Download t2 – modern atomic theory continued

CH 11
T2 – MODERN ATOMIC THEORY CONTINUED
1
You have mastered this topic when you can:
1) name and describe the atomic models developed by THOMSON, NAGAOKA, RUTHERFORD and BOHR.
2) identify the ATOMIC NUMBER of an ELEMENT and an ELEMENT from an ATOMIC NUMBER using the PERIODIC
TABLE.
3) define or describe NUCLEUS, LINE SPECTRUM, SUBATOMIC PARTICLE, PROTON, ELECTRON, TRANSITION,
GROUND STATE and PRINCIPLE QUANTUM NUMBER.
4) describe the relative location, mass and charge of a PROTON and an ELECTRON.
EARLY ATOMIC MODELS CONTINUED [Figure 1, pg. 37; 1.4, pgs. 37 – 45]
I) In 1870, William Crookes used cathode ray tubes to investigate atomic structure. A cathode ray tube is a sealed
glass tube containing trace amounts of gaseous element. Two electrodes, an anode and a cathode, are inserted
into opposite ends of the cathode ray tube and are connected to a power source, a source of electricity. When the
power source is activated the cathode becomes negatively charged and the anode becomes positively charged.
The cathode is the
negative terminal
The anode is the
positive terminal
Power source
A) When the power source was activated a glowing beam traveling between the cathode and the anode was
observed. This beam came to be known as cathode rays. Crookes knew that the cathode rays were emitted by
the cathode, and that they consisted of either light rays or glowing particles. He also knew that magnets have
no effect on light rays while they cause particles to move. This knowledge led him to place a magnet near a
beam of cathode rays to see if they moved. The magnet caused the cathode rays to bend, which led Crookes
to conclude that cathode rays were composed of particles not light rays. To remind us of this fact we’ll refer
to the beam as cathode particles. Crookes repeated these experiments using many different elements as the
cathode with each experiment producing the same results: All the elements tested produced a beam of cathode
particles that was deflected when a magnet was brought near it. Crookes was not able to determine how the
atoms of the elements used as a cathode produced the beam of cathode particles; this inspired further research
to identify the source of the cathode particles.
II) J. J. THOMSON: In the 1890’s, Thomson repeated the experiments done by Crookes 20 years earlier. Thomson
observed that every element used as a cathode produced a beam of cathode particles and that all beams of
cathode particles behaved the exact same ways. FIRST: Cathode particles traveled in a straight line from the
cathode, the negative terminal, to the anode, the positive terminal. SECOND: When he placed a positive charge
next to the cathode ray tube, the cathode particle beam was deflected (bent) toward it a specific distance. THIRD:
When he placed a negative charge next to the cathode ray tube, the cathode particle beam was deflected (bent)
away from it a specific distance. Thomson knew that opposite charges attract each other and like charges repel
each other. As a result he reasoned as follows: FIRST: Since the cathode particle beam was attracted to a positive
charge and repelled from a negative charge he concluded that cathode particles carry a negative charge. Since all
elements he tested emitted a cathode particle beam and all elements are composed of atoms, all atoms contain
these negatively charged cathode particles. SECOND: Every element tested produced cathode particles that
deflected (bent) the same distance toward a given positive charge and the same distance away from a given
negative charge placed next to the cathode ray tube. This meant that the negatively charged cathode particles
released from every element tested had identical mass and identical charge. These negatively charged cathode
particles were eventually named ELECTRONS, symbolized as e−, and were the first SUBATOMIC particle to be
discovered. SUBATOMIC means under or smaller than the atom. Since all electrons are identical and each
carries the same amount of negative charge, it was decided that each electron carries a charge of negative one
( −1) . This means that an atom containing four electrons carries a negative four charge.
A) Thomson knew that all atoms contained negatively charged electrons and that all atoms are electrically
neutral. This led him to develop the MUFFIN MODEL (plum pudding model) in 1903. Thomson’s Muffin
Model [Figure 4a, pg. 25] is summarized at the top of the next page.
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T2 – MODERN ATOMIC THEORY CONTINUED
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1. The atom consists of a neutrally charged sphere composed of a positively charged spongy matrix that is
mostly empty space.
2. All atoms contain negatively charged subatomic particles called electrons, which are suspended within
the positively charged spongy matrix.
3. Since all atoms are electrically neutral, the positive charge of the atom is equal to its number of electrons
(negative charge). i.e. If an atom contains seven electrons its spongy matrix supplies a positive seven
charge.
