physical changes and three examples of chemical changes you might
observe happening in our classroom.
9) Assume you collect an unknown gas, by water displacement, as part of your lab. For each of
the following gases, write down the name of the gas, and identify a test which could be used
to identify whether this was your unknown gas: H2, O2, H2O, CO2 (two tests). Specifically
expect what you would expect to see when you performed the test (what color change, etc.).
10) Briefly explain the contributions of the following scientists toward the development of
current atomic theory (including a brief description of their experiment, where appropriate):
a. Democritus
b. John Dalton
c. Joseph John Thomson
d. Robert Millikan
e. Ernest Rutherford
f. James Chadwick
g. Max Planck
h. Niels Bohr
11) The protons are ____________ charged particles found in the _____________ along with
__________, which have no charge. The total number of these nucleons in an atom is the
__________________ of the atom. The electrons are the _____________ charged particles
found _________________________ . Electrons can be gained by the atom to form
______________ or lost to form ______________ .
12) The thing that makes any element on the periodic table different from all other elements is the
number of _______________ in its nucleus. ______________ are atoms of the same type of
element which have different _____________ because they have different numbers of
_____________ .
13) Write nuclear symbols for the following isotopes: (a) 19 protons, 20 neutrons, 18 electrons;
(b) 26 protons, 31 neutrons, 26 electrons; (c) 16 protons, 17 neutrons, 18 electrons.
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name: _____________________
date: __________
14) Fill in the missing information in the chart below.
chemical
symbol
atomic
number
K+
mass
number
# of
protons
42
19
38
# of
electrons
# of
net charge
neutrons
1
38
–
F
P
50
11
15
31
20
15) Given this nuclear symbol:
–1
0
18
22
238
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U , write the hyphen notation & give the number of neutrons.
16) Although mercury commonly exists in 7 isotopes, an asteroid with only two specific isotopes
is found: Hg-199 with a weight of 198.967g and Hg-200 with a weight of 199.968g. You
find that 62.305% of the mercury is Hg-200 & 37.695% is Hg-199.
(1) What is the average atomic mass of mercury on this asteroid?
(2) Why do isotopes of the same element have different masses?
17) State the names of (a) and (b), and tell the difference between the two. How is (a) formed?
a.
b.
18) According to the quantum theory, first postulated by ________________ , when an object
(such as an electron) loses energy, it does not do so continuously. Instead, it radiates off
energy in small specific amounts called _____________. Such particles of light or other
radiation are called ______________. Radiation is absorbed and emitted only in whole
numbers of these particles. According to the ________ model of the hydrogen atom, when
the electron in a hydrogen atom relaxes from the 3rd energy level down to the ____ energy
level, it emits photons of visible light. When it relaxes from the 3rd to the 1st energy level, it
gives off photons of ____________________ .
19) The dots in the Lewis diagram represent the ___________ in the ___________ energy level.
20) What are the 4 quantum numbers of electrons and what does each one represent or indicate?
21) Orbital filling order: 1s2, 2s2, 2p6, ___, ___, 4s2, ___, ___, ___, ___, ___, 6s2
22) Describe the shapes of s, p, and d orbitals (one word will suffice).
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name: _____________________
date: __________
23) Give the orbital notation, electron configuration notation, and shorthand e.c.n. for:
a. O b. Cr c. Ar d. Pt
24) Circle the valence electrons in the orbital notation of all elements in the previous question.
25) Give the electron dot diagrams (Lewis dot diagrams) and state the # of valence electrons for:
a. P b. K c. Al3+ d. Mn e. Br –
26) Name 3 elements whose orbitals are half-filled.
27) Write the decay equation and state what type of radiation is emitted when uranium-238
decays into thorium-234. Write the decay equation and state what particle of radiation is
emitted when protactinium-234 decays into uranium-234.
