CHEMISTRY 1000 A
Fall 2004
ATOMS, MOLECULES AND CHEMICAL REACTIONS
Instructor: Dr. René T. Boeré
Review Materials
This .PDF document contains two copies of previous final exams from similar offerings of Chemistry 1000
that I am making available to currently registered students of my Chemistry 1000 class for the sole purpose
of self-study and review. The content of this exam may not reflect exactly the content of the course you are
taking. User beware! Indeed, several questions have been marked as not relevant to the Fall 2004
offering of the course. Such questions would be replaced with other, more relevant questions, if these tests
were devised for the Fall 2004 class.
The purpose of this document is to give you a good idea of how I structure the exams, and of approximately
how long they tend to be. They may remind you of material we covered near the beginning of the course,
which you may tend otherwise to neglect.
A good way to use these practice exams would be to write them under simulated test conditions – lock
yourself in your room for 3 hrs and try to write the thing with the text-book and your notes closed. You may
find that you are already well prepared for the final, or you may detect certain subject areas that you are weak
in. Then you could focus your studies on those areas. You could use one of the tests near the beginning of
your review work and the other near the end, so see for yourself how your studying has developed your
understanding of the material.
An answer guide to this practice exam will be released towards the end of the week. I strongly caution you
from just looking at my answers, however. You must try to work out your own answers to the best of your
ability. Then compare your efforts with my suggested answers as a way of checking any false assumptions
you may have made. On the whole, I think you will sense as you try to write the sample tests just where you
are strong and where you are not.
Best of luck studying for the final exam!
Prof. Boeré
THE UNIVERSITY OF LETHBRIDGE
DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY
CHEMISTRY 1000
Instructor: Dr. R.T. Boeré
Practice Final Examination #1
(3 Hours)
No. of Pages = 8 + 1
Caveat: Note that this exam is provided only to give you an idea of the kind of format you may face in this year's final exam, and as a
way of checking the state of your own preparation. It does not in any way limit the type or range of questions that will be on the
final. Questions may be drawn from any portion of the course. The definition of material to be studied is those sections of the
textbook identified in the eleven official assignments handed out during the course of the semester. Happy studying, and good luck.
Question
Mark
Question
Mark
/value
#1
Question
/value
#5
/4
#6
/10
#7
/14
#11
/8
#4
/8
#10
/9
#3
/value
#9
/6
#2
Mark
/19
#8
/14
Total
/11
/17
/120
1. Students must count the number of pages in this examination paper before beginning to write, and report
any discrepancy immediately to the invigilator. Answer all questions on the paper.
2. The value of each question is given in the margin as the number in square brackets.
3. Some data, including a periodic table, is provided on a separate page within this exam booklet; the data
sheet need not be returned with the completed examination.
4. Calculators and slide rules are permitted. All alphanumeric memories MUST be cleared before the exam
commences.
5. The exam booklet must be handed in before leaving the examination hall. No-one will be permitted to leave
the examination hall during the last 15 minutes of the exam period.
1.
a) Write acceptable names for each of the following:
[6]
(i) KHSO3
(ii) (NH4)2Te
b) Indicate which of the following is a hydrocarbon, which a carbohydrate, which a strong acid, which
a weak base. Only one example of each type is included! Leave the rest blank. Marks will be
subtracted for false identification.
i) H3PO4
ii) C21H28
iii) CsOH
iv) NH3
v) HNO3
vi) C16H28O14
vii) C3H7Br
viii) AgCN
2.
[9]
For each of the following, write a balanced chemical equation which includes correct formulas for all
species involved in the reaction. Include the solvent, water, only if it is a net reactant or product of the
reaction. Indicate the state of each reactant/product, e.g. (s), (l), (g) or (aq).
a) Chromium(II) nitrate in aqueous solution reacts with gaseous hydrogen sulfide to produce chromium(II)
sulfide.
b) Rubidium carbonate in aqueous solution reacts with aqueous perchloric acid and a vigorous gas evolution
occurs.
c) Manganese(II) nitrate in aqueous solution reacts with solid trisodium phosphate.
3.
a) Rhenium occurs in nature as a mixture of two isotopes:
amu, and
isotopes.
187
75
185
75
Re , which has an atomic mass of 184.953
Re , which has an atomic mass of 186.956 amu. Determine the percent abundance of the two
[8]
39
40
40
b) Potassium has three naturally occurring isotopes, 19
K , 19
K and 1941K , but 19
K has a very low natural
abundance. Which of the other two is the more abundant? Briefly explain your answer.
