Atomic Theory
Unit II
Lesson Date Topic WS
#
1. Early Atomic Theory 1
2. Bohr
Evidence 1
3. Bohr
Diagrams 1
4. Quantum Theory 2
5. Mass
Spectrometer/ Elegant Universe-1 3
6. Elegant
Universe-2/Periodic Chem 4
7. Periodic
Chemistry 5
8. Classifying
Matter Lab/ Elegant Universe-3
9. Classifying and Naming
Formulas 1 6
10. Classifying and Naming
Formulas 2 7
11. Electron
Dot Diagram Structural Formula 1
8
12. Electron
Dot Diagram 2
9
13. Practice
Test 1 10
14. Practice
Test 2 11
Worksheet #
1 Early
Atomic Theory
Briefly describe
each atomic theory listed below. Include
a diagram.
1. The Four-Element Theory
2.
3. The Thompson Atom
4. The Rutherford Atom
5. The Bohr Atom
For each of the
above, give some evidence that led to the atomic model. Briefly explain how each
evidence is accounted for in the atomic theory.
6. The Four Element Theory
(a) evidence (b)
explanation within theory
7.
(a) evidence (b) explanation within theory
8. The Thompson Atom
(a) evidence (b)
explanation within theory
9. The Rutherford Atom
(a) evidence (b)
explanation within theory
10. The Bohr Atom
(a) evidence (b)
explanation within theory
Draw Bohr atomic
diagrams for the following atoms. Be
sure to include protons, neutrons and electrons.
11. Oxygen 12.
Calcium
13. Silver 14.
Barium
15. Cs 16. I
17. Na 18. V
19. Cl- 20. Al3+
21. Se2- 22. Ca2+
Worksheet #2 Quantum Mechanics
1. What is the main difference between the
Bohr theory of the atom and the Quantum Mechanical Theory?
2. How many electrons will fill the
smallest orbital in quantum mechanical theory?
3. How is a 3s orbital different than a 2s
orbital in terms of shape and distance from the
nucleus?
4. Explain what happens to the energy when
an electron falls from a 3s orbital to a 2s
orbital.
Use your Quantum
Periodic Table to write quantum electron configurations for each element below.
5. F
6. K
7. C
8. Kr
9. S
10. Rb
11. Co
12. P
13. Ca
14. Al
15. Ag
16. 1s22s22p63s1
17. 1s22s22p63s23p5
18. 1s22s22p63s23p63d94s2
19. 1s22s22p63s23p63d104s24p64d105s25p5
20. 1s22s22p63s23p63d104s24p64d105s25p66s2
21. Give
the formula of four chemical species that are isoelectronic
(same electron configuration) as Ar.
Worksheet #3 Mass Spectrometry
Calculate the
average atomic mass for each element.
Round off to the correct number of sig
figs. Write down the atomic mass from
the periodic table rounded off to the same number of sig
figs. Show all your work.
Isotope Mass % Abundance Average
Mass
Atomic Mass(table)
1.
14N 14.0030744 99.6340
15N 15.000108 0.366001
2.
20Ne 19.992404 90.92
21Ne 20.993849 0.257
22Ne 21.991385 8.82
3.
46Ti 45.952633 7.93
47Ti 46.95176 7.28
48Ti 47.947948 73.94
49Ti 48.947867 5.51
50Ti 49.944789 5.34
4.
54Fe 53.93962 5.8202
56Fe 55.93493 91.660
57Fe 56.93539 2.1901
58Fe 57.93327 0.33001
5. Silver
has two common isotopes. One is 106.90508 amu and
51.35 % and the other is 48.65 %. If the
average atomic mass is 107.9730 amu, what is the atomic
mass of the other isotope.
6. Copper
has two common isotopes. One is 62.92959 amu and
69.09 % and the other is 30.91 %. If the
average atomic mass is 63.5472 amu, what is the
atomic mass of the other isotope.
7. Complete the
chart below.
