1. Atomic Structure Big Idea #2 Electrons and the Structure of Atoms
Chemistry I Notes Ch. 4
Atomic Structure
Big Idea #2
Electrons and the Structure of Atoms

2. 4.1 Defining the Atom A. Democritus-Atomos-the indivisible part
B. Lavosier–Law of Conservation of Matter
C.   Proust- Law of Constant Composition
D. Dalton – Atomic Theory of Matter
1.  Each element is composed of extremely small particles called atoms
2.  All atoms of a given element are identical, but they differ from the atoms of other elements.
3.  Atoms are neither created nor destroyed in any chemical reaction.
4.  A given compound has the same relative number and kinds of atoms

3. Sizing up the Atom
This liquid mercury illustrates Dalton’s concept of the atom.
Every drop, no matter its size has the same properties.
Even if you could make a drop the size of one atom, it would still have the chemical properties of mercury.
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4. Sizing up the Atom
If you were to grind a copper coin into a fine dust, each speck in the small pile of shiny red dust would still have the properties of copper.
If you could continue to make the copper dust smaller, you would eventually come upon a particle of copper that could no longer be divided and still have the chemical properties of copper.
This final particle is an atom.
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5. Sizing up the Atom Atoms are very small.
A pure copper coin the size of a penny contains about 2.4  1022 atoms.
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6. Sizing up the Atom Atoms are very small.
A pure copper coin the size of a penny contains about 2.4  1022 atoms.
By comparison, Earth’s population is only about 7  109 people.
If you could line up 100,000,000 copper atoms side by side, they would produce a line only 1 cm long!
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7. Sizing up the Atom
Despite their small size, individual atoms are observable with instruments such as scanning electron microscopes.
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8. If an atom has a radius of 1  m, how many of these atoms must be lined up in a row to produce a line 1 m long?
1  1010 (10,000,000,000) atoms of radius 1  m would need to be lined up in a row to produce a line 1 m long.
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9. The nuclear atom
When subatomic particles were discovered, scientists wondered how the particles were put together in an atom.
Most scientists—including J. J. Thompson—thought it likely that the electrons were evenly distributed throughout an atom filled uniformly with positively charged material.
In Thomson’s atomic model, known as the “plum-pudding model,” electrons were stuck into a lump of positive charge, similar to raisins stuck in dough.
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10. 4.2 Structure of the Nuclear Atom
Faraday- Atoms contain particles that have electric charge
Thomson- Cathode rays are composed of negatively charged particles - electrons (e-) mass=1/1874 amu
Becquerel & Curie – radioactivity
Rutherford- Discovered nucleus which contains positively charged Particles protons (p+) Disproved the plum pudding model.

11. Subatomic Particles Electrons
One electrode, the anode became positively charged.
The other electrode, the cathode, became negatively charged.
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12. Subatomic Particles Electrons
The result was a glowing beam, or cathode ray, that traveled from the cathode to the anode.
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13. Subatomic Particles Electrons
Thomson found that a cathode ray is deflected by electrically charged metal plates.
A positively charged plate attracts the cathode ray, while a negatively charged plate repels it.
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14. A cathode ray can also be deflected by a magnet.
Subatomic Particles
Electrons
A cathode ray can also be deflected by a magnet.
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15. Subatomic Particles
Protons and Neutrons
In 1886, Eugen Goldstein (1850–1930) observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays.
He concluded that they were composed of positive particles.
Such positively charged subatomic particles are called protons.
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16. Subatomic Particles
Electrons
The U.S. physicist Robert A. Millikan (1868–1953) carried out experiments to find the quantity of an electron’s charge.
From his data, he found that the charge on each oil droplet was a multiple of 1.60  10–19 coulomb, meaning this must be the charge of an electron.
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17. The Atomic Nucleus
This model of the atom turned out to be short-lived, however, due to the work of a former student of Thomson, Ernest Rutherford (1871–1937).
Born in New Zealand, Rutherford was awarded the Nobel Prize for Chemistry in His portrait appears on the New Zealand $100 bill.
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18. Florescent
Screen
Uranium
Lead block
Gold Foil