2
4.
a) Identify the missing component in each of the following nuclear reactions, and classify the type of
nuclear reaction that is involved.
[11]
i)
230
92
ii)
24
10
U
Ne →
iii)
51
24
iv)
119
51
60
30
v)
→
Cr
226
90
Th
+
24
11
Na
+
+
→
Sb +
0
−1
Zn →
60
29
Mn +
51
25
2 01n
e →
Cu
+
b) Predict the most likely mode of nuclear decay for the following isotopes. Explain your reasoning!
i)
137
53
ii)
37
18
iii)
5.
I
Ar
230
90
Th
In a rigid, inert, sealed container the following reaction is carried out:
Si (s)
+
2 CH3Cl (g)
→
(CH3)2SiCl2 (g)
[4]
Before the reaction commences, the total pressure in the flask is 3.000 atm. After the reaction has gone
to completion (assume 100% yield), the container is returned to the same temperature it was at before the
start of the reaction, and the pressure is now measured to be 2.000 atm. What was relative amount of
initial reactants in the flask? In other words, determine the ratio Si (s)/ CH3Cl (g).
3
6.
Answer only ONE of the following possible essay questions. No credit will be given for answering
more than one question, so if you change your mind, clearly strike out anything that should not be
graded.
[10]
Option A: Describe nuclear fission as a source of electric power. Focus your discussion on the
particulars of the Canadian CANDU design.
Option B: Describe in detail the concept of radiation safety. Include the nature of the hazard, how risk is
assessed, and what protective measures need to be taken to work with radioactive isoptopes in a
laboratory setting.
Option C: Discuss in detail the use of the 99mTc isotope as an imaging agent in nuclear medicine. Include
applications in use at the Lethbridge Regional Hospital.
4
7.
At 180.0 °C, 27.20 mg of the antimalarial drug quinine in a 250.0 mL container gave rise to a pressure of
9.48 mmHg. In a combustion analysis, 16.39 mg of quinine produced 44.41 mg of carbon dioxide, 10.83
mg of water, and 1.46 mg of nitrogen. Assuming that it is a compound of only C, H, N and O, calculate
the empirical and molecular formulae of quinine.
[19]
8.
a) Determine the complete ground-state electron configurations of the following atoms or ions:
i) N
[17]
ii) F
iii) Br
iv) Br–
b) With reference to the configurations in (a), answer the following, and justify your choice in detail:
i) Which of N or F has the larger radius?
ii) Which of Br or Br– has the larger radius?
5
iii) Which of F or Br has the larger ionization energy?
c)
Use “boxes” notation for the valence electrons of a ground state cobalt atom, Co. Assign a valid
combination of all four quantum numbers for each valence electron. Explain clearly whether such a
cobalt atom would be expected to be diamagnetic or paramagnetic.
9.
If energy is absorbed by a hydrogen atom in its ground state, the atom is excited to a higher energy state.
a) How much energy is required to excite a ground state hydrogen atom to the n = 4 state?
[8]
b) If, as is commonly the case, this excitation is caused by the absorption of a single photon per hydrogen
atom, what is the wavelength of the electromagnetic radiation required for this excitation process?
c) What is the longest wavelength of electromagnetic radiation capable of ionizing a hydrogen atom?
d) Sketch an outline diagram which clearly shows the shape of (a) the px , (b) the dz2 , and (c) the s orbital
on the xyz grids provided. Pay attention to the labels on the axes. Shade the orbital lobes to indicate
phase.
z
a)
z
b)
y
x
z
c)
y
x
y
x
6
10.
The compound oxygen difluoride is quite unstable, giving oxygen and hydrofluoric acid on reaction with
water vapour, according to the equation:
[14]
OF2(g)
+
H2 O
→
O2(g)
+
2 HF(g)
a) Write Lewis structures for each reactant and product
b) Use bond energies to estimate the enthalpy change, ∆Hrxn, accompanying this reaction.
c) If 24.21 g of OF2 are reacted with 7.803 g of water vapour, what is the maximum amount of HF that
can be produced?
d) How much energy would be released or absorbed during the reaction described in (c)?
e) If 14.01 g of HF is recovered in the actual reaction performed in (c), what is the %-yield of the reaction?
