Chemistry EOC Review – Answer Key
1) a. 68,200,000 cg
b. 5000 mm
c. 0.0125 kL
d. 0.548 dm
2) ~21.5 mL
3) production of heat & light, formation of a gas, formation of a precipitate, color change
4) 33.64 g / 4.28 ml = 7.86 g/ml = Fe
5) m.p. –183C = CH4 (methane)
6) H2 – splint test (burning splint ignites hydrogen)
O2 – splint test (oxygen reignites glowing splint)
H2O – cobalt chloride paper (paper changes from blue to pink in the present of water vapor)
CO2 – splint test (carbon dioxide extinguishes burning splint);
limewater test (limewater turns cloudy when carbon dioxide is bubbled through it,
producing CaCO3)
7) The protons are positively charged particles found inside the nucleus along with the neutrons, which
have no charge. The total number of these nucleons in an atom is the mass # / atomic mass of the
atom. The electrons are the negatively charged particles found orbiting in “clouds” around the
nucleus. Electrons can be gained by the atom to form anions or lost to form cations.
8) The thing that makes any element on the periodic table different from all other elements is the number
of protons in its nucleus. Isotopes are atoms of the same type of element which have different
masses because they have different numbers of neutrons .
9) a.
b.
c.
39

19 K
57
26 Fe
33 2 
16 S
10)
chemical
symbol
K+
Sr
FP
Ca
2+
11) uranium – 238
atomic
number
19
38
mass
number
42
# of
protons
19
88
38
# of
electrons
18
38
9
15
20
31
9
10
15
20
42
20
15
18
# of
net charge
neutrons
+1
23
50
0
11
-1
16
22
0
+2
238 – 92 = 146 neutrons
12) (1) 199.59g
(2) They contain the same # of protons (they are the same element), but different #s of
neutrons (electrons don’t matter, b/c they have virtually no mass), & thus different masses.
1
13) a. line spectrum – formed when atoms of an element are heated and their excited electrons
relax back down to lower energy levels (line spectra are unique to each element)
b. continuous spectrum
14) According to the quantum theory, first postulated by Max Planck, 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 quanta. Such particles of light or other radiation are called photons. Radiation is
absorbed and emitted only in whole numbers of these particles. According to the Bohr model of the
hydrogen atom, when the electron in a hydrogen atom relaxes from the 3rd energy level down to the
2nd energy level, it emits photons of visible light. When it relaxes from the 3rd to the 1st energy level,
it gives off photons of UV light.
15) The dots in the Lewis diagram represent the valence electrons in the highest (outermost) energy
level.
16) POMS
Principal quantum # - the energy levels (shells) surrounding the nucleus (1-7)
Orbital (Azimuthal) quantum # - indicates orbital shape (subshells = s, p, d, f)
Magnetic quantum # - indicates orientation of an orbital around the nucleus (x, y, z axes)
Spin quantum # - indicates the two possible states of an e– in an orbital (+½ , -½ , or )
17) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2
18) s – sphere
p – peanut
d – double peanut
19) a. O: 1s2 2s2 2p4 [He]2s2 2p4
b. Cr: 1s2 2s2 2p6 3s2 3p6 4s2 3d4
[Ar]4s2 3d4
c. Ar: 1s2 2s2 2p6 3s2 3p6 [Ne]3s2 3p6
d. Pt: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d8
[Xe]6s2 4f14 5d8
20) (valence electrons highlighted in red, above)
21) a. P
(5) b. K
(1) c. Al
(8) d. Mn
(2) e. Br
(8)
22) N, Mn, P, K, etc. (anything in group 2, 7, 15)
23)
238
92 U
234
91 Pa


234
90Th
234
92 U
+
+
4
2+
2 He
0 –
1 e
24)
Name
identity
Symbol
charge
penetrating ability
2
Alpha ()
helium-4 nuclei
Beta ()
electrons
Gamma ()
high energy
non-particle radiation
4
2
4
2
He 2+
He
4
2

0
1
e-
0
1
e

0
0

0
1
2+
low: stopped by paper,
clothing, several cm of air
1–
medium: stopped by
heavy clothing, metal
none
high: stopped by lead,
concrete
25) To say that the half-life of carbon–14 is 5730 years means that half of the carbon-14 atoms present in
a sample of matter will decay 5730 years.