4. Thomson called his atomic model the Muffin Model [Figure 4a, pg. 25]. A chocolate chip muffin is a
fairly accurate representation of Thomson’s muffin model because it is composed of chocolate chips, the
electrons (e−), suspended in the cake batter containing many ‘empty’ spaces, the spongy matrix.
III) HANTARO NAGAOKA developed the SATURN MODEL [Figure 4b, pg. 25] in 1904 using the same information
Thomson used to develop his Muffin Model.
A) Nagaoka’s Saturn Model is summarized here.
1. All atoms contain negatively charged electrons.
2. The atom consists of a relatively large positively charged central sphere surrounded by a ring of
negatively charged electrons much like the planet Saturn and its rings thus he called it the Saturn
Model. [Figure 4b, pg. 25]
3. Atoms are electrically neutral which means they contain a positive charge that is equal to number of
electrons (negative charge) within the atom.
IV) ERNEST RUTHERFORD
A) The models developed by both Nagaoka and Thomson explained the existing evidence equally well, thus both
models were equally valid. Ernest Rutherford’s studies of radioactive decay in the early 1900’s led him to
favour one over the other. His research led him to modify it creating his own model.
B) Radioactive decay results when an unstable radioactive atom breaks apart emitting radiation in the form of
particles and energy. Rutherford conducted many experiments investigating radiation using the apparatus
illustrated below. In it, the radiation released by a radioactive source was directed toward a focus plate. The
focus plate concentrated the radiation into a narrow beam, which was passed between positive and negative
terminals. When the focused beam of radiation passed between the positive and negative terminals it was split
into three distinct beams.
Positive terminal
ß-rays (Beta)
γ-rays (Gamma)
α-rays (Alpha)
Radioactive
source.
Negative terminal
Focus Plate
Photographic Film
1) Rutherford’s experiment revealed that radiation consists of three distinct rays.
a) Since α-rays were deflected toward the negative electrode he concluded they are positive.
b) Since ß-rays were deflected toward the positive electrode he concluded they are negative.
c) Since γ-rays did not bend toward either electrode he concluded they are neutral.
d) Further research revealed that α-rays and ß-rays are particles while γ-rays are energy waves. αparticles are small, high-energy positively charged particles. ß-particles are electrons, which are
negatively charged subatomic particles.
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B) Rutherford’s more famous experiment conducted in 1911 shot α-particles at a piece of gold foil several atoms
thick [Figure 5, pg. 25]. Based on Thomson’s muffin model, he predicted all the α-particles would pass
through the gold foil experiencing minor deflections of course. This hypothesis seemed reasonable because
according to Thomson’s muffin model, the accepted model of the day, atoms are spheres of mostly empty
space in which negatively charged electrons are suspended in a positive spongy matrix. The results of
Rutherford’s experiments, however, were very different. Most of the α-particles passed through the gold foil
unaffected while a significant number of them deflected off course at large angles with a stunning few being
deflected back toward the source of the α-particles [Figure 5, pg. 25]. How you would explain these results?
1) Rutherford explained the results by concluding the following [Figure 6, pg. 26]: FIRST: Because most of
the α-particles passed through the foil unaffected, the atom consists of mostly empty space. SECOND:
Since α-particles are positive and a significant number were deflected off course at large angles with a
small but significant few deflecting back toward their source, the atom’s positive charge must be isolated
in an extremely small core. These conclusions suggested an initial model where the neutrally charged
atom consists of essentially empty space with its positive charge concentrated in an extremely small core
he called the NUCLEUS around which electrons travel in circular ring shaped orbits like planets orbiting the
sun [Figure 6, pg. 26; Figure 1, pg. 38]. In 1914 further research led Rutherford to conclude that the
nucleus consisted of positively charged subatomic particles he called PROTONS, symbolized as p+.
Research revealed that a proton’s mass is 1.673 x 10−24 g and its charge is +1.602 x 10−19 C (C = Coulomb
which is the standard unit of electric charge). Review Nagaoka’s and Thomson’s models described above
then, based on Rutherford’s results, predict which model was more accurate.
C) RUTHERFORD’S NUCLEAR MODEL (1914). Rutherford’s model of the atom is called the NUCLEAR MODEL
[Figure 6, pg. 26] and it is summarized here.
1. All atoms contain positively charged subatomic particles called protons and negatively charged
subatomic particles called electrons.