7
11.
[14]
For the following molecules or ions determine:
a) The Lewis dot diagram; show formal charges; include important resonance isomers if there are any.
The central atom is underlined.
b) The VSEPR structure (sketch using the “wedges” notation); write down the name of the shape of the
molecule or ion. Include bond angles and indicate as exact or approximate (Hint: use the ~ symbol to
indicate approximate angles).
c) The hybridzation at the central atom.
Lewis (electron distribution)
VSEPR (shape)
Hybridization
PF5
CS32–
NF3
XeF4
8
DATA SHEET
Fundamental Constants
Bond Energies (kJ mol–1)
1.6605 × 10-24 g
6.022 × 1023
1.6022 × 10-19 C
9.1095 × 10-28 g
96,487 C mol-1
8.314 J K-1mol-1
0.08206 L atm K-1mol-1
6.626 × 10-34 J s
1.67252 × 10-24 g
1.6749 × 10-24 g
2.998 x 108 m s-1
101.325 kPa = 760.0 mmHg
1.097 x 107 m-1
Atomic mass unit
Avogadro's number
Electron charge (e)
Electron mass
Faraday constant (F)
Gas constant
Planck's constant
Proton mass
Neutron mass
Speed of light in vacuum
Standard atmosphere
Rydberg Constant
Bond type
H–H
O–F
O–H
O–O
O=O
H–F
F–F
Energy
436
184
464
138
498
569
159
Some common formulas
n=
m
MM
d=
PV = nRT
1
1
1
=R 2 − 2
λ
ninitial n final
m
V
E = hν =
∆E = E final − Einitial = Rhc
PT = PA + PB + PC
(
PA = X A ⋅PT
1
n
2
final
−
hc
λ
1
2
initial
n
λ =
c
ν
E = mc 2
λ=
h
mv
∆x ⋅∆( mv ) >
) (
∆H rxn = Σ ∆H 0f prod − Σ ∆H 0f react
h
2π
)
Periodic Table of the Elements
1
18
1.008
4.003
H
1
6.939
2
13
14
15
9.012
10.811
12.011
14.007
Li
Be
3
22.990
Na
11
39.102
Mg
12
40.08
K
Ca
19
85.47
20
87.62
3
44.956
38
137.34
4
47.90
21
88.905
Y
39
138.91
5
50.942
V
Ti
Sc
Sr
Rb
37
132.91
5
26.982
4
24.312
22
91.22
Zr
40
178.49
23
92.906
La
Hf
Ta
57
(227)
72
(257)
73
(260)
Ra
Ac
89
Rf
24
95.94
Mo
Ba
88
Mn
42
183.85
56
(226)
Fr
54.938
Cr
Nb
Cs
W
105
140.12
Tc
43
186.207
Re
55.847
Fe
26
101.07
Ru
44
190.2
9
Al
Si
13
69.72
14
72.59
11
12
58.933
58.71
63.546
65.37
Co
Ni
Cu
Zn
Ga
Ge
As
27
102.90
28
106.4
29
107.87
30
112.40
31
114.82
32
118.69
33
121.75
Rh
45
192.2
Pd
Ag
Cd
46
195.09
47
196.97
48
200.59
Pt
Ir
Os
Au
In
49
204.37
Tl
Hg
8
32.064
15
74.922
F
9
35.453
Cl
S
P
10
18.998
O
16
78.96
Se
34
127.60
17
79.904
He
2
20.183
Ne
10
39.948
Ar
18
83.80
Br
Kr
35
126.90
36
131.30
Sn
Sb
Te
I
Xe
50
207.19
51
208.98
52
(210)
53
(210)
54
(222)
Bi
Pb
Po
At
74
75
76
77
78
79
80
81
82
83
84
85
140.91
144.24
(147)
150.35
151.96
157.25
158.92
162.50
164.93
167.26
168.93
173.04
174.97
Sm
Eu
Gd
Tb
Dy
Ho
Tm
Yb
Ce
Pr
Nd
58
232.04
59
(231)
60
238.03
Th
25
(99)
8
7
30.974
17
Rn
86
Ha
104
90
7
51.996
41
180.95
55
(223)
87
6
6
28.086
15.9994
N
C
B
16
U
Pa
91
92
Pm
61
(237)
Np
93
62
(242)
Pu
94
63
(243)
Am
95
64
(247)
Cm
96
65
(247)
Bk
97
66
(249)
Cf
98
67
(254)
Es
99
Er
68
(253)
Fm
100
69
(256)
Md
101
70
(253)
No
102
Lu
71
(257)
Lr
103
9
THE UNIVERSITY OF LETHBRIDGE
CHEMISTRY 1000
Instructor: Dr. R.T. Boeré
(3 Hours)
DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY
Practice Final Examination #2
No. of Pages = 8 + 1
Caveat: Note that this exam is provided only to give you an idea of the kind of format you may face in this year's final exam, and as a
way of checking the state of your own preparation. It does not in any way limit the type or range of questions that will be on the
final. Questions may be drawn from any portion of the course. The definition of material to be studied is those sections of the
textbook identified in the eleven official assignments handed out during the course of the semester. Happy studying, and good luck.