28) a. Describe the three types of radiation (their names, charges, and what type of particle it is),
& their penetrating power (what it would take to stop them).
b. Label X, Y and Z with the appropriate type of radiation:
Y
X
Z
29) a. What does it mean to say that the half-life of carbon–14 is 5730 years?
b. What makes carbon–14 (or any other radioactive isotope of any element) radioactive?
30) What are three practical applications of radiation sources?
31) How is heat produced by the radioactive decay of uranium-235 used to generate energy?
32) Radioactive isotopes of _____________ might be used to test thyroid function. Radiation
with enough energy to knock electrons off of atoms, giving them a charge is known as
___________________________ .
33) Name 3 metals, 3 non-metals, and 3 metalloids.
34) Define the following:
a. group (family) b. period c. atomic radius
f. ionization energy
d. ionic radius e. electronegativity
35) All elements (main group elements: 1A–8A) in the same group have: (1) the same number of
_________________, (2) the same ___________________, and similar ________________.
36) State the group or period in which you would find these elements, and what charge they
would have as ions.
a. halogens b. alkali metals c. noble gases d. alkaline earth metals e. rare earth metals
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name: _____________________
date: __________
37) Describe properties you would expect of the Group 1A metal sodium. How would you
expect the properties of Group 2A metal calcium to differ from those of sodium?
38) a. What is the most reactive group of metals?
b. What is the most reactive group of non-metals? The least reactive?
39) Describe the trends for each of the following, both within a group and within a period:
a. ionic radius b. electronegativity c. ionization energy d. atomic radius
40) Explain the trend for atomic radius (again, within both a group and a period) in terms of
electrons, the pull of the nucleus, & shielding.
41) Rank the following sets of elements in order of increasing (1) atomic radius; (2) ionic
radius; (3) ionization energy; (4) electronegativity:
a. Cs, Pb, Ba
b. P, Sb, N
c. Cl, S, P
42) Define the following and give one example (words or pictures) of each:
a. ionic bond b. covalent bond c. polyatomic ion d. molecule
e. polar covalent bond f. non-polar covalent bond g. dipole-dipole forces
h. hydrogen bonding i. London dispersion (van der Waal’s) forces
43) Excluding metallic bonds, there are two basic types of bonds: __________ and __________.
In ionic bonds, the electrons are transferred from a __________ to a ________________. In
covalent bonds, electrons are __________. Two non-metals join together by ___________
bonding. A metal and a non-metal join together by __________ bonding. Elements form
ions which have the same valence electron configuration as ___________ (because this is the
most ____________ arrangement of electrons).
44) a. Compare and contrast properties of ionic & covalent bonds (including melting and boiling
points, volatility, hardness, electrolytic nature, etc.). You may create a chart if that helps.
b. You have two white solids powders in front of you – a covalent solid and an ionic solid.
What tests could you do to determine which one was ionic and which one was covalent?
(Think in terms of solubility, solution conductivity, melting point, etc.)
c. You have several crystalline solids in front of you that are different colors (not white) –
are they more likely to be ionic or covalent?
45) a. What are the only 2 elements that can triple bond?
b. What are the 7 elements which commonly exist as diatomic molecules in nature?
46) What is the relationship between bond length and bond energy? Which of the following
bonds would require the most energy to break: F – F, Cl – Cl , or I – I ?
47) For each of the following molecules: (1) Draw electron (Lewis) dot diagrams; (2) Determine
their shape (VSEPR geometry), and (3) State whether the molecule is polar or non-polar.
a. HF b. CCl2F2 c. NI3 d. H2S e. BI3
48) What intermolecular forces would exist in molecule (a) in the previous problem? What
intermolecular forces exist in molecule (e)?
49) a. State six properties of metals.
b. How does the nature of metallic bonding help explain these properties?
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name: _____________________
date: __________
50) Would CCl4 be soluble in water? Why or why not? What types of compounds are almost
always soluble in water?
51) a. List two examples of natural polymers and two examples of synthetic polymers.
b. List five examples of network covalent solids, including the three allotropes of carbon.
c. What is the difference between macromolecules and network covalent solids?