Protons Electrons Neutrons
28Si
29Si
30Si
8. Write a quantum electron configuration
for each of the following.
a) Ne
b) Mg
c) Ti
d) Cr
e) Sr
f) Ag
g) Br
9. What was the first atomic theory to
account for the Law of Conservation of Mass?
Explain how the theory accomplished
this.
10. What was the first atomic theory to
account for electromagnetic radiation (light)?
Explain how the theory accomplished
this.
11. What was the first atomic theory to
account for the small, dense nucleus?
Explain how the theory accomplished
this.
12. What was the first atomic theory to have a
wave theory for the electron?
Explain how the theory accomplished
this.
13. What was the first atomic theory to
account for positive and negative charges in matter? Explain how the theory accomplished this.
Worksheet #4 Periodic Chemistry
1. Define the following:
a) Oxidation
b) Reduction
c) Anion
d) Cation
e) Atom
f) Chemical family
g) Period
2. Why are noble gases stable?
3. Why are non-noble gases un-stable or
reactive?
4. Draw Bohr diagrams for the following
chemical species.
a) He b)
K
c) K+ d) S2-
e) P3- f) Li+
5. Fill in the chart below.
Symbol atom, Protons Neutrons
Electrons Valence stable or
cation or Electrons reactive?
anion
Mg2+
Cation 12
12 10 8 Stable
Mg
F
F-
Ne
C
Be
Be2+
N3-
Worksheet #5 Ionic Chemistry
1. Symbol Atom, Protons Neutrons Electrons Valence stable or
Cation or Electrons reactive?
Anion
Ga atom 31 39 31 3 reactive
Ga3+
Br
Br-
Kr
Ca
Ca2+
P
P3-
2. What happens to protons, electrons and
neutrons as you move form left to right within
a row on the periodic table?
3. Write half-reactions to show how each
atom forms an ion. Label each as oxidation or
reduction. The first two are done for you.
a) K →
K+ + 1e- oxidation
b) N2 + 6e- → 2N3- reduction
c) P
d) O2
e) Ca
f) Br2
g) I2
h) Al
i) Ba
j) Cs
k) Mg
l) Zn
m) Ga
n) Cl2
o) F2
4. Describe five properties of:
a) Metals
b) Non-metals
5. Draw Bohr diagrams for each of the
following.
a) Na b)
Na+
c) O d)
O2-
e) Ca f)
Ca2+
Worksheet # 6 Classifying and Naming Formulas 1
1. Complete the
table.
Salt Base Acid Covalent
Nonacid
Litmus
Conductivity
2. Put each formula into the table below.
Ca(OH)2 NH4OH CH3OH C12H22O11
HCl PI3 K2SO4 RbOH
H3PO4 NaOH CaCl2 Li2SO4
H2SO3 BaF2 BCl5 CH3COOH
H2CO3 CsOH S2Cl2 Fr2S
Salt Base Acid Covalent Nonacid
3. Draw Bohr diagrams for each of the
following.
a) K+ b)
P3-
4. Write half-reactions to show how each
atom forms an ion. Label each as oxidation or
reduction. The first two are done for you.
a) Ca → Ca2+ + 2e- oxidation
b) O2 +
4e- → 2O2- reduction
c) I2
d) N2
e) Cs
f) Ba
g) Al
h) F2
i) H2
j) Na+ +
1e- ®
Na(s)
k) N3- ® N2 + 6e-
l) Ca2+
m) Ba2+
Worksheet # 7 Classifying and
Naming Formulas 2
1. Complete the
following table by classifying and naming each compound
Formula Classification Name
CuS(s)
H3PO4(s)
P2O5(s)
NH4OH(s)
Al2O3(s)
MgSO4(s)
HCl(g)
HCl(aq)
H2SO4(l)
H2SO4(aq)
NI3(s)
N3Cl3(s)
CO(g)
K2CrO4(s)
H2Cr2O7(aq)
H2O(l)
CrCO3(s)
HBr(g)
P3O5(s)
.
Complete the
following table by classifying and naming each compound.