Question
Mark
Question
Mark
/value
#1
Question
/value
#5
/12
#6
/25
#7
/10
#11
/10
#4
/8
#10
/9
#3
/value
#9
/8
#2
Mark
/18
#8
/16
Total
/10
/20
/146
1. Students must count the number of pages in this examination paper before beginning to write, and report
any discrepancy immediately to the invigilator. Answer all questions on the paper.
2. The value of each question is given in the margin as the number in square brackets.
3. Some data, including a periodic table, is provided on a separate page within this exam booklet; the data
sheet need not be returned with the completed examination.
4. Calculators and slide rules are permitted. All alphanumeric memories MUST be cleared before the exam
commences.
5. The exam booklet must be handed in before leaving the examination hall. No-one will be permitted to leave
the examination hall during the last 15 minutes of the exam period.
1.
a) Write acceptable names for each of the following:
(i) NaHSO4
(ii) Fe(OH)3
[8]
(iii) Mg(NO2)
(iv) CdF2
b) Indicate which of the following is a hydrocarbon, which an alcohol, which a weak acid, which a
strong base. Only one example of each type is included! Leave the rest blank. Marks will be subtracted
for false identification.
i) H2SO4
ii) C6H6
iii) FeBr3
iv) KOH
v) HF
vi) C2H5OH
vii) Zn(OH)2
viii) CH2O
2.
Predict the course of the following equations and state the products. Balance each equation and classify
it as (i) acid-base, (ii) gas-forming, (iii) precipitation, or (iv) oxidation-reduction.
[9]
3.
a)
Pb(NO3)2 (aq)
b)
HNO3 (aq)
c)
Ca (s)
+
+
KBr (aq) à
+
RbOH (aq) à
O2 (g)
à
A 0.400 g sample of sodium azide, NaN3, is heated and decomposes according to the equation:
2 NaN3 (s) → 2 Na (s) + 3 N2 (g)
The gas is collected over water (i.e. by bubbling the gas through water into an inverted container initially
filled with water), and the pressure inside is equalized to the atmospheric pressure. If the volume of gas is
measured as 225 mL at a temperature of 22°C and a barometric pressure of 0.980 atm, what is the %-yield
of the reaction?
[10]
2
4.
Identify the missing component in each of the following nuclear reactions, and classify the type of nuclear
reaction that is involved. Each box corresponds to a single nuclear particle.
[10]
i)
19
10
ii)
230
90
Th →
iii)
246
96
iv)
37
18
v)
5
Ne →
59
26
0
+1
+
e
+
226
88
Ra
Cm +
Ar
+
Fe →
→
0
−1
254
102
No +
4 01n
e →
59
27
Co
+
Not covered in 2004 - ignore! a) Draw the structure of each of the following organic molecules
(i) 2-methyl-3-ethylpentane
(ii) 3,3-dimethyl-1-butanol
[12]
(iii) 2-methylhexanoic acid
b) Draw all possible isomers with the molecular formula C4H10O. Marks will be subtracted for duplicate
structures.
3
6.
a)
Technetium is not found naturally on earth; it must be synthesized in a nuclear facility. In the form of
sodium pertechnetate, NaTcO4, it has uses in nuclear medicine.
In what group and period of the periodic table is the element found?
b)
Write out the full electron configuration for a ground state Tc atom.
