Download Learning Guide 11: Atomic models

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LEARNING GUIDE 11:
ATOMIC MODELS
BEFORE YOU START
Learning the atomic models involves appreciating the history of chemistry, understanding the importance of laws,
and helps students understand theories may be changed as new data are collected and discussed.
There are numerous models and you must memorize their different features. Students should be able to compare
and contrast these models and explain how they were developed.
STATE CONTENT STANDARDS
State Standard 1: The periodic table displays the elements in increasing atomic number and shows how
periodicity of the physical and chemical properties of the elements relates to atomic structure.
MAJOR CONCEPTS
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Lesson 37: Subatomic particles
Lesson 38: Dalton’s atomic theory
o Students learn the major atomic models
o Describe Democritus’ ideas about atoms
o Describe the size of an atom
Lesson 39: Describe J. J. Thomson’s model of the atom
Lesson 40: Rutherford’s nuclear model
Lesson 41: Bohr model: Introduction
Lesson 42: Wave model: Introduction
Lesson 43: Standard model
VOCABULARY
Define each vocabulary word.
Nuclear atom
Subatomic
Excited state
Nucleus
Photon
Ground state
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READING
Part I: Important Facts (List seven important facts found during the reading)
Section 3.3
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Section 3.5
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Section 11.1
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Part II: Summaries (3.6, 11.5, and 11.6 are not included here but you should read them.)
Using one paragraph per section, summarize the three sections you just read.
Section 3.3
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Section 3.5
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Section 11.1
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READING GUIDE
3.3: Title: ______________________________________________________
1. What does the law of constant composition mean?
2. Although Dalton’s atomic theory was rejected at first, he used his theory to do WHAT that soon
resulted in his theory being accepted by scientists?
3.5: Title: ______________________________________________________
3. Because of the success of Dalton’s theory, scientists came to believe what two things?
4. How did J.J. Thomson know that the particles he had discovered were negatively charged?
5. Today, the particles discovered by Thomson are called __________________.
6. Why did Thomson then conclude that atoms must also have positively charged particles?
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7. Describe J. J. Thomson’s plum pudding model of the atom.
8. Rutherford was a student of ______________________ and designed an experiment in which
___________________-charged alpha () particles were directed at a _______ _________
_______.
9. What did Rutherford expect the  particles to do?
10. What were the results of Rutherford’s experiment?
11. What must have caused the large deflections of the  particles?
12. Why did most of the  particles pass directly through the foil?
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Learning Guide: 11
13. Describe Rutherford’s nuclear atom.
14. The neutron was discovered in 1932 by Rutherford and _________________________.
CHAPTER 11
11.1: Title: ____________________________________________________
15. Rutherford’s experiments involved bombarded metal foil with _____ _____________.
16. According to Rutherford, the nucleus is composed of positively charged particles called
______________ and neutral particles called _______________________.
17. What did Rutherford suggest about how the electrons were moving?
18. Rutherford was unable to explain why the negative __________________ aren’t attracted into
the positive _____________________, causing the atom to collapse.
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LECTURE NOTES
Lesson 37: Subatomic particles
Text 3.6
We’ll work backwards for a moment, accept there is an atom, and remind you of the three
major subatomic particles located in an atom that you learned about in 8th grade.
Symbol
Location
Charge
Relative mass
Actual mass
Electrons
e-
Outside nucleus
Negative
1/1840
9.11 x 10-28
Protons
p+
In nucleus
Positive
1
1.673 x 10-24
Neutrons
n0
In nucleus
Neutral
1
1.675 x 10-24
A. Distinguishing Among Atoms
1. Atomic Number, Mass Number, and Electrons
a. Atomic Number (Z)
i.
The number of protons in the nucleus of each atom of that element
ii.
Atoms are identified by their atomic number
iii.
Because atoms are neutral,
# protons = # electrons
iv.
Periodic Table is in order of increasing atomic number
b. Mass Number (A)
i.
The total number of protons and neutrons in the nucleus of an isotope
c. Electrons
i.
The volume of an atom is from the area in which the electrons move
ii.
The chemical properties of an atom arise from the electrons.
B. Calculating the number of electrons, protons, and neutrons (Introduction)
A = protons + neutrons
Z = protons
Therefore, to get the number of neutrons, subtract A - Z
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Lesson 38: Dalton’s Atomic Theory
Text: 3.3
John Dalton (1766 – 1844) explained observations such as the law of constant composition
(a compound always has the same composition) using his atomic theory. The predictive
value of the theory led to its eventual acceptance.
A.
Defining the Atom
1. Atomic Theory
a. All matter is made up of very tiny particles called atoms
b. Atoms of the same element are chemically alike
c. Individual atoms of an element may not all have the same mass. However, the
atoms of an element have a definite average mass that is characteristic of the
element
d. Atoms of different elements have different average masses
e. Atoms are not subdivided, created, or destroyed in chemical reactions
1e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass.
The volume of the hydrogen nucleus is about one trillion times less than the volume of the hydrogen atom, yet the nucleus contains almost all the mass in the form of one
proton. The diameter of an atom of any one of the elements is about 10,000 to 100,000 times greater than the diameter of the nucleus. The mass of the atom is densely
packed in the nucleus.
The electrons occupy a large region of space centered around a tiny nucleus, and so it is this region that defines the volume of the atom. If the nucleus (proton) of a hydrogen
atom were as large as the width of a human thumb, the electron would be on the average about one kilometer away in a great expanse of empty space. The electron is
almost 2,000 times lighter than the proton; therefore, the large region of space occupied by the electron contains less than 0.1 percent of the mass of the atom.
