Download Development of the Atomic Theory

Atomic Theory
Unit 3 Development of the Atomic Theory
1.
2.
3.
4.
5.
Where is the mass of the atom concentrated? In the nucleus
What is located in the nucleus? Neutrons and protons
What is the negative particle that orbits the nucleus? electrons
What is the sum of the protons and neutrons called? Mass number
With a neutral atom, what two items are equal in number? Protons and
electrons
6. What is the term for any charged particle? ion
7. What is the term for the positively charged ion? cation
8. What is the term for the negatively charged ion? anion
9. What type of ion is formed when electrons are gained? anion
10. What type of ion is formed when electrons are lost? cation
11. How does atomic theory today differ from Dalton’s theory? Today we
know about ions and isotopes
12. Which model of the atom is based on the solution to the Schrodinger
equation? Quantum mechanical model
How is this different from the planetary model? Electrons are not in fixed
orbitals
Match each term from the experiments of J.J. Thomson with the correct
description.
13. anode d
a. an electrode with a negative charge
14. cathode a
b. a glowing beam between electrodes
15. cathode ray b
c. an electrode with a positive charge
16. electron d
d. a negatively charged particle
17. The diagram shows electrons moving from left to right in a cathode-ray tube. Draw
an arrow showing how the path of the electrons will be affected by the placement of
the negatively and positively charged electrodes.
H. Cannon, C. Clapper and T. Guillot
Klein High School
Atomic Structure
18. Indicate the letter of each sentence that is true about atoms, matter and electric
charge.
a. All atoms have an electric charge.
b. Electric charges are carried by particles of matter.
c. Atoms always lose or gain charges in whole-number multiples of a single
basic unit.
d. When a given number of positively charged particles combine with an equal
number of negatively charged particles, an electrically neutral particle is
formed.
19. Indicate the letter next to the number of units of positive charge that remain if a
hydrogen atom loses an electron.
a. 0 b. 1
c. 2
d. 3
20. The positively charged subatomic particle that remains when a hydrogen atom loses
an electron is called a(n) _proton____.
21. What charge does a neutron carry? neutral
22. Indicate the letter of each sentence that is true about the nuclear theory of atoms
suggested by Rutherford’s experimental results.
a. An atom is mostly empty space.
b. All the positive charge of an atom is concentrated in a small central region
called the nucleus.
c. The nucleus is composed of protons and neutrons.
d. The nucleus is large compared with the atom as a whole.
e. Nearly all the mass of an atom is in its nucleus.
23. According to Bohr, electrons cannot reside at __
in the figure below
a. point A
b. point b
c. point c
d. point d
24. According to the quantum mechanical model, point D in the above figure represents
(a) the fixed position of an electron
(b) the farthest position from the nucleus that an electron can be found
(c)A position where an electron probably exists
(d) a position where an electron cannot exist
25. What does the atomic number of an atom represent? Number of protons
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Unit 4 Electrons in the atom
Select the best possible response.
1.The type of charge on the nucleus is
a. negative
b. neutral
c. positive
2. The number of protons in a neutral atom having 18 electrons
is_18.
3. How many neutrons are in the isotope P-29? 14
A neutral atom of bromine has a mass number of 80.
4. It has_35_ protons.
5. It has_35_ electrons.
6. It has_45_ neutrons.
7. Its nuclear composition is_35 protons and 45 neutrons___ .
One isotope of hydrogen has a mass of 3.
8. How many protons does it have? 1
9. How many neutrons does it have? 2
An atom with a +1 charge has atomic number 19 and mass number 39.
10. It has_19__ protons.
11. It has 18___electrons.
The anion of oxygen has a -2 charge.
12. How many electrons does it have? 10
_____.
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Atomic Structure
An atom has a charge of +1 and has 10 neutrons with a mass number of 21.
14. What is its atomic number? 11
15. How many electrons does it have? 10
An atom has 10 protons, 8 neutrons, and 12 electrons.
16. What is its charge? 217. What is its mass number? 18
18. In forming this ion, the neutral atom gained what? 2 electrons
An atom has 15 protons, 16 neutrons, and 16 electrons.
19. What is its mass number? 31
20. What is the overall charge on this ion? 1A neutral atom has a mass number of 24 and 11 protons.
21. How many protons does it have? 11
22. How many neutrons does it have? 13
An atom with a -2 charge has atomic number 16 and mass number 32.
23. It has_16___ protons.
24. It has_18___ electrons.
25. It has__16__neutrons.
Calculate the following:
26. There are 4 isotopes of sulfur, S-32, 95.002%; S-33, 0.76%; S-34,
4.22%; and S-36, 0.014%. What is the average amu of sulfur?