2. All atoms are electrically neutral, thus they contain equal numbers of protons and electrons.
3. Atoms consist of mostly empty space with protons located in an extremely small nucleus and electrons
traveling in randomly oriented circular orbits like bees circling their hive.
4. The nucleus contains approximately 99.95 % of the atom’s mass and approximately 0.05 % of its
volume while the orbits contain approximately 99.95% of the atom’s volume and approximately 0.05
% of its mass.
e−
Electrons reside in randomly
oriented circular orbits.
2p+
Protons in very small nucleus
e−
V) HENREY MOSELEY discovered an important property of the nucleus in 1913. His experiments revealed that the
atoms of each element have a unique amount of positive nuclear charge. Since no two elements have the same
positive nuclear charge, Moseley knew that the positive nuclear charge could be used to identify elements. As a
result, he created the term ATOMIC NUMBER = Z. The atomic number of an element is equal to the positive
nuclear charge of the atom. In 1914 Rutherford concluded that the atom’s positive nuclear charge was housed
in subatomic particles he called protons (p+). He theorized that each proton, (p+) contains a single positive
charge. Since atoms have a unique amount of positive nuclear charge, they must also have a unique number of
protons. This means the atomic number of an element is equal to the number of protons (p+) found in the
nucleus of its atoms.
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A) The atomic number, which equals the number of protons, is found at the top left corner of each box on your
periodic table. It is used to identify an element because it is exclusive to the element: i.e. The atomic number
of oxygen is Z = 8, therefore all oxygen atoms have exactly 8 protons, not 7 nor 9.
B) Required Practice 1: Answer these questions on your own paper. {Answers are on page 7 these of notes.}
1. Complete this table.
a
b
c
d
e
Z
3
Element
Sulphur
21
Kr
74
2. Name the element that has:
a. Z = 4
b. Z =13
c. Z = 24
d. Z = 40
e. Z = 84
3. Name, describe and draw the atomic models developed by Thomson, Nagaoka and Rutherford.
4. How are the models of Thomson, Nagaoka and Rutherford similar and how are they different?
5. Draw Rutherford models for elements having Z = 1 to 5.
NEILS BOHR [Figure 1, pg. 37; 1.4, pgs. 37 – 45]
I) NEILS BOHR knew that when elements are energized by heat or electricity they release flashes of light energy.
When the flash of light released from an energized atom is passed through a prism it is broken apart producing a
pattern of different coloured lines known as a LINE SPECTRUM. [Figure 6, pg. 40]
A) Study the line spectra of the 10 elements given in Appendix C6 on page 637 of your text. NOTICE that each
element produces is own unique line spectrum. The combination of colours in hydrogen’s line spectrum is as
unique to hydrogen as your fingerprints are to you; no other element has one like it. This means that elements
can be identified by their line spectrum.
B) Bohr’s (~1913) experimental analysis of hydrogen explained the LINE SPECTRA of elements and modified
Rutherford’s Nuclear Model. Rutherford’s Nuclear Model stated that each atom has positively charged
protons (p+) residing in a tiny central nucleus with negatively charged electrons (e–) traveling around the
nucleus in randomly oriented circular orbits like bees circling their hive (see the diagram at the bottom of
page 3 of these notes). Bohr realized there were problems with this model. He knew that oppositely charged
particles attract each other so he wondered why the negatively charged electrons (e–) didn’t fall into the
positively charged nucleus. The answer to Bohr’s question came as a result of his study of the relationship
between elements and light [pgs. 37 – 40]. Bohr knew that energized atoms release light, a form of energy,
and that electrons carry energy. As a result he reasoned that electrons are involved producing the flash of
light that creates an element’s line spectrum. The element with the simplest atomic structure is hydrogen. Its
line spectrum is the ninth from the top in Appendix C6 on page 637 of your text. Bohr’s research led him to
develop a theory that explained hydrogen’s line spectrum. In doing so, he modified Rutherford’s Nuclear
Model and sparked the creation of one of the most profound theories in chemistry and physics – QUANTUM
MECHANICS.
II) EXPLAINING HYDROGEN’S LINE SPECTRUM
A) Recall that Rutherford’s Nuclear Model of hydrogen has one proton in the nucleus and one electron traveling
it in a randomly oriented circular orbit as illustrated below.
p+
e−
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B) When an hydrogen atom is energized by heat or electricity, it emits a flash of light. To explain this flash of
light, Bohr reasoned that hydrogen’s single electrons (e−) absorbs the heat or electrical energy and becomes
excited. Since electrons, like all matter, have a natural tendency to exist with the lowest amount of energy
possible, the excited high-energy electron will release the energy it gained as a flash of light.
