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Chemistry First Semester Final Exam Review
Chapter 1 & 2
1)
2)
What is the difference between chemical and physical properties?
Physical properties can change w/o changing the actual substance. Chemical
properties refer to the chemical tendencies of one specific substance
Write C for chemical property or P for physical property.
__P___ Copper sulfate is blue.
__C___ Silver nitrate reacts with sodium chloride.
__P___ Oxygen is a gas.
___C__ Argon is an inert element.
__P___ The density is 3.2 g/ml.
__P___ It is a liquid at room temp.
3) Write C for chemical change or P for physical change.
__P___evaporation/condensation
__P___crushing
__C___rusting
__P___subliming
__C___burning/smoke
__C___souring (as in milk)
__P___ dissolving a solid in a liquid
4) What are the indications that a chemical change (reaction) has occurred?
Light, heat, color, gas, precipitate
5) a. Define a chemical change: when one substance is converted into an entirely new
substance
b. How do properties of compounds compare to the properties of the elements they are
made up of?
Their properties are entirely different
6)
What is an alloy? Do the components keep the properties they had before they were
mixed? Are there chemical bonds?
A mixture of 2 elements (usually metals), they do keep some properties, no, the metal
atoms are just next to each other
7) Write HE if the following mixture is heterogeneous and HO if it is homogeneous:
_HE__beach sand
_HO__diet coke
_HO__jello
_HO__salt dissolved in water
8) Define Precipitate:
A solid that forms and settles from a solution
9) Write E for element, C for compound or M for mixture.
10)
_C____HgO
__C____CO2
__C___steam
__M____air
_M____alloys
__C____salt
__E___24 K gold
__M____vegetable soup
Describe the processes of distillation and filtration. What property is used in each to
actually separate substances?
Distillation: boiling point
filtration: size
11) WHAT IS.......
Heterogeneous mixture_Matter that has parts with different properties?
Homogeneous mixture__Matter that has identical properties throughout?
elements________Substances that can't be further decomposed by ordinary chemical
means?
compounds______Substances composed of 2 or more elements chemically combined?
Pure substance__Homogeneous material consisting of one particular kind of matter?
Chapter 3
12)
Convert the following .
a) 2500 km = __2.5 x 108_____________ cm
c) .3 m = ____30___________ cm
b) 300 ml = _300______________ cm3
d) 200kg = __2.0 x 108_______________ mg
13) Read the following scales to the correct decimal place!
7.0
7.5

_7.30 a)
25
_24.60_ b)
24
.4
.5

_0.43_ c)
14) In a lab experiment, the following results were obtained in measuring the mass of a
50.00 g cylinder of metal: 50.00g, 45.99 g, 55.98g, 49.01g, and 51.02g. Are these
results accurate, precise, both or neither? EXPLAIN.
Neither 55.98-45.99 = 9.99 varies more than just hundredths place – not precise
50.00-50.40 = 0.40 varies more than just hundredths place – not accurate
15) How many significant figures in each number? Underline the uncertain digit in
letter f.
a)52 __2____
e)500
__1_____
b)0.001 __1______
f)7.0 x 10-3 __2_______
c)300.0 __4______
d)30600__3______
g) 7.005 x 1012 __4________
16) Given that the density of an object is .57g/cm3 and that the mass is 40.0g, Calculate
the volume of the object? Round the answer to the correct number of significant
digits.
D = m/v 0.57 g/cm3 = 40.0 g/v
V = 70. cm3
17) Round the following numbers to 3 significant figures.
5374__5370_________
7645000_7650000_________ 0.4789___0.479_________
18) Write the following numbers is scientific notation.
0.000506 _5.06 x 10-4________
5070000 __5.07 x 106____________
32 __3.2 x 101_________
0.0307 _3.07 x 10-2_____________
19) Use dimensional analysis to answer the following question:
How many H are in 25 U if 2H = 3F, 5F = 1.2K, and 1K = 0.6U?