26) Four practical applications of radiation include:
1. electricity
2. radioactive carbon-14 dating
3. radioisotopes used in medicine (radiotracers, ionizing radiation, gamma sterilization)
27) Heat from the fission reaction boils water, and the steam turns a turbine, which generates electricity.
28) Radioactive isotopes of iodine might be used to test thyroid function. Radiation with enough energy
to knock electrons off of atoms, giving them a charge is known as ionizing radiation.
29) (answers will vary…)
30) a. group (family) – columns going down – all elements in the same group have the same
outer electron configuration / valence e- (& therefore, similar properties)
b. period – rows across periodic table set up to reflect regularly recurring properties of elements
c. atomic radius – ½ the distance between the nuclei of identical atoms joined in a molecule
d. ionic radius – ½ the diameter of an ion in an ionic compound
e. electronegativity – ability of an atom in a chemical compound to attract electrons (F = 4.0)
f. ionization energy – energy required to strip away one electron from an atom
31) All elements (main group elements: 1A–8A) in the same group have: (1) the same number of valence
electrons, (2) the same oxidation number, and (3) similar chemical properties.
32) a. halogens – group 17 (–1)
b. alkali metals – group 1 (+1)
c. noble gases – group 18 (0)
d. alkaline earth metals – group 2 (+2)
e. rare earth metals – (lanthanides) top row of the F-block (charges vary, mainly +2)
33) Na – soft, silvery metal; easily cut with a knife; very reactive with air & water
Ca – harder, higher melting point, slightly less reactive than Na
34) a. Group 1 (1A)
b. Group 17 (7A); Group 18 (8A)
35) a. ionic radius – increases as you go left and down
b. electronegativity – increases across (right) and up
c. ionization energy – increases across (right) and up
d. atomic radius – increases go left and down
3
36) Atomic radii increase as you go down a group b/c the outer electrons occupy higher # (and energy)
sublevels, farther from the nucleus. More energy levels also means more shielding from inner shell esuch that the nucleus cannot hold onto outer shell electrons as tightly.
Atomic radii decrease across a period because you are adding a proton and an electron each time you
go to the right, but the electrons are in the same energy level, thus there is no additional shielding
from inner shell e-. More protons will therefore mean a stronger pull on the electrons in the same
energy level, pulling them in tighter.
37)
A.R.
a. Pb, Ba, Cs
b. N, P, Sb
c. Cl, S, P
I.R.
Pb, Ba, Cs
N, P, Sb
Cl, S, P
I.E.
Cs, Ba, Pb
Sb, P, N
P, S, Cl
E-neg
Cs, Ba, Pb
Sb, P, N
P, S, Cl
38) a. ionic bond – metal bonded with a non-metal or polyatomic ion; the metal loses an e-, the
non-metal gains the e- & the 2 ions are attracted to each other (ex: NaCl  Na+ Cl–)
b. covalent bond – 2 atoms (non-metals) share electrons (ex: NH3, H20, etc.)
c. polyatomic ion – a charged group of covalently bonded atoms (ex: NH4+, SO42-)
d. molecule – a group of 2 or more atoms held together by covalent bonds (ex: CO2, H2O)
e. polar covalent bond – difference in electronegativity between bonded atoms is 0.3 – 1.7
(ex: C – H)
f. non-polar covalent bond – difference in electronegativity between atoms is 0 – 0.3
(ex: 2 identical atoms bonded together: H – H ( = H2 ))
g. dipole-dipole forces – forces of attraction between polar molecules
(ex: solution of I – Br )
+ h. hydrogen bonding – attraction between a hydrogen bonded to a strongly electronegative
atom (N, O, F) & an unshared pair of electrons on another strongly
electronegative atom (ex: H2O)
i. London Dispersion Forces – result from all the electrons being on one side of the atom 
(van der Waals Forces)
instantaneous & induced dipoles (ex: only intermolecular
force which occurs in all atoms & molecules)
39) Excluding metallic bonds, there are two basic types of bonds: covalent and ionic. In ionic bonds, the
electrons are transferred from a metal to a non-metal. In covalent bonds, electrons are shared. Two
non-metals join together by covalent bonding. A metal and a non-metal join together by ionic
bonding. Elements form ions which the same valence electron configuration as noble gases (because
this is the most stable arrangement of electrons).