1) Passing the flash of light through a prism produces the four coloured line spectrum that is unique to
hydrogen. [See the ninth line spectrum on page 637.] Bohr’s reasoning explained the source of the light
flashes from energized hydrogen but not the different colours of its line spectrum. His curiosity continued:
How does hydrogen’s single excited high-energy electron create four different coloured lines?
a) Bohr knew that electrons (e−) are energy carriers and that light travels in waves. [Figure 5, pg. 39:
NOTICE the four coloured waves released by energized hydrogen.] Waves have peaks and valleys. The
distance from one peak to the next is called a wavelength. High-energy light waves have a short
wavelength while low-energy light waves have a long wavelength. Light consists of many colours,
some of which are seen in a rainbow: red, orange, yellow, green, blue, indigo and violet. Each colour of
light has a unique fixed wavelength, which means it carries a unique fixed amount of energy.
Hydrogen’s line spectrum consists of four different coloured lines: violet, indigo, blue-green and red.
[Figure 5, pg. 39: NOTE: The light labeled green in the text should be labeled blue-green.] Carefully
study the four coloured waves given in Figure 5 on page 39 of your text. NOTICE that each colour of
light has a different wavelength. When arranged from shortest to longest wavelength, the four colours
are ordered: violet, indigo, blue-green and red. Since light waves with short wavelengths are highenergy waves while light waves with long wavelengths are low-energy waves, violet light carries more
energy than indigo light, which carries more energy than blue-green light, which carries more energy
than red light.
C) Bohr explained hydrogen’s line spectrum by modifying Rutherford’s Nuclear Model as outlined here. The
different colours of light found in hydrogen’s line spectrum are caused by its single electron first absorbing
then emitting different quantities of energy. Bohr theorized that an electron (e−) is an energy carrier capable
of carrying different fixed quantities of energy. Since an electron carries different fixed quantities of energy,
its energy is said to be QUANTIZED. The quantity of energy carried by an electron determines its distance
from the nucleus. Since an electron can carry different fixed quantities of energy, it can reside at different
fixed distances from the nucleus, each distance being determined by the quantity of energy the electron
carries. Bohr called those distances ENERGY LEVELS, although chemists and physicists continue to call them
orbits. Each hydrogen atom has an infinite number of energy levels (orbits). The quantity of energy each
orbit is capable of carrying increases as its distance from the nucleus increases, this means low-energy orbits
are close to the nucleus while high-energy orbits are far from the nucleus. Since each orbit holds a unique
fixed quantity of energy, each orbit is given a unique designation called the PRINCIPLE QUANTUM NUMBER,
symbolized n [Figure 3, pg. 38]. The first orbit is designated n = 1, the second is n = 2, the third is n = 3, etc.
The larger the principle quantum number, n, the greater the energy carried by the orbit. The n = 1 is the
lowest energy orbit; n = 2 is a higher energy orbit and n = 3 is an even higher energy orbit. When an
electron, e-, is in the lowest-energy orbit possible, it is said to be in its GROUND STATE. An electron cannot
exist outside an orbit, thus when it gains or emits energy it must ‘jump’ between orbits. The process of
‘jumping’ between orbits is called TRANSITION [Figure 4a, pg. 38].
1) IN SUMMARY: When an electron (e−) in its ground state is energized, it becomes excited and must
transition to a higher-energy orbit [Figure 4a, pg. 38]. The more energy the electron gains the farther
from the nucleus it will transition. Since electrons have a natural tendency to exist carrying the lowest
amount of energy possible, they will transition back to the lowest-energy orbit available to them as soon as
possible. In order to transition back to the lowest-energy orbit available, the excited high-energy electron
must release the energy it absorbed as a flash of light [Figure 4b, pg. 38]. Transition between different
high energy orbits and low energy orbits releases different amounts of energy which appears as different
colours of light. Since each colour represents a different wavelength of light and thus a different amount of
energy, the colour of each line of the line spectrum indicates the energy difference between the two
orbits involved in the electron’s transition. This information was used to confirm that each individual
orbit is capable of carrying a different quantized (fixed) amount of energy (the difference in energy
between the orbits involved in a transition) and the amount of energy each orbit is capable of carrying.