25U x 1 K x 5 F x 2 H = 250 = 115.7 = 120H
0.6 U 1.2 K 3 F
2.16
20) In a lab experiment, the mass of a 20.00 g weight was to be determined. The student
measured the mass to be 18.60g. Calculate the percent error. %E = (M-A) x 100/A
18.60 – 20.00 x 100 = 7.00 %
20.00
Chapter 4/25
21)
In the space before the question, write the part of the atom which gives the
characteristic listed. (Your choices are protons, neutrons, electron cloud, valence
electrons, or nucleus.)
_electron cloud__________________a) produces most of the atom's volume.
_protons__________________b) determines the identity of an atom.
_nucleus__________________c) is responsible for most of the atom's mass
_electrons__________________d) determines the chemical properties of an atom
_neutrons__________________e) acts as "nuclear glue" to hold the nucleus together
22)
A nitrogen atom has 7 protons and 7 neutrons. Determine a) the atomic number b)
the mass number.
a) 7
b) 14
23) a) How many protons, neutrons and electrons for the following elements? b) write the
name of each of these isotopes.
32
209
p+
a) 16 S
16
b)
82Pb
no
16
127
e–
16
82
82
24) Complete the following table:
particle
alpha
beta
gamma
isotopic symbol
42He
0-1e
00γ
mass (AMU)
4
0
0
Greek Sy.
α
β
γ
charge
+2
-1
0
25) Complete the following transmutation reactions:
a) Pu – 235 gives off a beta particle to become ___235Am________
b)
Am-245 gives off an alpha particle to become __241Np_________
26) The element chlorine is made up of 2 isotopes. One of these has a mass of 34.97, the
other has a mass of 36.97. The light isotope has an abundance of 75.5%, the heavy
isotope has an abundance of 24.5%. Calculate the average atomic mass of chlorine.
(34.97 x 0.755) + (36.97 x 0.245) = 35.46
27)
a) Why do isotopes have different mass numbers? The # of neutrons are different
b) Which of these two are isotopes of each other?
14
7N
15
7N
14
6
C
28) Complete the following half-life problems:
a) Given that the half-life of carbon-14 is 5730 years, consider a sample of fossilized
wood that, when alive, would have contained 24 g of carbon-14. It now contains
1.5 g of carbon-14. How old is the sample?
24  12  6  3  1.5
4 x 5730 = 22920 yrs
b) How much of a 100 g sample of Pu-239, T1/2 is 24,100 years, will remain after
96,400 years?
96400 = 4
100  50  25  12.5  6.25
6.25 g
24100
c)
Determine the half life from the graph.
600  1 hr
300  3 hr
600
grams
T1/2 = 2 hrs
300
0 1 2 3 4 5 6
Hours
Chapter5
29) Rank yellow light, IR, green light, and U.V. in order of decreasing energy.
UV, green, yellow, IR
30) The frequency of electromagnetic radiation is _α__ proportional to the energy. (Use
the Greek letter symbol to represent the proper proportion.
31) Which of the following are in the ground state and which are in an excited state?
Indicate with an E or G in the space.
_E___ 1s22s22p63s13p1
_E____ 1s22s22p43s2
__G___ [Ne]3s23p1
__G___ [Ne]3s23p6
32)
Electrons jump to higher energy levels when they (lose/gain) energy. When they fall
back to their ground state, energy is (gained/released). Because the difference in
energy levels is constant, the amount of energy is a fixed amount. This results in
photons of light which have a fixed _frequency__ and energy. These photons appear
on a screen as bands of light.
33)
Which electron jump would absorb the most energy?
a. 5  6
b. 6  7
c. 3  4
34)
d. 1  2
Which electron jump would release the most energy?
a. 3  2
b. 4  2
c. 5  2
d. 7  6
35)
Draw an s and a p-orbital. How many orientations are there for each in space?
s – 1 orientation
p – 3 orientations
36)
What does the shape of an orbital represent?
37)
38)
Where electrons have the highest probability of being found
Write the electron configuration for sulfur. How many electrons are in each energy
level?
1s22s22p63s23p4
1–2 2–8 3-6
For the 3rd energy level, how many of the following quantities are possible?
_1___a) s orbitals
_3___b) p orbitals
__5__ c) d orbitals
__0__ d) f orbitals
39) How many electrons can each of the following sublevels hold?