52) Identify (name) the three allotropes of carbon illustrated below:
______________
_______________
_____________
53) Name the following compounds:
a. SF6 b. Ba(OH)2 c. K2CO3 d. P2O5 e. NH4C2H3O2 f. Na2O
h. H2SO4 i. HCl j. Cu3PO4 k. Pb(ClO3)4
g. FeS
54) Consider all the compounds listed in the previous problem:
a. Which of the compounds in the previous question are covalent?
b. Which compounds are ionic?
c. Which compounds are ionic, but also contain covalent bonds?
55) Write chemical formulas for each of the following:
a. phosphorus trichloride b. magnesium phosphate c. cobalt (II) nitrate
d. dinitrogen trioxide e. ammonium perchlorate f. copper g. sulfite
h. silver chromate i. phosphoric acid j. peroxide k. acetic acid
56) What is the percent composition of Pb3(PO4)2?
57) An oxide of vanadium is 56.00% vanadium and
44.00% oxygen. What is its empirical formula?
If the molecular mass of the true compound is
364g, what is its molecular formula?
58) Use the solubility chart to the right.
(1) At what temperature will 110g NaClO3
dissolved in 100g of water crystallize?
(2) How many grams of KNO3 can be dissolved
in 100g of water at 50C?
(3) At what temperature will 50g of NH4Cl
dissolved in 100g of water crystallize?
(4) Which compound is least soluble at 45C?
(5) Would a solution containing 90g of
Pb(NO3)2 be saturated at 30C?
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name: _____________________
date: __________
59) You want to dissolve as much carbon dioxide as possible into a 2 liter bottle of Mountain
Dew. What conditions of temperature (high or low) and pressure (high or low) should you
use to dissolve the maximum amount of CO2? Would your answers be the same if you were
trying to dissolve as much sugar as possible in the Mountain Dew? Why or why not?
60) Write & balance each formula equation, and identify the reaction type (synthesis, etc.):
a. barium nitrate + iron (II) sulfide react to yield barium sulfide plus iron (II) nitrate
b. zinc plus lead (III) sulfate yield zinc sulfate and lead
c. magnesium hydroxide (when heated) yields magnesium oxide and water
d. aluminum and oxygen produce aluminum oxide
e. potassium phosphate plus iron (II) chloride yield potassium chloride and iron (II) phosphate
61) Assign state symbols [(s) or (aq) only] to the reactants and products in (a) & (e) from the
previous problem, using solubility tables.
62) Write ionic and net ionic equations, and identify any spectator ions for the balanced equation
in (e) in the previous problem.
63) Using the activity series, predict whether or not the following reactions will occur. If they
will occur, predict the products of the reaction (but you do not need to balance the equation).
a. Zn(s) + HCl(aq) 
b. Mg(s) + NaBr(aq) 
c. Cl2(s) + AlI3(aq) 
64) What is the mass of 23.88 moles of Li?
65) How many molecules in 0.157 moles of N2O5?
66) How many grams are in 2.60 x 1025 formula units of Cu3(PO4)2?
67) Given 1.85 moles of Cu3(PO4)2:
a. How many formula units of Cu3(PO4)2 does this represent?
b. How many moles of phosphate ions are present in 1.85 moles of Cu3(PO4)2?
c. How many moles of copper atoms in 1.85 moles of Cu3(PO4)2?
d. How many phosphorus atoms in 1.85 moles of Cu3(PO4)2?
68) Using the following reaction: C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g), what mass of water
vapor can be formed from the combustion of 275g of C3H8?
69) Considering the balanced equation from the previous problem:
a. What volume of carbon dioxide will be produced from the combustion of 3.75 L of
C3H8(g) with sufficient oxygen at the same conditions of temperature and pressure?
b. How many grams of CO2 will be formed by the reaction of 6.00 L of C3H8(g) with
sufficient oxygen at STP?