32.1 amu
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Atomic Structure
27. There are three common isotopes of chromium, Cr-50, 4.345%
abundance; Cr-52, 83.789%; Cr-53, 9.50% abundance; what is the average
amu of chromium?
50.8 amu
28. Calculate the wavelength of the yellow light emitted by a sodium lamp if
the frequency of the radiation is 5.10 x 1014 Hz (5.10 x 1014 s-1).
588 nm
29. What frequency is radiation with a wavelength of 5.00 x 10-6 cm? In
what region of the electromagnetic spectrum is the radiation?
6.00 x 1015 Hz
30. What is the energy of a photon of green light with a frequency of 5.80 x
1014 s-1?
3.84 x 10-19
31. Suppose your favorite AM radio station broadcasts at a frequency of
1150 kHz. What is the wavelength in cm of the radiation from the station?
26100 cm
Atomic Structure Practice
Element/Isotope
1) Zinc
2) Aluminum
3) Calcium
4) Sulfur
5) Bromine
6) Gold
7) Silver
8) Platinum
9) Uranium – 236
10) Plutonium – 246
11) Potassium – 39
12) Mercury – 201
13) Titanium - 48
14) Titanium - 46
Atomic
number
30
13
20
16
35
79
47
78
92
94
19
80
22
22
Mass
number
65
27
40
32
80
197
108
195
236
246
39
201
48
46
Protons Neutrons Electrons
30
13
20
16
35
79
47
78
92
94
19
80
22
22
35
14
20
16
45
118
61
117
144
152
20
121
26
24
30
13
20
16
35
79
47
78
92
94
19
80
22
22
Atomic Structure Exercise
I. Use the periodic table to compute the number of electrons. Neutrons and
protons in the following:
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Atomic Structure
A. Cr
24p, 24e, 28n
B. Cl 17p,17e, 18n
C. Mg 12p, 12e, 12n
D. Ir 77p, 77e, 115n
E. Si 14p, 14e, 14n
F. Ne 10p, 10e, 10n
II. Moseley used x rays to determine the atomic numbers of the elements
identify each of the following elements by name.
A. 1 Proton H
B. 4 Protons Be
C. Protons
D. 12 Protons Mg
E. 20 Protons Ca
F. 30 Protons Zn
III. How many protons are in the nucleus of each of the following elements
A. Uranium
B. Selenium
C. Helium
D. Magnesium
92
34
2
12
IV. Give the number of neutrons in each of the following isotopes
A. Titanium-46 B. Nitrogen-15
24
8
C.
34
S 16
18
D. 85Cu29
56
The Spectra of Elements
Everybody knows that a few drops of soup or milk spilled onto a gas burner will
change the blue gas flame into a mixture of colors, predominantly yellow. These colors
can be used to identify the elements present in the substance dropped into the flame.
We will observe the colors produced by several known substances.
The color observed in the flame is the result of atoms of the element absorbing
energy from the flame then reemitting it. The energy absorbed has an energy and
wavelength in the visible region of the electromagnetic spectrum. What your senses
detect as light is actually radiation that is part of a continuum that makes up the
electromagnetic spectrum. The electromagnetic spectrum is a continuum of energies
ranging from the high-energy gamma and x-rays to low-energy radio and microwaves.
Radiation includes any energy emitted in all directions from a single source, not just
nuclear decay. Light is one form of energy that can be radiated; - the sun constantly
radiates energy into space as its matter is converted into energy. We see this light in a
variety of colors, depending on the wavelength of the energy when it reaches us. The
waves of light with the longest wavelength (red) are refracted (bent) to a lesser degree
than the shorter waves (indigo and violet) as they pass through the atmosphere. The
visible range of the spectrum is from about 400 nm to 700 nm. The wavelength () is the
distance between successive peaks of a wave. The number of waves that pass a given
point in space per second is the frequency (f). The unit is inverse time, sec -1 and is also
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called hertz (Hz). All light moves through a vacuum at the same rate of speed, about 3
X 108 m/s (the speed of light – c). Frequency and wavelength are inversely related to
each other, as the product of the two equals the speed of light: c = f. Since the electrons
of an atom can only absorb certain amounts of energy, the wavelength of the energy
emitted after excitement of the atoms is characteristic of the element. When viewed
through a spectroscope, the spectrum of energy emitted by a sample of a particular
element can be observed.
1. Use the spectroscope to view each of the gas tubes set up in the back
of the lab. Each colored band corresponds to a different amount of
energy (a different wavelength) of light emitted by an electron in the
element. Each element has a unique spectrum based on the amount of
energy emitted by its electrons
2. Record the details of the colored lines formed on the scale of your
spectroscope and the wavelength of each line. Using the equations we
have discussed in class, solve for the frequency and energy of each of the
gases observed. If a substance has multiple lined, calculate frequency and
energy for each line. Express each answer to three significant digits.