_2___e) s orbital
40)
_6___f) p orbitals
_10__ g) d orbitals
_14__ h) f orbitals
Rank the following in order of increasing energy.
4s , 1s , 3p , 3d , 6s
1s, 3p, 4s, 3d, 6s
41) Write the following out in full electron configuration.
a) Mn 1s22s22p63s23p64s23d5
b) As
42)
1s22s22p63s23p64s23d104p3
Which of the following is impossible? Why?
a) 1s22s22p63s2 b) 1s22s22p63s23p64s3
c) 1s22s22p65s1
d) 1s22s22p63s24s1
____b – s can only hold 2 electrons_________________________________________
43)
44)
Name the elements whose neutral atoms possess the following electron configuration.
a)1s22s2_beryllium__________
b) 1s22s22p4_oxygen___________________
c) ends with a 3d8 _nickel__________
d) ends with 4p3 _arsenic______________
a) Which sublevel has the lowest energy? (s, p, d, or f)
b) What is the highest sublevel with electrons in it in the element Silicon? 3p
46)
Why are the electron configurations of chromium and copper exceptions to the filling
rule?
one electron is promoted from the s to the d level to get a ½ filled and a full sublevel
(more stable)__________________________________________________________________
47)
Write the kernal electronic configuration for the following elements and write the
number of valence electrons each has in the space before the symbol.
__1__ a) Rb [Kr]5s1
48)
__6__ b)
Se [Ar]4s23d104p4
Define electronegativity. _the tendency of an atom to attract electrons to itself in a
chemical bond_______________________________________________________________
Chapter 6
49)
The modern periodic table is arranged by increasing _atomic
number____________________________
50)
Elements are arranged into _groups__ having similar chemical properties and
electronic configurations. Therefore, Se would be chemically more like (As/S).
(circle one)
51)
Write the symbol of a) any transition element Fe b) any alkaline earth metal Ba #2
c) any noble gas element Ne #18 d) the lightest alkali metal Li #1 e) the heaviest
halogen. At #17 Give the group number for parts b - e.
52)
Identify where the metals, nonmetals, metalloids and noble gases are on the P.T.
(Draw a picture)
53)
How many valence electrons are in _2____ a) calcium _6____ b) sulfur
54)
What are the periodic trends of IE (ionization energy) across and down?
IE decreases down_____________________________________________
IE increases across_____________________________________________
55)
Give the major reason why atomic radii decrease across a period and increase down a
group?
Across – more protons, same energy level, greater nuclear pull_________________
Down – electrons are being added to energy levels, less pull____________________
56)
The most reactive metal is francium and the most reactive nonmetal is fluorine.
57)
Metals generally want to (lose, gain) electrons and non-metals usually have ( fewer,
more) valence electrons.
58)
As the size of an atom increases its attraction for outer electrons (increases/decreases)
making the atom have (high/lower) ionization energy
59)
For each set, circle the atom or ion which has the larger volume.
a) Cl+, Cl, Cl–
b) Na, Na+
Chapter 7, 8 , 9
60)
According to the octet rule, phosphorous will gain 3_ electrons when it bonds.
61)
Decide whether the following are descriptions of non polar covalent bonds (C), ionic
bonds (I), or polar covalent bonds (P).
__P___ Is the result of unequally shared electrons.
__P___ Has dipoles.
__I___ Electrons are transferred from one atom to another.
__C___ Electrons are shared evenly.
..
+
:
__I___ Na Cl :–
..
__C___ N2
62)
(atoms of the same element)
Write the symbols of the following ions:
__S-2___ a) sulfide ion
63)
_Ca+2____ b) calcium ion
__Fe+3___ c) iron (III) ion
How can you tell if two elements will form ionic bonds?
If they are a metal and a nonmetal
64)
Write the electron configuration for the F- ion and the Ca2+ ion.
F- : 1s22s22p6
Ca+2 : 1s22s22p63s23p6
65) Why are metals good conductors?
Their electrons can easily slide past each other.
66)
Draw the Lewis Dot structure for Ca and Oxygen separately and then as a compound.