[25]
c)
Write out a full set of quantum numbers (n, l, ml, ms) which is valid for any one of the valence d electrons
of Tc.
d)
Technetium emits a γ(gamma) ray with an energy of 0.140 MeV. (1 MeV = 1 million electron volts
where 1 eV = 1.60218 × 10–19 J). What are the wavelength and frequency of a γ
-ray photon with an
energy of 0.140 MeV?
e)
To make normal NaTcO4, the metal is dissolved in nitric acid according to the equation:
7 HNO3 (aq) + Tc (s) à HTcO4 (aq) + 7 NO2 (g) + 3 H2O (l)
and the product, HTcO4, is treated with NaOH to make NaTcO4. Write a balanced equation for the
reaction of HTcO4 with NaOH. If you begin with 4.5 mg of Tc metal, how much NaTcO4 can be made?
What mass of NaOH (in mg) is required to convert all of the HTcO4 into NaTcO4?
4
f)
Discuss briefly the use of the 99mTc isotope as an imaging agent in nuclear medicine. Include applications
in use at the Lethbridge Regional Hospital.
7.
Maleic acid is prepared by the catalytic oxidation of benzene. It is a dicarboxylic acid, that is, it has two
carboxylic acid groups (an organic functional group), and is therefore a compound of C, H and O.
a) Combustion of 0.125 g of the acid gives 0.190 g of CO2 and 0.0388 g of H2O. Calculate the empirical
formula of the acid.
]
b) A 0.261 g sample of the acid requires 34.60 mL of 0.130 M NaOH for complete titration (so that the
H+ ions from both the carboxylic acids are used). What is the molecular formula of the acid?
c) Draw a Lewis structure for the acid (i.e. an organic structural formula showing all non-bonded electron
pairs.) Hint: the molecule can be thought of as a fragment of benzene.
d) Describe the hybridization used by the C atoms, and the bond angles around the C atoms.
5
8.
a)
Answer the following questions:
i) The quantum number n describes what
property(ies) of an atomic orbital?
[20]
ii) The quantum number l describes
what property(ies) of an atomic orbital?
iii) The quantum number ml describes
what property(ies) of an atomic orbital?
b)
What type of orbital corresponds to l = 3?
c)
Each drawing below represents one specific type of atomic orbital. Give the letter designation, the l
value, and specify the number of nodal planes for each orbital in the spaces indicated below:
z
(i)
z
(ii)
y
x
z
(iii)
y
x
y
x
Letter designation:
__________
___________
___________
l-value:
__________
___________
___________
# of nodal planes:
__________
___________
___________
d)
Which of the following orbitals cannot exist according to the atomic quantum theory: 2s, 3p, 2d, 3f, 5p,
6p?
e)
Which of the following is not a valid set of the first three quantum numbers?
n l
ml
3 2 1
2 1 2
4 3 0
f)
What is the maximum number of orbitals that can be associated with each of the following sets of
quantum numbers? ("None" is a possible answer.) Show how you get your answer!
i) n = 2 and l = 1
ii) n = 3
iii) n = 3 and l = 3
iv) n = 2, l = 1 and ml = 0.
6
9.
Not covered in Fall 2004 – ignore this question, please!
Construct a Born-Haber cycle and use it to calculate the lattice energy for crystalline potassium chloride,
KCl (c), in kJ mol–1.
10.
Describe the bonding in methylene amine, H2C=NH, under the following headings. Use sketches and
text, as appropriate. A computer model of this molecule is shown below:
[8]
a) The Lewis structure and the shape of the molecule, with approximate bond angles.
[10]
b) The σ (sigma) bonds. Include the hydrogen atoms and lone pairs. Identify the atomic or hybrid
orbitals employed in each bond.
c) The π (pi) bond(s). Identify the atomic or hybrid orbitals employed in each bond.
7
11.
[16]
For the following molecules or ions determine:
a) The Lewis dot diagram; show formal charges; include important resonance isomers if there are any.
The central atom is underlined.
b) The VSEPR structure (sketch using the “wedges” notation); write down the name of the shape of the
molecule or ion. Include bond angles and indicate as exact or approximate (Hint: use the ~ symbol to
indicate approximate angles).
c) The hybridization at the central atom.