40) (1)
ionic
metal gains, non-metal loses
electrons
high
melting/boiling points
stronger
bond strength
non-volatile
volatility
brittle, hard, crystalline
hardness
molecular (covalent)
shared
low
weaker
often, but varies
varies
4
electrolytic
conductivity
usually non-electrolytes
(also remember covalent = 2 non-metals; ionic = metal + non-metal/polyatomic ion)
(2) Covalent compound would be less likely to dissolve in water and even if it did, would probably
not conduct electricity in solution (might dissolve in non-polar solvent). Ionic compounds usually
dissolve in water, and if they do, they will conduct electricity. Covalent compound also likely has
lower melting and/or boiling point than ionic compound.
(3) Ionic compounds more likely to have colored crystals.
41) a. (C:::C CC, C:::N CN)
b. H2, N2, O2, F2, Cl2 , Br2, I2 (BrINClHOF)
42) The shorter the bond, the higher the bond enthalpy (energy). F – F.
43) a.
b.
linear
tetrahedral
(polar)
(polar)
c.
triangular
pyramidal
(polar)
44) a. H-bonding (dipole-dipole), van der Waal’s;
d.
e.
bent
(polar)
triangular
planar
(non-polar)
e. van der Waal’s
45) Six properties of metals: high electrical conductivity, malleability, thermal conductivity, ductility,
high melting & boiling point, and luster.
If the delocalized electrons are free to move around, that allows heat and an electric current to flow
easily. Also, because of the delocalized electrons, metallic bonding is not directional, but uniform
throughout the solid. One plane of metal ions can slide past one another without encountering any
resistance or breaking any bonds.
46) No. It is non-polar. Polar compounds (most ionic compounds).
47) a. DNA, proteins / PVC, nylon
b. diamonds, graphite, buckminsterfullerenes, silicon, silicon dioxide
c. Polymers are essentially composed of a relatively small unit (a monomer) that repeats over
and over again. Network covalent solids are similar to macromolecules but they contain no
discrete molecular units (like a monomer). Rather they form a continuous network of
covalently bonded atoms.
48)
diamonds
graphite
buckminsterfullerenes
49) a. sulfur hexafluoride b. barium hydroxide c. potassium carbonate
d. diphosphorus pentoxide e. ammonium acetate f. sodium oxide g. iron (II) sulfide
h. sulfuric acid i. hydrochloric acid j. copper (I) phosphate k. lead (IV) chlorate
5
50) a. a, d, h, i
b. b, c, e, f, g, j, k
c. b, c, e, j, k
51) a. PCl3 b. Mg3(PO4)2
h. Ag2CrO4 i. H3PO4
c. Co(NO3)2 d. N2O3
j. O22– k. HC2H3O2
52) Pb: 76.6%
O: 15.8%
53) E.F. = V2O5
P: 7.6%
e. NH4ClO4
f. Cu
g. SO32–
M.F. = V4O10
54) According to this solubility chart:
(1) ~35C
(2) ~80 g
(3) ~50C
(4) KClO3
(5) yes
55) High P / Low T
No – solids are more soluble at high T; pressure doesn’t affect solubility of solids
56) a. Ba(NO3)2 + FeS  BaS + Fe(NO3)2 (double replacement)
b. 3Zn + Pb2(SO4)3  3ZnSO4 + 2Pb (single replacement)
c. Mg(OH)2  MgO + H2O (decomposition)
d. 4Al + 3O2  2Al2O3 (synthesis)
e. 2K3PO4(aq) + 3FeCl2(aq)  6KCl(aq) + Fe3(PO4)2(s) (double replacement)
57) a. Ba(NO3)2(aq) + FeS(s)  BaS(aq) + Fe(NO3)2(aq)
e. 2K3PO4(aq) + 3FeCl2(aq)  6KCl(aq) + Fe3(PO4)2(s)
58) ionic equation:
6K+(aq) + 6Cl–(aq) + Fe3(PO3)2(s)  6K+(aq) + 2PO43–(aq) + 3Fe2+(aq) + 6Cl–(aq)
net ionic equation:
Fe3(PO3)2(s)  2PO43–(aq) + 3Fe2+(aq)
+
–
spectator ions:
K (aq) & Cl (aq)
59) a. ZnCl2 + H2
b. NR
c. AlCl3 + I2
60) 165.8 g Li
61) 9.45  1022 molecules N2O5
62) 16,400 g Cu3(PO4)2 (16.4 kg)
63) a.