67)
Circle the following elements if they tend to form cations. Put a box around anions.
a) Ca
68)
b) Cl
c)
Sn
d. O
Electrons involved in chemical bonding come from (incomplete/full) (outer/inner)
energy levels.
69)
Use the octet rule to predict whether the element selenium will gain or lose electrons
when it bonds and write the kernal (abbreviated) notation of the selenium ion.
Gain Se-2
[Ar]4s23d104p6
70) What are the vertical and horizontal periodic trends for electronegativity?
Vertical – decreases
71)
horizontal - increases________________________________
Classify the following as ionic, polar, or covalent.
a) C-O
polar
b) H-H
covalent
c) Na-F
ionic
72)
73)
74)
75)
What is the relationship between size and the amount of dispersion?
Bigger size = greater dispersion
Compare the 3 types if IMFs in terms of strength.
H-bonding > dipole-dipole > dispersion
How would the viscosity of 3 liquids compare if each had a different type of IMF?
H-bonding most viscous, dispersion least viscous
Write the formula of the given compound. (You may need a polyatomic ion chart)
Then, say if the compound is ionic or molecular.
_Li2O ionic______________________________ a) lithium oxide
_Ca(OH)2 ionic______________________________ b) calcium hydroxide
_Pb(OH)2 ionic______________________________ c) lead (II) hydroxide
_NO3 molecular______________________________ d) nitrogen trioxide
_H2S molecular______________________________ f) hydrosulfuric acid
_S2O5 molecular______________________________ e) disulfur pentoxide
_ZnCO3 ionic______________________________ f) zinc carbonate
_HClO2 molecular______________________________g) chlorous acid
76)
Write the name of the following chemical compounds.
__silver nitride_____________________________ a) Ag3N
__sodium sulfate_____________________________ b) Na2SO4
__triphosphorus pentanitride_________________ c) P3N5
__hyposulfurous acid________________________ d) H2SO2
__dinitrogen monoxide_______________________ e) N2O
__copper (I) oxide____________________________ f) Cu2O
__copper (II) sulfate pentahydrate____________ g) CuSO4 . 5H2O
__hydronitric acid___________________________ h) H3N
77) Complete the chart
Molecule
Dot Structure
Polar?
Geometry
Hybrid
IMF
H2Se
yes
bent
sp3
D -D
KCl
(ionic)
XXXXXX
XXXXXX
XXXXXX
XXXXXX
XXXXXX
XXXXXX
XXX
XXX
XXX
CO3-2
no
Trigonal
planar
sp2
no
tetrahedral
sp3
disp
no
linear
sp
disp
yes
pyramidal
sp3
XXX
XXX
XXX
Disp
(resonance)
SiCl4
CO2
NH3
78)
79)
H-bond
Compare the properties of the four types of solids (ionic, molecular, metallic and
network)
Ionic – brittle, conducts when melted or dissolved, high melting point, s, metal and
nonmetal, transfer of electrons, ex: NaCl
Molecular – soft, never conducts, low melting point, s.l,g, all nonmetals, share
electrons, IMF’s, ex: H2O
Metallic – always conduct, usually solid except Hg, all metal, “sea of electrons”, ex: Au
Network – do not cnduct, solid, very high melting point, held together by covalent
bonds, ex: diamond
Compare the length and strength of single, double and triple bonds.
Single – longest, weakest
Double
80)
Triple – shortest, strongest
When do we draw resonance structures?
When the double bond could go to more than one atom. When a double/triple bond
can rotate to another site.
Chapter 10
81) What is the molar mass of:
_58.7 g/mol______ a) Ni
_159.6 g/mol______ b) CuS04
82)
How many atoms in.....
a) 14g S
b) 2 moles of Sn
14 g S x 6.02 x 1023atoms = 2.6 x 1023atoms 2 mol Sn x 6.02 x 1023 atoms = 1.2 x 1024 atoms
32.1 g
1 mol
83)
a) How many moles in 48.6 g Ca
48.6 g Ca x 1 mol = 1.22 mol Ca
40 g
b) What is the mass of 6 moles of NaOH?
6 mol NaOH x 40 g = 240 g NaOH
1 mol
84)
85)
86)
87)
What is the mass of 6.72 L of carbon dioxide gas at STP?