19. The Atomic Nucleus Rutherford’s Gold-Foil Experiment
In the experiment, a narrow beam of alpha particles was directed at a very thin sheet of gold.
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20. The Atomic Nucleus Rutherford’s Gold-Foil Experiment
Rutherford’s results were that most alpha particles went straight through, or were slightly deflected.
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21. The Atomic Nucleus Rutherford’s Gold-Foil Experiment
Rutherford’s results were that most alpha particles went straight through, or were slightly deflected.
What was surprising is that a small fraction of the alpha particles bounced off the gold foil at very large angles.
Some even bounced straight back toward the source.
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22. 4.3 Distinguishing Among Atoms
   A. All atoms are composed of three subatomic particles
1. Protons – Positively charged particles found in
the nucleus. Mass approx 1 amu given by atomic
number (Z#).
2. Neutrons – Neutral charge also found in nucleus
Mass approx 1 amu. Number given by subtraction of the
Z# from the A#. Neutrons = A-Z
3. Electrons - Negatively charged particles found outside
the nucleus in the electron cloud. Mass approx 1/1837
amu in neutral atoms the number of electrons = number
of protons.

23. 4.3 Distinguishing Among Atoms cont…
Moseley- Each element has a unique number of protons, so the atomic number or Z# identifies it.
Mass number or A# is an element’s atomic mass rounded to the nearest whole number.
Isotopes are atoms of the same element with different numbers of neutrons.
Average atomic mass given in the periodic table is the weighted average of the masses of all isotopes of that given element.
Isotopes are identified with symbols that indicate the mass number as a superscript.