Express the energy in ergs.
In the next part of our investigation, we will use a procedure known as a flame test to
become familiar with the colors produced by several common elements and use these
observations to better understand the affect of energy on atoms and ions.
Procedure:
a.
b.
Light a Bunsen burner. Make sure that the gas burns with
a clear blue flame and has an easily distinguishable inner blue cone.
Obtain samples of ions as provided by your teacher.
These will be metal nitrates, dissolved in a solvent to allow easy
distribution of the ion.
c.
Using the spray bottles, introduce a small amount of each
ion (one squirt) into the flame and observe the result.
d.
After you have identified the color produced by each ion,
mix two of the substances and see if you can detect the presence of
each of the substances or if the colors mingle so that one cannot be
distinguished from the other.
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Atomic Structure
e.
Color Key:
Na = Yellow
green
Sr = Deep red
Obtain an unknown sample and observe the colors
produced. Try to identify the element present in the unknown sample.
K = Violet
Ca = Yellow-red
Cu = Blue-
Li = Crimson
Ba = Green-yellow
1. What is the order of the colors in the spectrum from lowest to
highest wavelength?
2. What is the order of metallic ions from lowest to highest
wavelength?
3. What correlations if any do you see between the electron
configuration of the metallic elements and the energy given off
by exciting their electrons
Further Analysis
In 1913, Neils Bohr hypothesized that the electron in the hydrogen atom is allowed to
have only certain amounts of energy. These energy levels would be orbits in which the
electron in hydrogen would circle the nucleus of hydrogen. He believed that the
electron would move from one energy level to another and would give off light when it
jumped from a higher energy level to a lower energy level. The amount of energy
would be different for jumps between different levels.
1) The diagram below shows a sketch of some of the possible orbits of the
hydrogen electron and their corresponding energy values. If we think of the
energy values as being on a number line, which orbit has the greatest value?
2) When an electron jumps from E2 to E1 the amount of energy in the light would
be E2 – E1; or (-3.4 eV) – (-13.6 eV) which would equal 10.2 eV.
What would be the amount of energy in the light when an electron jumps from:
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a) E3 to E4 .66 eV
b) E4 to E2 2.55 eV
c) E3 to E2 1.89 eV
d) E5 to E2 2.86 eV
3) Using Bohr’s model, we would assume that the electron would only move
between certain orbitals or energy levels. Every possible jump corresponds to
light of a different energy.
a. How many different energies of light can be emitted from hydrogen
when the electron jumps down to E2 from E3, E4, E5, and E6?
b. How many bands of light did you observe when you viewed the
hydrogen tube through the spectrometer?
c. How do you think these two observations are related?
4) The colors corresponding to jumps to the E1 level have higher energies or lower,
than those to the E2.
a. As the electron jumped from the E6 all the way to the E1 level, how
many different energies would be emitted?
b. Did you see all of these bands? Explain why or why not.
c. Bohr did not find bands corresponding to the jumps to the E3 level. Why
do you think that was?
Conclusions:
In this activity we saw the evidence that energy is emitted in particular patterns
depending on the distance that an electron travels between one orbital and another. The
number of electrons in an atom and the number of energy levels those electrons occupy
determine the number of bands of light and the color of the light seen.
What if any conclusions can you draw about the relative numbers of electrons in the
atoms of our test atoms?
Arrangement of Electrons in Atoms
1. What is the difference between the earlier models of the atom and the modern
quantum mechanical model? No fixed orbitals
2. What is a quanta and who developed the concept of quantum energy? Packet of
energy that it takes to move an electron from one energy level to the next; Planck
3. How many quantum numbers are used to describe the energy state of an electron
in atom?
a) 1
(b) 2
(c) 3
d) 4
4. What is the Heisenberg uncertainty principle?
It is not possible to predict both position and momentum at the same time
5. The energy level of an electron is the region around the nucleus where an electron
is likely to be found____
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6. In general, the higher the electron is on the energy ladder, the _further_ it is
to/from the nucleus?