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−
2) Bohr used this theory to explain hydrogen’s line spectrum. Hydrogen’s single electron (e ) exists in its
ground state, the first orbit, n = 1. When an hydrogen atom is energized by subjecting it to heat or
electrical energy, its single electron (e−) absorbs some of the energy and becomes excited causing it to
transition (‘jump’) from its ground state, the n = 1 orbit, to either the n = 3, n = 4, n = 5 or n = 6 orbit.
The greater the quantity of energy it absorbs the farther from the nucleus is its transition. Since every
electron has a natural tendency to exist in the lowest possible state of energy, Hydrogen’s excited highenergy electron will emit the energy it absorbed by transitioning (‘jumping’) back to its ground state, the
n = 1 orbit. Transition of hydrogen’s electron from the four different high-energy orbits it can occupy to
the lowest-energy orbit available as it moves to its ground state orbit causes it to release different amounts
of energy in the form of flashes of different colours of light.
a) Study Figure 5 on page 39 of your text. The colours produced are explained here.
Transition from the n = 6 orbit to the n = 2 orbit causes the electron to emit a flash of violet light.
Transition from the n = 5 orbit to the n = 2 orbit causes the electron to emit a flash of indigo light.
Transition from the n = 4 orbit to the n = 2 orbit causes the electron to emit a flash of blue-green light.
Transition from the n = 3 orbit to the n = 2 orbit causes the electron to emit a flash of red light.
Transition from the n = 2 orbit to the n = 1 orbit causes the electron to emit a flash of infrared
radiation. We cannot see infrared radiation thus it does not appear in hydrogen’s line spectrum.
III) THE BOHR MODEL, AN INTRODUCTION TO QUANTUM MECHANICS
A) Bohr used the knowledge he acquired from his analysis of hydrogen’s line spectrum to modify Rutherford’s
Nuclear Model. Bohr’s atomic model, often called the SOLAR SYSTEM MODEL, is summarized here with the
diagram given at the top of the next page.
1. Atoms are composed of equal numbers of positively charged protons and negatively charged
electrons.
2. Protons reside in the tiny nucleus, which contains approximately 99.95% of the atom’s mass and
approximately 0.05 % of its space.
3. Electrons travel around the nucleus within an infinite number of concentric circular energy levels,
each identified by a principle quantum number, n [Figure 3, pg. 38]. Energy levels contain
approximately 0.05% of an atom’s mass and approximately 99.95% of its space.
4. Electrons cannot exist outside an energy level, thus they transition (jump) between energy levels as
they gain or lose fixed quantities of energy.
5. Electrons carry fixed quantities of energy that determines the distance they are from the nucleus and
thus which orbit they occupy. Low-energy electrons reside in inner low-energy orbits while highenergy electrons reside in outer high-energy orbits.
Protons in the tiny nucleus.
Electrons in orbits.
Electrons in orbits.
e
Atoms have an infinite
−
number of un-occupied
2p
+
circular orbits in which they
can exist.
e
−
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B) Required Practice 2: Using correct terminology, answer these questions on your own paper. {Answers are
on page 8 of these notes.}
1. Explain how chemists find line spectra of elements useful.
2. Define the terms ground state and transition.
3. Describe what happens to an electron when it is energized and when it releases energy.
4. What does the principle quantum number refer to?
5. Which orbit has the least energy: n = 2, n = 3, n = 4?
6. What do different colours of an element’s line spectrum tell us?
7. How are the models of Rutherford and Bohr similar?
8. How are the models of Rutherford and Bohr different?
C) Read section titled “Applications of Emission Spectra: Flame Tests” on page 42 of your text. Flame Tests
for four different metals are shown in Figure 8 page 44. Since each flame test produces a different colour,
each picture shows the results of burning a different element. Why do different elements produce different
coloured flame?
1) Each element produces a different coloured flame because the atoms of each element have a unique
combination of electrons, which occupy orbits differently. Electrons transition between orbits as they
absorb or emit energy. Because the atoms of each element have a unique number of electrons occupying
different orbits, different electrons within the atoms of each element will transition between different
energy levels. In other words, within the atoms of each element a unique combination of electrons is
transitioning between a unique combination of orbits. No two elements have the same combination of
electrons transitioning between the same combination of orbits thus providing each element an individual
coloured flame when burned.