2. Formula Classification Name
HI (aq)
(NH4)3PO4(s)
NCl3(l)
Ba(OH)2(s)
Rb2SO4(s)
CuCl2(s)
Al2O3(aq)
N3Cl3(aq)
CO(g)
H2SO3(aq)
CuSO4
.
6H2O(aq)
H3PO3(s)
Mg3(PO4)2(aq)
HCH3COO(aq)
HF(aq)
N2O5(aq)
Na3PO4
. 5H2O(aq)
Ni(NO3)3(aq)
SO(g)
Use your Quantum
Periodic Table to write quantum electron configurations for each element below.
3. Sr
4. V
5. 1s22s22p63s2
6. 1s22s22p63s23p3
7. 1s22s22p63s23p64s23d4
Pick the best answers. Answers can be used more than once.
Answers: Four Element Theory
Thomson’s Atomic Theory
Bohr’s Atomic Theory
Quantum Mechanical
Theory
8. The 1st
model of the atom to explain the gold foil experiment
9. The 1st
model to explain light
10. The 1st
model to account for the Law of Constant Composition
11. The 1st
model to have a small, dense nucleus
12. The 1st
model to have an electron as a wave
13. Non-scientific
Theory
14. The 1st
model to have electrons
15. The 1st
model to account for the Law of Conservation of Mass
16. Modern theory
of the atom
17. The 1st
model to claim the atom is mainly “empty space”
Worksheet #8 Electron Dot Diagrams
Draw structural
and electron-dot diagrams for each.
Structural Dot-Diagram
CH4
CI4
S2
P2
C2Cl6
C2F4
NF3
CS2
N2Cl2
HCN
CH4N2O
(Symmetrical)
C6H6
(Cyclic)
CF4
N2Cl4
NBr3
Name
each compound
1.
CH3COOH(aq)
2.
HBr(aq)
3.
HF(g)
4.
HNO3(aq)
5.
HClO4(aq)
Write
the quantum electron configurations for the following.
6.
Cl-
7.
Sr2+
8.
I
Write
a dissociation equation for each to show how each ionizes in water.
9.
CH3COOH(l)
10.
HNO3(l)
11.
Al2(SO4)3(s)
12.
Co3(PO4)2(s)
Name
each compound above.
13.
14.
15.
16.
17. Classify and name the following compounds.
NaOH
BaF2
BCl5
CH3COOH
H2CO3
CsOH
S2Cl2
BaCl2
Worksheet
# 9 Electron Dot
Diagrams
Draw structural
and electron-dot diagrams for each.
ClO3- PO43-
IO3- BrO3-
CN- NO3-
SO42- CaCO3
Li2SO4 CCl4
NI3 NSCl
NH4+ H3O+
NaCl ClO3-
Draw structural
and electron-dot diagrams for each.
BrO4- PO33-
IO4- NO3-
HCN SO32-
CO32- CaS
Na2SO4 NCl3
N2 O2
Cl2 C2H6
C2H4 C2H2
Draw
electron dot diagrams for each ionic compound
LiCl
Na2O
K2S
BaO
GaH3
Worksheet
# 10 Practice Test 1
1. Classify
as stable or reactive.
Na
N+ Ne Cl- S2- S3-
P P3- Ca
Ca+2 NaCl N3-
2. Describe
a metal and a nonmetal in terms of gaining or losing electrons.
3. Why
are noble gases always stable?
4. Determine
the number of valence electrons for:
Ca Ca2+ Cl Cl-
O O2- Al
5. Draw
a Bohr diagram for
Ca Ca2+
6. Determine the number of protons, neutrons,
and electrons in each.
Protons Neutrons Electrons
S
S2-
Al
Al3+
42Ca
7. Classify
as ionic or covalent compounds.
HCl CH3OH H2O NH4OH
NaCl MgSO4 CoCl2 H3PO4
NH3 P2O5
Ba(OH)2
8. Classify
the above compounds into
Acids Covalent Non-acids Salts Bases.
9. Calculate
the average atomic mass for magnesium using the following percentage abundance data.
24Mg
78.70% (24.00 amu)
25Mg
10.13% (25.00 amu)
26Mg
11.17% (26.00 amu)
10. Write
the formulas for each ionic compound.