6.72 L CO2 x 44 g CO2 = 13.2 g CO2
22.4 L
What are the percent compositions of the indicated elements in the following
compounds?
a) Mo in MoCl
b) Cd in Cd(HCO3)2
mm = 131.5 g/mol
mm = 234.4 g/mol
96
x 100 = 73%
112.4 x 100 = 48%
131.5
234.4
A binary compound is made up of 36 g of carbon and 48 g oxygen. What is its
empirical formula?
36 g C x 1 mol = 3 mol C/3 = 1
CO
12 g
48 g O x 1 mol = 3 mol O/3 = 1
16 g
a) A compound has an empirical formula of CH2 and a molecular mass of 112 grams.
What is the molecular formula of this compound?
efm: 14 g
112/14 = 8
C8H16
b) Could C2H6 be an empirical formula?
NO
88)
A binary compound has a molecular mass of 86 grams. If a sample of this compound
contains 36 g of C and 7 g of hydrogen, what is its molecular formula?
36 = 3 7 = 7
C3H7
molecular = C6H14
12
1
Chapter 11
89) Write, complete, and balance the following reactions:
a) sodium phosphate + barium oxide —>
2 Na3PO4 + 3 BaO  Ba3(PO4)2 + 3 Na2O
b) solid potassium reacts with magnesium chloride to produce...
2 K + MgCl2  2 KCl + Mg
c) sodium oxide decomposes to yield ..
2 Na2O  4 Na + O2
90)
91)
What do the following stand for?
a) (aq)
b) (s)
c) (l)
d) (g)
aqueous
solid
liquid
gas
Balance the following equations and classify them as synthesis, decomposition, single
replacement, double replacement or combustion reactions by placing the proper
letter(s) in the space before the equation.
__S___a) N2
+ 3 H2
__S or C___b) 2 C
__DR___c) FeCl3
__SR___d) 2 K
+
+
—> 2 NH3
O2
—> 2 CO
3 KOH
—>
Fe(OH)3
+ 3 KCl
+ 2 H2O —> 2 KOH + H2
__D___e) 2 NaHCO3
—>
Na2O
+
H2O
+
2 CO2
92)
List the symbols of the seven elements that are diatomic in nature.
H, N, O, F, Cl, Br, I
93)
Complete the following equations and balance them.
a)
2 Al(ClO3)3 + 3 Mg
b)
2 Al
+ 6 HCl
—> 3 Mg(ClO3)2 + 2 Al
—> 2 AlCl3 + 3 H2
c) 3 Ca(OH)2 + 2 H3PO4
—> Ca3(PO4)2 + 6 H2O
94)
When equation 93 a) is correctly balanced, how many total oxygen atoms are
represented on the reactant side of the equation? How many must appear on the
product side? Explain why this is true.
18  18 Law of conservation of matter
95)
Write the net ionic equation for the reaction between AgNO3 and NaBr.
Identify the spectator ions.
Complete: AgNO3 + NaBr  AgBr(s) + NaNO3(aq)
Ionic: Ag+1 + NO3-1 + Na+1 + Br-1  AgBr(s) + Na+1 + NO3-1
Net ionic: Ag+1 + Br-1  AgBr(s)
Spectator ions : Na+1 and NO3-1
Chapter 12
96)
Use this equation for the following problems.
2 Fe (s)
+ 3 I2(s)
—> 2 FeI3 (s)
a) balance the equation
b) when five moles of iodine are used, how many moles of iron(III) iodide are
produced?
5 mol I2 x 2 mol FeI3 = 3.3 mol FeI3
3 mol I2
c) If 25g of iron are used, calculate how many grams of FeI3 can be produced?
25 g Fe x 1 mol Fe x 2 mol FeI3 x 436.5 g FeI3 = 195.4 = 2.0 x 102 g FeI3
55.845 g Fe 2 mol Fe
1 mol FeI3
97) For the following equation:
CH4(g) + 2 O2(g) —>
CO2(g)
+ 2 H2O(g)
a) Balance it.
b) If you are interested in producing 50 liters of CO2, how many liters of O2 would
you have to start with?