70) Given the following reaction: Zn + 2HNO3  Zn(NO3)2 + H2
a. If 52.0g of zinc react with 27.0g of nitric acid, what mass of zinc nitrate is formed?
b. What is the limiting reactant?
c. How many grams of excess reactant remain?
d. Will this reaction actually occur? Why or why not (how do you know)?
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name: _____________________
date: __________
71) Methanol can be produced by reacting carbon monoxide with water. If 75.0g CO reacts to
produce 68.4g methanol, what is the percent yield of methanol?
(CO + 2H2O  CH3OH + O2)
72) What is the concentration (molarity) of 0.375 moles of NaCl dissolved in 0.500 L of water?
73) What is the molarity of a solution made by dissolving 22.15g of (NH4)2SO4 in enough water
to make 50.00mL of solution?
74) How many milliliters of solution are necessary to make a 1.25 M solution from 58.3g of
acetic acid (ethanoic acid)?
75) You need to dilute 10.0 mL of a 2.00 M NaOH solution down to a 0.500 M solution. What
will be your final solution volume?
76) How much water must you add in order to dilute 15.0 mL of 6.00 M HCl down to a 0.100 M
solution?
77) What are the 5 assumptions of the Kinetic-Molecular Theory of gases?
78) Consider the following gases: Ne, CO2, HCl.
a. Which would you expect to behave most like an ideal gas and why?
b. Under what conditions do real gases behave most like ideal gases?
79) If the pressure exerted on 245mL sample of gas at constant temperature is increased from
325mm Hg to 550mm Hg, what is the new volume of the gas?
80) A sample of gas has a volume of 140.0mL at 67C. To what temperature must the gas be
lowered to reduce its volume to 50.0 mL at constant pressure?
81) A sample of gas at 47C and 780.mm Hg pressure occupies a volume of 2.20L. What volume
would this gas occupy at 107C and 600.mm Hg pressure?
82) Calculate the pressure (in atm) of 38.5g of CO2 taking up 750.mL of volume at 57C.
83) How many grams of oxygen (O2) would be needed to fill up a deflated basketball with a
volume of 525 mL at 25C and 133 kPa?
84) Given this balanced equation: 3CH4(g) + 4Al(OH)3(s)  Al4C3(s) + 12H2O(l)
If 14.0L of CH4 at 725mm Hg and 37C react, what mass of aluminum carbide is produced?
85) A sample of gas containing O2, CO2, and N2 is at 766 mm Hg. The PCO2 = 0.285 atm, and the
PN2 = 36.6 kPa. What is the partial pressure of O2 (in mm Hg)?
86) A sample of chlorine gas is collected by water displacement at 20.0C. If barometric pressure
at the time of the experiment is 752 mm Hg, what is the partial pressure of the Cl2 collected?
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name: _____________________
date: __________
87) (1) Tell whether the following reactions are endothermic or exothermic and then give the H
for each. (2) Describe if the reaction will feel hot or cold. (3) Do these reactions represent an
increase or decrease in entropy?
a. CaCO3(s) + 176kJ  CaO(s) + CO2(g)
b. SO3(g) + H2O(l)  H2SO4(aq) + 129.6 kJ
88) a. Draw your own potential energy diagram for an exothermic reaction. Label both axes and
assign potential energy values (in kJ) to the reactants, activated complex, and products.
b. Label H, EA and EA’ on your PE diagram.
c. Determine H for your reaction (forward and reverse), and determine the forward and
reverse activation energies.
d. Which one (reactants, activated complex, products) is least stable? Which one is most
stable? How do you know? All that energy given off in the course of this exothermic
reaction – where does it come from?
e. Draw how the PE diagram would look different if a catalyst was used in the reaction.
89) List the 6 factors that affect the rate of reaction and specifically how they affect reaction rate.
90) Everything in nature tends toward lower ________________ & higher _______________.