Magnesium chloride
Silver phosphate
Cobalt III oxide
Zinc phosphate
Calcium nitride
Copper I nitrate.
11. Name
each ionic compound.
Fe2(SO4)3 CoCl3
Na2O AgCl
Na3PO4 CaF2
NH4OH Ca(NO3)2
K2Cr2O7 MgCrO4
12. Name each covalent compound.
P2O5 N2O3
CO CO2
SO2 P3O5
C6H6 H2SO4(l)
HCl(l) HNO3(l)
H2CrO4(aq) HF(aq)
H3PO4(l) H2CO3(aq)
13. Complete the chart below.
Protons Neutrons Electrons Reactive
# of valence
or stable? Electrons
Li+
S2-
22Mg2+
14. Write
dissociation equations showing how each salt or acid dissolves in water and forms ions. The first one is done for
you.
MgCl2 → Mg2+ +
2Cl-
FeCl3
Co2(SO4)3
HClO3
H2Cr2O7
16. Name
each acid:
HCl (aq) HClO3 (aq)
HNO3 (aq) HBr(aq)
17. Classify
each as formula units or molecules.
NaF H2O
CH4 NH4Cl
CaSO4 CH3OH
H2SO4
18. Indicate
the solutions that conduct electricity.
NaCl(s) NaCl(aq)
Ca(OH)2(aq) HCl(aq)
NH4OH(aq) C6H12O6(aq)
19. Consider the following electron
configuration 1s22s22p6. Determine the element
and some ions that have the above
electron configuration.
Element
Cation
Anion
20. Describe
why NaCl(s) doesn’t conduct
electricity. Describe what happens to NaCl(s) when
it is dissolved in water. Why does it
conduct electricity?
21. Write
the quantum electron configuration for each (1s22s2p.....)
He Ar
Na Na+
Cl Cl-
K K+
Br Br-
22. What two particles make up most of the
mass within an atom?
23. I
am an atom with 35p 40n. Who am I?
24. I
am a cation with 56p 81n
& 54e. Who am I?
25. Define
isotopes.
26. In
27. Define
ionic and covalent bonding
28. How
many valence electrons are in the calcium ion?
29. How
many valence electrons are in the fluoride ion?
30. What
is the name of the family that has and electron configuration of:
a) s2p5
b) s1
c) s2p2
31. Consider the following electron
configuration 1s22s22s22p63s23p6
. Determine the element
and some ions that have the above electron configuration. These are called isoelectronic.
Element:
Cations:
Anions:
Draw electron
dot diagrams for the following using brackets for ions. Write a dissociation
equation first.
32. NaCl
33. Li2O
34. CaF2
Name and
classify each compound as an acid, molecular, salt, or base.
35. CuSO4(aq)
36. P2O4(s)
37. H2SO4(aq)
38. H2Cr2O7(aq)
39. H2Cr2O7(l)
40. Ca(OH)2
. 5H2O
41. HBr(aq)
42. Calculate the average atomic mass for neon
if there are three naturally occurring isotopes
and they are:
20Ne
mass =
19.9924404 amu
abundance = 90.92 %
21Ne
mass =
20.993849 amu
abundance = 0.2570
%
22Ne
mass =
21.991385 amu
abundance = 8.820
%.
Show some work if you want some
marks. Round to an appropriate number of sig figs.
43. Name each compound below. Assume all
compounds are aqueous.
Ca(OH)2 NH4OH
CH3OH C12H22O11
HCl PI3
K2SO4 RbOH
H3PO4 NaOH
CaCl2 Li2SO4
SiO2 BaF2
BCl5 CH3COOH
H2CO3 CsOH
S2Cl2 Fr2S
Fe2(SO4)3 ZnCl2
Co3(PO4)2 Ag2Cr2O7
Worksheet # 11 Practice
Test 2
Balance each equation.
1. ___C16H34 + ___O2 → ___CO2 + ___H2O
2. ___Ga + ____H2SO4 → ____H2 +
____Ga2(SO4)3
Write a balanced equation including phase symbols.