Lewis (electron distribution)
VSEPR (shape)
Hybridization
BH3
SCN–
OF2
ClF4–
8
DATA SHEET
Fundamental Constants
Atomic mass unit
Avogadro's number
Electron charge (e)
Electron mass
Faraday constant (F)
Gas constant
1.6605 × 10-24 g
Planck's constant
23
6.022 × 10
Proton mass
-19
1.6022 × 10 C
Neutron mass
-28
9.1095 × 10 g
Speed of light in vacuum
96,487 C mol-1
Standard atmosphere
8.314 J K-1mol-1
Rydberg Constant
0.08206 L atm K-1mol-1
Vapour pressure of water at various temperatures
Vapour pressure (mmHg)
T (°C)
20
17.5
21
18.7
22
19.8
23
21.1
24
22.4
25
23.8
Useful formulae
m
m
n=
d=
MM
V
Standard Enthalpy Data
Process
(kJ mol–1)
–436.747
∆H°f of KCl (c)
Bond energy of Cl2 (g)
+243.36
Sublimation of K (s)
+89.24
st
1 IE of K (g)
+419
1st EA of Cl (g)
–349
+121.68
∆H°f of Cl (g)
PT = ∑ pi
PV = nRT
6.626 × 10-34 J s
1.67252 × 10-24 g
1.6749 × 10-24 g
2.998 x 108 m s-1
101.325 kPa = 760.0 mmHg
1.097 x 107 m-1
pi = X i ⋅PT
hc
λ
E = hν =
i
1
∆E = E final − Einitial = Rhc
n
2
final
1
−
h
mv
λ=
2
initial
n
(
h
4π
∆x ⋅∆( mv ) >
λ =
c
ν
E = mc 2
) (
∆H rxn = Σ ∆H 0f prod − Σ ∆H 0f react
∆Hrxn = Σ( BE bonds broken) − Σ( BE bonds formed )
n = M ⋅V
1
18
Chem 1000 Standard Periodic Table
4.0026
1.0079
He
H
1
6.941
2
14
10.811
12.011
15
16
17
18.9984
14.0067
15.9994
Be
B
C
N
O
3
22.9898
4
24.3050
5
26.9815
6
28.0855
7
30.9738
Na
K
19
85.4678
9.0122
13
Li
11
39.0983
Mg
12
40.078
20
87.62
Rb
Sr
38
137.327
Cs
Fr
87
3
44.9559
Ca
37
132.905
55
(223)
)
Ba
47.88
Sc
21
88.9059
Y
39
La-Lu
56
226.025
Ra
4
Ac-Lr
88
5
6
7
50.9415
51.9961
54.9380
55.847
58.9332
V
Cr
Mn
Fe
Co
26
101.07
27
102.906
Ti
22
91.224
Zr
40
178.49
23
92.9064
Nb
41
180.948
24
95.94
Mo
42
183.85
25
(98)
Tc
43
186.207
8
Ru
44
190.2
9
Re
Os
Ir
75
(262)
76
(265)
77
(266)
105
106
107
108
109
138.906
140.115
140.908
144.24
(145)
150.36
151.965
La
Ce
Pr
Nd
Pm
57
227.028
58
232.038
59
231.036
60
238.029
61
237.048
Ac
89
Th
90
Pa
91
U
92
Np
93
Pu
94
Pt
Si
14
72.61
P
15
74.9216
Zn
Ga
Ge
As
29
107.868
30
112.411
31
114.82
32
118.710
33
121.757
Au
Cd
48
200.59
In
49
204.383
Hg
Tl
Sn
50
207.19
Sb
51
208.980
Pb
Bi
F
9
35.4527
S
16
78.96
Se
34
127.60
Cl
17
79.904
Ne
10
39.948
Ar
18
83.80
Br
Kr
35
126.905
36
131.29
Te
I
Xe
52
(210)
53
(210)
54
(222)
Po
At
78
79
80
81
82
83
84
85
157.25
158.925
162.50
164.930
167.26
168.934
173.04
174.967
Tm
Yb
Rn
86
Mt
104
62
(240)
28
106.42
Ag
W
Sm
65.39
Al
13
69.723
Cu
47
196.967
74
(263)
Hs
63.546
Ni
Pd
Ta
Bh
58.693
46
195.08
73
(262)
Sg
12
Rh
Hf
Db
11
45
192.22
72
(261)
Rf
10
8
32.066
2
20.1797
Eu
63
(243)
Am
95
Gd
64
(247)
Cm
96
Tb
65
(247)
Bk
97
Dy
66
(251)
Cf
98
Ho
67
(252)
Es
99
Er
68
(257)
Fm
100
69
(258)
Md
101
70
(259)
No
102
Lu
71
(260)
Lr
103
9