ANSWERS TO THE REQUIRED PRACTICE
Required Practice 1 from page 4
1. Answers in the table are in bold.
a
b
c
d
e
Z
3
16
21
36
74
Element
Lithium
Sulphur
Scandium
Kr
Tungston
2a. Be; 2b. Al; 2c. Cr; 2d. Zr; 2e. Po. 3. Thomson: Muffin Model: Atoms consists of a neutrally charged
sphere composed of a positively charged spongy matrix that is mostly empty space in which negatively charged
subatomic particles called electrons which are suspended. Since all atoms are electrically neutral, the positive charge
of the atom is equal to its number of electrons. Nagaoka: Saturn Model: All atoms contain negatively charged
subatomic particles called electrons located in a ring that circles a large positively charged central core. All atoms
are electrically neutral therefore their positive charge is equal to number of electrons. Rutherford: Nuclear Model:
All atoms contain positively charged subatomic particles called protons which are housed in a central nucleus and
negatively charged subatomic particles called electrons which are housed in randomly oriented orbits that surround
the nucleus. The nucleus contains approximately 99.95 % of the atom’s mass and approximately 0.05 % of its
volume while the orbits contain approximately 99.95% of the atom’s volume and approximately 0.05 % of its mass.
See your teacher to check your drawings. 4. Similarities: atoms are electrically neutral, contain positive and
negative charge, and contain subatomic particles called electrons. The atoms are electrically neutral which means an
atoms positive charge is equal to its number of electrons. Differences: Thomson’s model has electrons suspended in
a spongy matrix, while Nagaoka’s and Rutherford’s models have electrons orbiting a central positive core.
Rutherford’s model has the positive charge housed in subatomic particles called protons, which are located in an
extremely tiny central nucleus. 5. See your teacher.
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Required Practice 2 from page 7
1. Line spectra can be used to identify the elements that a substance is composed of. 2. Ground state is the lowest
energy level that an electron can occupy. Transition is the jumping of electrons between orbits. 3. When an
electron is energized it transitions to a higher-energy orbit; when an electron loses energy it transitions to a lowerenergy orbit. 4. Principle quantum numbers refer to the different energy levels an atom contains. 5. The n = 2
orbit has the least energy. 6. The colours of each line of a line spectrum indicate the different wavelengths of light
that are released as electrons transition toward the ground state. The different wavelengths are an indication of the
difference in energy between the two energy levels involved in the transition toward an electrons ground state.
7. Similarities: Atoms are electrically neutral containing equal numbers of positive charge subatomic particles called
protons and negative charged subatomic particles called electrons. Protons are found in a tiny central nucleus with
electrons circling in orbits. The nucleus contains approximately 99.95 % of the atom’s mass and approximately 0.05
% of its volume while the orbits contain approximately 99.95% of the atom’s volume and approximately 0.05 % of
its mass. 8. Differences: Rutherford’s model has orbits containing electrons arranged randomly around the nucleus
like bees circling their hive. Bohr’s model has an infinite number of circular orbits called energy levels that an
electron can occupy. Low energy orbits are close to the nucleus while high energy orbits are far from the nucleus.
Electrons can carry fixed quantities of energy, which determines the orbit they occupy. Energized electrons will
transition (jump) to a higher energy orbit then transition release the energy as they transition back to a lower energy
level.
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ASSIGNMENT
At the top of your assignment, please print “T2 – Modern Atomic Theory Continued, your LAST then First
name, block and date. Show all your work for questions requiring calculations; marks will not be awarded for final
answers only. Complete these questions in the order given here. [Marks indicated in italicized brackets.]
A. Name, draw and use bullet form to describe in full detail the atomic models of Thomson, Nagaoka, Rutherford
and Bohr. [22]
B. Define or describe these terms. [4]
1. Transition
2. Ground state
3. Atomic number
4. Subatomic particles
C. Match the number of the statement with the letter of the correct model. Correct model letters may be used more
than once. [4]
1. That atom contains a positive spongy matrix.
A. Saturn model
2. The atom contains a tiny nucleus inside layered orbits.
B. Solar system model
3. Electrons can transition between orbits.
C. Muffin model
4. The atom contains electrons embedded within the positive charge.
D. Dalton’s model
5. The atom contains electrons in randomly oriented orbits.
E. Nuclear model
6. Atoms are solid.
7. The atom contains a large positive core.
8. Electrons prefer to exist in their ground state.
D. Page 7: Are You Ready #15. [2]
E. Page 42: Practice #1, 2 & 3. Be sure you use correct terminology. [4]
F. Page 45: Practice #5, 6, & 7. [4]
G. Name these elements. [2]
1. Z = 9
2. Z =21
3. Z = 30
4. Z = 47
[42 marks in total]
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