3. Solid carbon
reacts with chlorine gas to produce liquid tetracarbon
decachloride.
Write chemical formulas for each ionic or molecular
compound.
4. Strontium sulphide
5. triphosphorous hexoxide
6. Osmium IV sulphide
Name each chemical formula
7. Sn(CO3)2.5H2O
8. Si3F8
9. NaHCO3
Classify the following as acids, bases, salts, and molecular
non-acids. Name each.
10. Sn(SO4)2
11. Ca(OH)2
12. CH3COOH
13. S2O5
Round off each measured number to three significant figures.
14. 0.0056349
15. 539663
Add or subtract the measured quantities.
16. 153.267 + 0.53493
17 4.5631 x 1024 +
2.36 x 1023
Simplify the following rounding to the correct number of
significant figures.
18. (5.6 x 10 -24) (5.37894 x 10-25)(6.532 x 1015)
(2.059378 x
1024)(5.23
x 1022)(9.37894 x 10-13)
Use unit analysis and the conversion factors to perform the
following conversions:
2.210 lb =
1.000 kg 14 lb = 1 stone (defined)
2000 lb = 1 ton (defined) 1.61 km = 1.00 mile
4.54 L =
1.00 gallon 16
oz =
1 lb (defined)
19. 236 oz to
stone
20. 8.53 stone to
oz
21. 25.6 mi/h to
km/s
State the model of the atom is best described by each
statement below.
22. First model to account for The Law of
Conservation of Mass.
23. The first
theory to explain the emission of photons.
24. First model to account for positive and
negative charges.
25. First model to account for the wave
properties of electrons.
26. First model to include a small dense
nucleus.
27. Describes the
atom as a small dense nucleus surrounded with electrons, which are not in orbitals.
28. Describes the
atom as a small dense nucleus surrounded with electrons, which are in spherical orbitals.
29. Describes the
atom as a spherical atom that is indestructible and combines in simple whole number ratios to form
compounds.
30. Describes
electrons as being contained in a complex 3D orbitals
as negative clouds of vibrational energy
31. Non-scientific
theory that delayed modern theories of the atom for 1800 years and was shown to be incorrect.
What did the evidence tell us about
the nature of the atom?
32. 99 % of
alphas in the gold foil experiment were completely un-deflected.
33. 1 % of alphas
in the gold foil experiment were radically deflected.
34. Flame
spectroscopy of an element produces an emission spectrum consisting of 4 photons.
35. A beam of
negative particles is produced in a Crooke’s tube.
36. There are
five naturally occurring isotopes of Germanium. Complete the chart to show the number of protons neutrons
and electrons.
protons neutrons electrons At. Mass
Abundance
70Ge 69.92428 20.52%
72Ge 71.92174 27.43%
73Ge 72.9234 7.760%
74Ge 73.92115 36.54%
76Ge 75.9214 7.760%
37. Calculate the
average atomic mass of Germanium. Show some work if you want some marks. Round to an appropriate
number of significant figs.
38. Write the
quantum electron configurations for the following atoms or ions.
39. F
40. Ga
41. Br-
42. Rb+
Determine the element that has the following electron
configuration.
43. 1s22s22p63s2
Determine a cation that has the
following electron configuration.
44. 1s22s22p63s23p6
Determine an anion that has the following electron
configuration.
45. 1s22s22p63s23p6
Complete the following chart.
Symbol p e n valance
el. stable/unstable atom/cation/anion
46. Na
47. P-3
48. Xe
49. Sr+2
Name and classify each compound as an acid, molecular, salt,
or base.
50. CuSO4(aq)
51. P2O4(s)
52. H2SO4(aq)
53. H2CO3(aq)
54. H2CO3(l)
55. Ba(OH)2.
2H2O
56. HF(aq)
Write structural diagrams and electron-dot diagrams for
each.
57. CCl4
58. S2
59. NH3
60. CO2
Write electron-dot diagrams for each.
61. NaCl
62. SO42-
63. Na3PO4
64. NO3-
65. NH4+
66. IO4-