24. Subatomic Particles
Protons and Neutrons
In 1932, the English physicist James Chadwick (1891–1974) confirmed the existence of yet another subatomic particle: the neutron.
Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton.
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25. Properties of Subatomic Particles Relative mass (mass of proton = 1)
Interpret Data
The table below summarizes the properties of these subatomic particles.
Properties of Subatomic Particles
Particle
Symbol
Relative charge
Relative mass (mass of proton = 1)
Actual mass (g)
Electron e– 1–
1/1840
9.11  10–28
Proton p+ 1+ 1
1.67  10–24
Neutron n0
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26. Subatomic Particles
Although protons and neutrons are extremely small, theoretical physicists believe that they are composed of yet smaller subnuclear particles called quarks.
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27. How can there be different varieties of atoms?
CHEMISTRY & YOU
How can there be different varieties of atoms?
Just as there are many types of dogs, atoms come in different varieties too.
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28. Atomic Number and Mass Number
Elements are different because they contain different numbers of protons.
An element’s atomic number is the number of protons in the nucleus of an atom of that element.
The atomic number identifies an element.
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29. Atomic Number and Mass Number
The total number of protons and neutrons in an atom is called the mass number.
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30. Atomic Number and Mass Number
If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition.
Number of neutrons = mass number – atomic number
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31. Atomic Number and Mass Number
The composition of any atom can be represented in shorthand notation using atomic number and mass number.
Au is the chemical symbol for gold.
The atomic number (Z #) is the subscript.
The mass number (A#) is the superscript.
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32. Atomic Number and Mass Number
You can also refer to atoms by using the mass number and the name of the element.
Au is the chemical symbol for gold.
Au may be written as gold-197.
197 79
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33. Determining the Composition of an Atom
Sample Problem 4.2
Determining the Composition of an Atom
How many protons, electrons, and neutrons are in each atom?
a. Be b. Ne c. Na 9 4 20 10 23 11
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34. Subatomic Particles
Protons and Neutrons
You can think through this problem using four simple ideas about matter and electric charges.
Atoms have no net electric charge; they are electrically neutral.
Electric charges are carried by particles of matter.
Electric charges always exist in whole-number multiples of a single basic unit; that is, there are no fractions of charges.
When a given number of negatively charged particles combines with an equal number of positively charged particles, an electrically neutral particle is formed.
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35. Analyze List the knowns and the unknowns.
Sample Problem 4.2
Analyze List the knowns and the unknowns. 1
Use the definitions of atomic number and mass number to calculate the numbers of protons, electrons, and neutrons.
KNOWNS
Beryllium (Be)
atomic number = 4 mass number = 9
Neon (Ne)
atomic number = 10
mass number = 20
Sodium (Na)
atomic number = 11
mass number = 23
UNKNOWNS
protons = ?
electrons = ?
neutrons = ?
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36. atomic number = number of protons
Sample Problem 4.2
Calculate Solve for the unknowns. 2
Use the atomic number to find the number of protons.
atomic number = number of protons
a. 4 b. 10 c. 11
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37. atomic number = number of electrons
Sample Problem 4.2
Calculate Solve for the unknowns. 2
Use the atomic number to find the number of electrons.
atomic number = number of electrons
a. 4 b. 10 c. 11
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38. number of neutrons = mass number – atomic number
Sample Problem 4.2
Calculate Solve for the unknowns. 2
Use the mass number and atomic number to find the number of neutrons.
a. number of neutrons = 9 – 4 = 5
b. number of neutrons = 20 – 10 = 10
c. number of neutrons = 23 – 11 = 12
number of neutrons = mass number – atomic number
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39. Evaluate Do the results make sense?
Sample Problem 4.2
Evaluate Do the results make sense? 3
For each atom, the mass number equals the number of protons plus the number of neutrons.
The results make sense.
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40. Isotopes Isotopes How do isotopes of an element differ?
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41. Isotopes There are three different kinds of neon atoms.
How do these atoms differ?
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42. Isotopes All have the same number of protons (10).
All have the same number of electrons (10).
But they each have different numbers of neutrons.
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43. Isotopes
Isotopes are atoms that have the same number of protons but different numbers of neutrons.
Neon-20, neon-21, and neon 22 are three isotopes of neon.
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44. Atomic Mass Atomic Mass
How do you calculate the atomic mass of an element?
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45. Atomic Mass The mass of even the largest atom is incredibly small.
Since the 1920s, it has been possible to determine the tiny masses of atoms by using a mass spectrometer.
The mass of a fluorine atom was found to be x 10–23 g.
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46. Atomic Mass
Such data about the actual masses of individual atoms can provide useful information, but in general these values are inconveniently small and impractical to work with.
Instead, it is more useful to compare the relative masses of atoms using a reference isotope as a standard.
The reference isotope chosen is carbon-12.
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47. 4.3 Distinguishing Among Atoms cont…
Atomic mass unit or AMU is the unit of mass used for atomic masses and mass numbers of elements. It is equal to 1/12 of the mass of a Carbon-12 atom
Atomic mass given in the periodic table is calculated by taking the mass of each isotope and multiplying by its percentage abundance in nature then summing.
Average atomic mass = (mass of isotope A x % abundance A) + (mass of isotope B x % abundance B) + (mass of isotope C x % abundance C) etc…

48. Interpret Data Natural Percent Abundance of
Stable Isotopes of Some Elements
Name
Symbol
Natural percent abundance
Mass (amu)
Atomic mass
Hydrogen H
99.985
0.015
negligible
1.0078
2.0141
3.0160
1.0079
Helium He
0.0001
4.0026
Carbon C
98.89
1.11
12.000
13.003
12.011
Oxygen O
99.759
0.037
0.204
15.995
16.995
17.999
15.999
Chlorine Cl
75.77
24.23
34.969
36.966
35.453 35 17 37 1 2 3 12 6 13 4 18 8 16
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49. Calculating Average Atomic Mass
What is the average atomic mass of Chlorine if
is 75.53% abundant with an atomic mass of amu and is 24.47% abundant with an atomic mass of ?
This is the weighted average of the atomic masses of these two isotopes of chlorine

50. Atomic Mass
The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element.
A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature.
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51. Calculating Average Atomic Mass
What is the average atomic mass of Silicon if
is 92.21% abundant with an atomic mass of amu, is 4.70% abundant with an atomic mass of amu, and is 3.09% abundant with an atomic mass of amu?
The weighted average of the atomic masses of these three isotopes of silicon is amu