7. A quantum of energy is the amount of energy required to
a. move and electron from its present energy level to the next lower one
b. maintain an electron in its present energy level
c. move an electron from its present energy level to the next higher one
8. True or False: the electrons in an atom can exist between energy levels. F
9. True or False: The quantum mechanical model of the atom estimates the
probability of finding an electron in a certain position. F
10. Which name describes the major energy levels of electrons?
a) atomic orbitals
b) quantum mechanical numbers
c) quanta
d) principal quantum number
11. What formula represents the maximum number of electrons that can occupy a
principal energy level (n = principal quantum number)?
a) 2n2
b)n2
c) 2n
d) n
12. A spherical electron cloud surrounding an atomic nucleus would best represent _.
a. an s orbital
b. a p orbital
c. a combination of two different p orbitals
d. a combination of an s and a p orbitals
13. An energy level of n = 4 can hold
(a) 32
(b) 24
14. An energy level of n = 2 can hold
(a) 32
(b) 24
electrons.
(c) 8
electrons.
(c) 8
(d) 6
15. An electron for which n = 4 has more
(a) spin
(c) energy
(b) stability
(d) wave nature
(d) 6
than an electron for which n = 2.
Behavior of electrons in the atom
1. How did de Broglie conclude that electrons have a wave nature?
2. Identify each of the four quantum numbers and the properties to which
they refer.
3. How did the Heisenberg uncertainty principle contribute to the idea that
electrons occupy "clouds," or "orbitals"?
4. Complete the following table.
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Atomic Structure
Principal quantum number,
n
1
2
3
4
Number of Sublevels
Types of Orbitals
1
2
3
4
S
S, p
S, p, d
S, p, d, f
5. The way in which electrons are arranged around the nuclei of atoms is called
.electron configuration
Match the name of the rule used to find the electron configurations of atoms with the
rule itself.
6. aufbau principle
a. When electrons occupy orbitals of equal
__b__
energy, one electron enters each orbital until
all the orbitals contain one electron with
parallel spins
__c__
7. Pauli exclusion
b. Electrons enter orbitals of lowest energy
principle
first
_a___
8. Hund’s rule
c. an atomic orbital may describe at most two
electrons moving in opposite directions
9. In the shorthand method for writing an electron configuration, what does the sum
of the superscripts equal? Number of electrons in the atom
Write the electron configuration and orbital notation for each of the following atoms.
10. Nitrogen 1s22s22p3
11. Aluminum 1s22s22p63s23p1
12. Argon 1s22s22p63s23p6
13. Which guideline, Hund's rule or the Pauli exclusion principle, is violated in the
following orbital diagrams?
Pauli
Hunds
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Atomic Structure
14. What is the relationship between the principal quantum number and the electron
configuration? Tells the number of energy levels
15. How does the figure above illustrate Hund's rule? One in each of the 2p levels
before two in any
16. How does the figure above illustrate the Pauli exclusion principle? If 2 in the
orbital they spin in opposite directions
17. True of False: The aufbau principle works for every element in the periodic table.
F
18. Filled energy sublevels are more _stable___ than partially filled sublevels.
19. Half-filled sublevels are not as stable as _filled____ levels but are more stable
than other configurations.
20. Write the electron configuration of the following atoms:
a. carbon 1s22s22p2
b. potassium 1s22s22p63s23p64s1
c. gallium 1s22s22p63s23p64s23d104p1
d. copper 1s22s22p63s23p64s23d9
21. What is an electron dot structure? Shorthand way to show the valence electrons
22. Draw the electron dot structure of each of the following atoms.
a. argon
..
: Ar :
..
b. calcium
c. iodine
Ca:
.
:I:
..
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Atomic Structure
23. Write the electron configurations for these metals and circle the electrons lost
when each metal forms a cation.
a. Mg 1s22s22p63s2
b. A1 1s22s22p63s22p1
c. K 1s22s22p63s23p64s1
Match the noble gas with its electron configuration.
_c__
__a_
_b__
_d__
24. argon
25. helium
26. neon
27. krypton
a.
b.
c.
d.
1s2
1s2 2s2 2p6
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p63d10 4s2 4p6
13. What is the electron configuration called that has 18 electrons in the outer energy
level and all of the orbitals filled?
Pseudo nobel gas
14. Write the electron configuration for zinc. 1s22s22p63s23p64s23d10
15. Fill in the electron configuration diagram for the copper(I) ion.
Electron Review
1. Write the electronic configurations, orbital notation, Lewis dot structure for the
following:
a. S
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Electron Configuration
b. C
c. P
d. Ca
e. Zn
f. Fe
g. A1
2. How many dots would appear in the Lewis electron dot diagram for an atom whose
electron configuration ended 4S2 3d10 4p3 ? 5
3. How many unpaired electrons does the Lewis dot structure of N have? 3
4. How many pairs of electrons does the Lewis dot structure of O have? 2
5. Why does copper have 1 valence electron? Half filled orbitals are more stable than
others
Charting Oxidation Number
Complete the following chart. You may wish to use the periodic table in your text.
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Electron Configuration
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