b.
c.
d.
1.11  1024 formula units of Cu3(PO4)2
3.70 mol of PO43– ions
5.55 mol of Cu atoms
2.23  1024 atoms of Cu
64) 449 g H2O(g)
6
65) a. 11.3 L CO2(g)
b. 35.4 g CO2(g)
66) a.
b.
c.
d.
40.5 g Zn(NO3)2
(HNO3)
38.0 g (of Zn)
Yes. Zn is above H2 on the activity series.
67) TY = 48.9g C2H5Cl / PY = 48.5%
68) 0.750 M NaCl
69) 3.35 M (NH4)2SO4
70) 0.777 L (777 mL)
71) 0.0400 L (40.0 mL)
72) 885 mL
73) 5 assumptions of the Kinetic-Molecular Theory (KMT) of gases:
1. Gases have negligible volume.
(distances separating gases are relatively great; volume of gas particles themselves is )
2. Gas particles are in constant motion, moving rapidly in straight lines in all directions. Thus they
have kinetic energy.
3. Collisions between gas particles or between gas particles and container walls are elastic  no KE
lost
4. There are no forces of attraction or repulsion between gas particles.
5. The average KE of a gas is directly proportional to the temperature (K) of the gas.
74) a. Ne – noble gas, low mass, very little attraction for neighboring Ne atoms
b. high T, low P
75) 145 mL
76) 121 K (–152C)
77) 3.40 L
78) 31.6 atm
79) 0.902 g O2
80) 25.2 g Al4C3
81) 275 mm Hg
82) 735 mm Hg
83) a. (1) endothermic / H = +176kJ; (2) cold; (3) increase in entropy
7
b. (1) exothermic / H = –129.6kJ; (2) hot; (3) decrease in entropy
84) (answers will vary…)
85) Factors affecting reaction rate:
1) The Reactants (the nature of the reactants)
 By “nature of the reactants”, we can mean things like the number and type of bonds that must
be broken and reformed in the course of a reaction (and the general structural complexity of
the reactant and product molecules).
2) Surface Area
 More surface area = greater # of collisions possible = faster reaction.
3) Concentration
 Increasing the concentration of a reactant usually increases the rate of the reaction.
4) Pressure
 Pressure is only a factor on reaction rate for reactions involving gases. Increasing pressure on
a reaction where reactants are gases will increase the rate of the reaction.
5) Temperature
 Increase in T  increase in KE  increase in reaction rate.
 Why? Because more KE  (1) greater frequency of collisions & (2) more particles with
sufficient activation energy to react when collisions occur.
6) Presence of Catalysts
 Most catalysts speed up reactions by lowering the activation energy required for a reaction
(often by splitting the reaction up into multiple steps, which have lower activation energies.
This speeds up a reaction because when reactant molecules collide at any given temperature,
they are more likely to have sufficient activation energy to react (since less is required).
86) Everything in nature tends toward lower enthalpy & higher entropy.