91) CaO(s) + H2O(g) + 38.5 kJ  Ca(OH)2(s)
Will this reaction be spontaneous at 25C, if it involves an entropy change of 0.150 kJ/molK?
92) The __________________________ of an open system is constantly _________________
due to _________________ .
93) A calorimeter is filled with 125 g of water at an initial temperature of 20.7C. A 2.43 g
sample of solid KOH is added, and the temperature increases to 24.1C.
a. Determine the quantity of heat absorbed by the water for this solution process.
b. Determine the molar enthalpy of solution for KOH.
94) If 474.8 joules of heat are needed to raise the temperature of 35.0 g of xylene from 10.4C to
20.6C, what is the specific heat of xylene (be careful to get the units correct)?
95) What is the enthalpy change associated with taking 53.8 g of ice (at 0C), warming it to the
boiling point of water, and then boiling all of it off?
96) Draw a phase diagram and a heating curve (freezing & boiling point graph) for substance “X”
whose melting point is 11C and boiling point is 78C. (For the phase diagram, you must
also include the triple point (at –5C & 0.5 atm) and critical point (at 110C & 20 atm);
assume that an increase in pressure causes substance “X” to freeze). Refer to notes for help.
97) Explain, in terms of equilibrium and vapor pressure, what happens when you take a sealed
beaker, filled with water in equilibrium with its vapor at room temperature, and place it in a
warm, sunny window.
98) a. List the six strong acids & vinegar, giving both their names and formulas.
b. List seven strong bases (formulas only).
99) Write an equation in which HCl functions as a Bronsted acid, but not an Arrhenius acid.
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name: _____________________
date: __________
100)
Consider the following solutions:
Solution A: 3.8  10–3 M HCl
a.
b.
c.
d.
e.
Solution B: 4.75  10–5 M NaOH
Calculate the pH, pOH, [H3O+], and [OH–] for both of these solutions.
What color would solution A become if several drops of litmus indicator were added.
What color would solution B become if several drops of phenolphthalein were added.
What gas would be produced if Na2CO3(s) or NaHCO3(s) were added to Solution A.
If the gas from (d) were bubbled through lime water, what color would it turn?
101)
In the following reaction, identify the acid, base, conjugate base & conjugate acid, and
circle the stronger acid and base:
H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4– (aq)
102)
Why are strong acids & bases also strong electrolytes?
103)
Write and balance the neutralization reaction between NaOH and H2SO4.
104)
If 15.0mL of 0.0250 M aqueous HNO3 is required to neutralize 10.0mL of Ca(OH)2 to
the endpoint of the titration, what is the molarity of the Ca(OH)2?
105)
Give oxidation numbers for all atoms in:
a. LiCl
106)
a.
b. CH4
c. KMnO4
d. MgCr2O7
e. HClO2
f. CuO
g. IO3–
Balance redox reactions for these simple metal-ion reactions and state whether oxidation
or reduction is occurring:
Mn+2  Mn+7
b.
Fe+2  Fe
c.
S2–  S
107)
Write the overall ionic and net ionic equations for the following reaction:
Cu(s) + HNO3(aq)  Cu(NO3)2(aq) + NO(g) + H2O(l)
108)
a.
b.
c.
Using the previous reaction:
Assign oxidation numbers to all elements.
Identify the oxidizing and reducing agents.
Balance the oxidation and reduction half reactions, and add them together to get a
balanced oxidation-reduction reaction.
109)
Draw the electrochemical cell described by the following reaction:
2Fe(s) + 3Cu(NO3)2(aq)  2Fe(NO3)3(aq) + 3Cu(s)
a.
b.
c.
d.
e.
110)
Using the activity series, label the anode & cathode.
Show direction of electron flow & cation flow.
Tell what is happening to the mass of the copper electrode.
Tell what is happening to the concentration of iron ions in solution.
Calculate the voltage for this cell.
State three practical applications of redox reactions and briefly describe how they work.
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