2. Sizes of Atoms
a. Atomic radius
i. 40 to 270 picometers (pm)
1. 1 pm = 10-12m
ii. Most of the atomic radius is due to the electron cloud
b. Nuclear radius
i. 0.001 pm
ii. density is 2x108 metric tons/cm3
1. 1 metric ton = 1000kg
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B.
Models of the Atom
A model is a representation of nature, an attempt to communicate an explanation.
Scientist
Year
Model
Experiment
Focus
Democritus
~ 400
B.C.E.
None
Suggested Atom
Dalton
1808
Solid sphere, tiny, indivisible,
indestructible particles
Weather data
Thomson
1897
Plum pudding
Cathode Ray Tube;
also invented mass
spectrometer
Electrons
Plank
1900
Energy emitted in discrete
quantities
Radiation from solids
Quanta
Rutherford
1911
Nuclear Atom; also called the
planetary model
Gold foil
Nucleus
Bohr
1913
Bohr Model, electrons travel in
discrete orbits
Spectrum of Hydrogen
Excited and Ground state
Einstein
1905
Wave mechanical model
Photoelectric Effect
Photons
Schrödinger
1926
Wave mechanical model
Schrödinger cat;
thought experiment
Schrödinger equation
Heisenberg
1929
Wave mechanical model
Heisenberg uncertainty
principle
Murray GellMann; George
Zweig
1970s
Standard Model
Quarks and leptons
(matter)
* There were many other models developed during this time period but we’ll only focus on these particular ones.
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C.
Contributed to the Models of the Atom
Maxwell
1873
Visible light consists of
electromagnetic waves
Planck
1900
Energy emitted in discrete
quantities
Chadwick
1932
Provides description of
light
Radiation from solids
Quanta; Plank’s constant
Identified subatomic
particle
Neutron
Lesson 39: The Structure of the Atom / Thomson’s model
Text: 3.5
A. The Electron
a. Discovery
i. Joseph John Thomson (1897)
1. Cathode ray tube produces a ray with a constant charge to mass
ratio
2. All cathode rays are composed of identical negatively charged
particles (electrons)
B. Plum-pudding model
C. Inferences from the properties of electrons
i. Atoms are neutral, so there must be positive charges to balance the
negatives
ii. Electrons have little mass, so atoms must contain other particles that
account for most of the mass
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Lesson 40: Rutherford’s model
Text 11.1
1a. Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass.
An atom consists of a nucleus made of protons and neutrons that is orbited by electrons. The number of protons, not electrons or neutrons, determines the unique properties of
an element. This number of protons is called the atomic number. Elements are arranged on the periodic table in order of increasing atomic number. Historically, elements were
ordered by atomic mass, but now scientists know that this order would lead to misplaced elements (e.g., tellurium and iodine) because differences in the number of neutrons for
isotopes of the same element affect the atomic mass but do not change the identity of the element.
D. Structure of the Nucleus
1. The Nucleus
a. The Rutherford Experiment (1911)
b. Alpha particles (helium nuclei) fired at a thin sheet of gold
i. Assumed that the positively charged particles were bounced back if they
approached a positively charged atomic nucleus head-on (Like charges
repel one another)
Results from gold foil experiment
1. Very few particles were greatly deflected back from the gold sheet
a. nucleus is very small, dense and positively charged
b. most of the atom is empty space
2. Structure of the Nucleus
a. Protons
i.
Positive charge, mass of 1.673x10-27kg
ii.
The number of protons in the nucleus determines the atom's identity
and is called the atomic number (Z)
b. Neutrons
i.
James Chadwick (1932)
ii.
No charge, mass of 1.675x10-27kg
c. Nuclear Forces
i.
Short range attractive forces:
a. neutron-to-neutron, proton-to-proton,
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Unanswered Questions
What are the electrons doing? How are the electrons arranged? How do electrons
move? Why aren’t electrons (negatively charged) attracted to the positive nucleus?
Lesson 41: Bohr’s model
Text 11.5
The energy of the electrons is restricted to certain discrete values; that is, the energy is
quantized.
Consider the rungs of a ladder. There is no “in between” on a ladder. Your foot is
either on a rung or it is not.
The electrons move from each orbital. Photons (packets of light) are either absorbed or
released. If it is at the lowest it’s called ground state and the highest is called excited state.
Image from http://imagine.gsfc.nasa.gov/docs/teachers/lessons/xray_spectra/background-atoms.html
The flame test shows the spectra changes based on the elements.
Lesson 42: Wave model Introduction
Text 11.6
The Bohr model explained the hydrogen spectrum very well but it failed to explain the
spectra of all other atoms. Additional spectra analysis of elements supported a new model of
the atom, called the wave model. In this model, orbits do not exist. Instead, orbitals that
match spectra are discussed. These orbitals are s (sharp), p (principal), d (diffuse), and f
(fundamental). Orbitals do not describe the path or motion of the electron. Instead, they
describe the probability of finding an electron at a particular time.
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QUESTIONS AND PROBLEMS
1. How is Thomson’s model of the atom different from Dalton’s model of the atom? Draw a
picture of each.
2. How is the Rutherford model different from the previous models?
3. What was wrong with Rutherford’s model of the atom? Why did it need to be modified?
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JOURNAL
What points in the material strike you as important?
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What new material have you learned?
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List several questions about what you learned?
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How does what you learned relate to other information that you have learned in this course?
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How did the class work stimulate your thinking?
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