87) G = –6.20 kJ/mol; spontaneous
88) The total useful energy of an open system is constantly declining due to entropy.
89) a. 1.8 kJ
b. –41.6 kJ/mol (of KOH)
90) Cxylene = 1.33 J/g·C (1.33 J g–1 C–1)
91) 162 kJ
92) (see answers…)
93) If the temperature in this system increases (you warm up the beaker), the vapor pressure will increase
for two reasons:
 More evaporation  more vapor molecules running into container walls (i.e. – the water vapor
molecules are more concentrated).
 More heat = more KE available to vapor molecules = particles colliding with container walls
more frequently and with more force.
 However, assuming all the liquid water does not evaporate, the system will eventually reestablish
equilibrium between evaporation and condensation, as the warmer air will eventually become
8
saturated with water vapor again. Until equilibrium is reach, we say that the forward reaction
(evaporation, here) is favored.
94) a. HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid),
HNO3 (nitric acid), H2SO4 (sulfuric acid), HClO4 (perchloric acid),
HC2H3O2 or CH3COOH (acetic acid)
b. LiOH, NaOH, KOH, RbOH, CsOH, FrOH, Ba(OH)2
95) HCl(aq) + NH3(aq)  NH4+(aq) + Cl–(aq)
96)
a.
b.
c.
d.
e.
Solution A: 3.8  10–3 M HCl Solution B: 4.75  10–5 M NaOH
Solution A: pH = 2.42; pOH = 11.58; [H3O+] = 3.8  10–3 M; [OH–] = 2.6  10–12 M
Solution B: pH = 9.68; pOH = 4.32; [H3O+] = 2.11  10–10 M; [OH–] = 4.75  10–5 M
red
pink
CO2
milky white
97) H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4–(aq)
acid
base
conjugate conjugate
acid
base
98) Strong acids and bases are also strong electrolytes because they ionize or dissociate 100%
in aqueous solution; the more ions, the better solution is able to conduct electricity.
99) 2NaOH(aq) + H2SO4(aq)  Na2SO4(aq) + 2H2O(l)
100)
101)
0.0188 M Ca(OH)2
Give oxidation numbers for all atoms in:
+1 –1
–4 +1
+1 +7 –2
+2 +6 –2
a. LiCl
b. CH4
c. KMnO4
d. MgCr2O7
+1 +3 –2
+2 –2
+5 –2
e. HClO2
f. CuO
g. IO3–
102)
a. Mn+2  Mn+7 + 5e–
oxidation
103)
3Cu(s) + 8HNO3(aq)  3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)
3Cu(s) + 8H+(aq) + 8NO3–(aq)  3Cu+2(aq) + 6NO3–(aq) + 2NO(g) + 4H2O(l)
3Cu(s) + 8H+(aq) + 2NO3–(aq)  3Cu+2(aq) + 2NO(g) + 4H2O(l)
104)
b. 2e– + Fe+2  Fe
reduction
c. S2–  S + 2e–
oxidation
0
+1 +5 –2
+2 +5 –2
+2 –2
+1 –2
3Cu(s) + 8HNO3(aq)  3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)
0
+1
+5 –2
+2
+5 –2
+2 –2
+1 –2
steps 1 & 2: 3Cu(s) + 8H+(aq) + 8NO3–(aq)  3Cu+2(aq) + 6NO3–(aq) + 2NO(g) + 4H2O(l)
steps 3 & 4: +5
+2
0
+2
9
reduction: NO3–  NO
steps 5 – 6:
reduction:
step 7:
step 8:
oxidation: Cu  Cu2+
NO3–  NO
oxidation: Cu  Cu2+
3e– + 4H+ + NO3–  NO + 2H2O
Cu  Cu2+ + 2e–
2  (3e– + 4H+ + NO3–  NO + 2H2O)
3  (Cu  Cu2+ + 2e–)
6e– + 8H+ + 2NO3–  2NO + 4H2O
+
3Cu  3Cu2+ + 6e–_________
3Cu + 8H+ + 2NO3–  2NO + 4H2O + 3Cu2+
10