50 L CO2 x 1 mol CO2 x 2 mol O2 x 22.4 L O2 = 100 L O2
22.4 L CO2 1 mol CO2 1 mol O2
c) If you start with 20.0 g CH4 and have excess O2, how many molecules of CO2 are
made?
20.0 g CH4 x 1 mol CH4 x 1 mol CO2 x 6.02 x 1023 molecules = 7.5 x 1023 molecules
16 g CH4
1 mol CH4 1 mol CO2
98. For the following reaction:
Zn
+
2 HCl
—>
ZnCl2
+
H2
a) Balance it.
b) How many grams of Zn are required to produce 4 liters of H2 at STP?
4 L H2 x 1 mol H2 x 1 mol Zn x 65.39 g Zn = 11.7 g = 10g
22.4 L H2 1 mol H2 1 mol Zn
c) Which would be the limiting reactant if 50 g of each reactant were added together?
50 g HCl x 1 mol HCl x 1 mole H2 = 0.69 moles H2 Less! HCl is limiting.
36.45 g HCl
2 mol HCl
50 g Zn x 1 mol Zn x 1 mol H2 = 0.77 moles H2
65.3 g Zn
1 mol Zn
Chapter 13 & 14
99)
100)
What are the parts of the kinetic molecular theory?
1. small hard spheres w/ insignificant volume
2. rapid, random motion – different speeds
3. elastic collisions
4. particles can trade energy upon collision
See page 385 for more
a) Sketch a kinetic energy distribution graph for two different temperatures and
label the curve of the lowest temperature. b) If you could see the particles of a gas
at room temperature, describe their motion.
a.
b. some fast, some slow, colliding and
transferring energy, large spaces
between them
101)
102)
103)
104)
105)
What is vapor pressure? What is the relationship between the temperature and
the vapor pressure? What is the relationship between the vapor pressure and the
boiling point?
The pressure exerted by a liquid that has evaporated.
Higher temperature  higher vapor pressure
Higher vapor pressure  lower boiling point
How is boiling point related to the atmospheric pressure?
Higher atmospheric pressure  higher boiling point
What is a unit cell?
Smallest group of particles that is repeated
What is an amorphous solid? How is it different from crystalline solids? Describe
properties of each type. Compare the arrangement of particles in solids, liquids
and gases.
A super cooled liquid. Particles are randomly arranged as in a liquid, no real
boiling point
Solids – orderly and close
Liquids - random and close
Gas – random and spread out
Sketch a phase diagram and label its axes, phases, and phase changes.
106)
Determine the pressure of a trapped gas using a manometer if the difference in Hg
levels on the two sides is 10 mm Hg and it is lower on the side open to the
atmosphere. (Air pressure is 675 mm Hg on that day.)
665 mmHg
107)
How are volume and pressure related? If the volume of a gas is 20 ml at a pressure
of 30 mm Hg and the pressure is increased to 50 mm Hg, what is the new volume?
Indirectly
(20 mL)(30 mmHg) = V2(50 mmHg)
V2 = 12 mL = 10 mL
108)
How is the volume of a gas related to the temperature? If a temperature is doubled
what happens to the volume?
Directly, doubled 1 = x
x=2
1
2
What is the difference between real gases and ideal gases?
All gases are real. Gases behave “ideally” (follow gas laws) at high temperatures
and low pressure
What is the total pressure of a mixture of H2 , O2 and Ne if PH2 = 0.20 atm, PO2 =
0.4 atm and PNe = 0.2 atm.
0.20 + 0.4 + 0.2 = 0.8 atm
How is the rate of diffusion related to molecular weight?
Lighter  faster
If pressure is cut in half and the temperature is doubled, what happens to the
volume?
P1V1 = P2V2
(2)(1) = (1)V2
V2 = 4 quadrupled
T1
T2
1
2
What volume do two moles of O2 occupy at S.T.P.? Of N2?
2 x 22.4 = 44.8 L O2
44.8 L N2
109)
110)
111)
112)
113)
114)
If you have 2 moles of a gas, what would be the volume of the gas at 27° C and
800 mmHg?
(800/760)V = (2)(0.0821)(300)
V = 46.8 L = 50 L