52. Calculating Atomic Mass
Sample Problem 4.5
Calculating Atomic Mass
Element X has two naturally occurring isotopes. The isotope with a mass of amu (10X) has a relative abundance of percent. The isotope with a mass of amu (11X) has a relative abundance of percent. Calculate the atomic mass of element X.
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53. Analyze List the knowns and the unknown.
Sample Problem 4.5
Analyze List the knowns and the unknown. 1
The mass each isotope contributes to the element’s atomic mass can be calculated by multiplying the isotope’s mass by its relative abundance. The atomic mass of the element is the sum of these products.
KNOWNS
UNKNOWN
atomic mass of X = ?
Isotope 10X:
mass = amu
relative abundance = 19.91% =
Isotope 11X:
mass = amu
relative abundance = 80.09% =
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54. Calculate Solve for the unknowns.
Sample Problem 4.5
Calculate Solve for the unknowns. 2
Use the atomic mass and the decimal form of the percent abundance to find the mass contributed by each isotope.
for 10X: amu x = amu
for 11X: amu x = amu
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55. Calculate Solve for the unknowns.
Sample Problem 4.5
Calculate Solve for the unknowns. 2
Add the atomic mass contributions for all the isotopes.
For element X, atomic mass = amu amu
= amu
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56. Electrons and the Structure of Atoms
BIG IDEA
Electrons and the Structure of Atoms
Atoms of the same element have the same number of protons, which is equal to an atom’s atomic number.
But atoms of the same element can have different numbers of neutrons.
Atoms of the same element with different numbers of neutrons are isotopes.
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57. 4.3 Distinguishing Among Atoms cont…
Ions – Atoms that have lost or gained one or more electrons.
Positive ions have lost electrons (e-)
Negative ions have gained electrons (e-)
The symbol indicates the number of electrons lost or gained.
How many protons neutrons and electrons do these contain?

58. Changes in the Nucleus
Nuclear Stability – Changes in the nucleus of the atom can alter the identity of the atom.
For small atoms (Z# 1-20) an equal number of protons and neutrons keeps the nuclei stable.
As atoms increase the number of protons greater numbers of neutrons are needed to maintain stability. Why?
If the nucleus of an isotope is unstable it will undergo radioactive decay.
All elements with a Z# greater than 83 are radioactive.

59. Changes in the Nucleus cont…
Types of radioactive decay
A. Alpha – Helium-4 nuclei, mass 4 amu, 2+ charge, poor penetration
B. Beta – electron, essentially no mass, 1- charge some penetration
C. Positron -+electron, essentially no mass, 1+ charge, some penetration
D. Gamma-electromagnetic radiation no mass or charge, high penetration
E. Neutron – mass 1 amu, no charge, some penetration

60. Changes in the Nucleus cont…
Nuclear equations – keep track of mass and charge in nuclear reactions.
Do the practice problems!

61. QUESTION

62. ANSWER 2) Sc Section 18.1 Nuclear Stability and Radioactive46 21 Sc
Section 18.1 Nuclear Stability
and Radioactive
Decay (p. 841)
Make sure to memorize the abbreviations for the
subatomic particles.
HMCLASS PREP: Table 18.2

63. QUESTION

64. ANSWER 2) Th Section 18.1 Nuclear Sta bility and Radioactive
234 90 Th
Section 18.1 Nuclear Sta
bility and Radioactive
Decay (p. 841)
Just as chemical equations need the same
number of each type of atom on each side,
nuclear equations need the same number of
each type of nucleon on each side.

65. QUESTION

66. ANSWER 2) Ar Section 18.1 Nuclear Stability and Radioactive Decay (p.40 18 Ar
Section 18.1 Nuclear Stability and Radioactive
Decay (p.
841)
The electron is “captured” from the core
electrons swarming around the nucleus.
Remember to place the electron on the left side
of the reaction.