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Chemistry
Julien
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Chapter 4: Atomic Structure
Atomic and Molecular Structure
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. As a basis for
understanding this concept:
e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass.
h.*Students know the experimental basis for Thomson’s discovery of the electron, Rutherford’s
nuclear atom, Millikan’s oil drop experiment, and Einstein’s explanation of the photoelectric
effect.
11. Nuclear processes are those in which an atomic nucleus changes, including radioactive decay of
naturally occurring and human-made isotopes, nuclear fission, and nuclear fusion. As a basis for
understanding this concept:
g.*Students know protons and neutrons have substructures and consist of particles called quarks.
Investigation and Experimentation
1. Scientific progress is made by asking meaningful questions and conducting careful
investigations. As a basis for understanding this concept and addressing the content in the other
four strands, students should develop their own questions and perform investigations. Students
will:
b. Identify and communicate sources of unavoidable experimental error.
c. Identify possible reasons for inconsistent results, such as sources of error or uncontrolled
conditions.
k. Recognize the cumulative nature of scientific evidence.
II. Structure of the Nuclear Atom.
A. Subatomic particles.
 Chemists design materials to fit specific needs.
1. Electrons.
a. Electrons—
b. J. J. Thomson discovered electrons when he connected a high voltage source to the cathode and
anode of a vacuum tube and produced a beam of electrons.
c. Cathode ray—
d. Thomson discovered that a negative plate would repel the beam of electrons and that a paddle
wheel placed on a track in the line of the electron beam would spin the paddle wheel.
e. No matter which gas Thomson used in his vacuum tube, the charge-to-mass ratio was the same.
f. Conclusions—
Chemistry
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Chapter 4 Notes
g. Thomson proposed that the atom was like a plum pudding with the electrons being like the plum
pieces found in the pudding.
h. Robert Milliken sprayed negatively charged particles of oil between oppositely charged plates
and was able to calculate the charge and mass of an electron.
2. Protons and neutrons.
a. Eugen Goldstein hypothesized that the beam of negative particles should produce a beam of
positive particles that would travel in the opposite direction of the negative particles and
discovered that beam.
b. Proton—
c. James Chadwick was able to discover a particle, with about the same mass as the proton and no
charge.
d. Neutrons—
e. Protons and neutrons have been broken down in particle accelerators and produce particles
known as quarks.
Properties of Subatomic Particles
Particle
Symbol
Relative
Charge
Relative Mass
(mass of proton = 1)
Actual Mass
(g)
Electron
Proton
Neutron
B. The atomic nucleus.
1. Rutherford’s gold-foil experiment.
a. Ernest Rutherford, a student of J. J. Thomson, designed a test to get a picture of the nucleus of an
atom by firing a nuclear gun at a very thin gold foil, hoping to pick up the image of the nucleus
on a fluorescent screen.
b. Most of the particles passed through the foil to the great surprise of the research team.
c. Conclusion—
2. The Rutherford atomic model.
a. Rutherford’s model of the atom concluded that the atom was mostly empty space and that the
positive charge and most of the mass was concentrated in the small, positive center region.
b. Nucleus—
 In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed
around the nucleus and occupy almost all the volume of the atom.
III. Distinguishing Between Atoms.
A. Atomic number.
1. Scientists discovered that different elements contained different amounts of positive charge.
Chemistry
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Chapter 4 Notes
 Elements are different because they contain different numbers of protons.
Atoms of the First Ten Elements
Name
Symbol
Atomic
Number
Protons
Neutrons
Mass
Number
H
He
Li
Be
B
C
N
O
F
Ne
1
2
3
4
5
6
7
8
9
10
1
2
3
4
5
6
7
8
9
10
0
2
4
5
6
6
7
8
10
10
1
4
7
9
11
12
14
16
19
20
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Number
of
Electrons
1
2
3
4
5
6
7
8
9
10
Practice Problems
15. Complete the table.
Element
K
Atomic
Number
19
S
V
Protons
Electrons
19
5
16
23
16. How many protons and electrons are in each atom?
a. fluorine (atomic number = 9)
b. calcium (atomic number = 20)
c. aluminum (atomic number = 13)
B. Mass number.
1. Mass number—
 The number of neutrons in an atom is the difference between the mass number and atomic number.
Number of neutrons = mass number – atomic number
2. The atom is composed of a very dense nucleus containing protons and neutrons and has electrons
that take 1840 of them to equal the weight of a proton.
Chemistry
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Chapter 4 Notes
3. Identifying the atoms of different elements can be done with a notation system that puts the mass
number as a superscript to the left of the element’s symbol and the atomic number as the subscript to
the left of the element’s symbol.
A
Z
E
A = mass number, Z = atomic number, E = element symbol
Practice Problems
17. How many neutrons are in each atom?
16
a. 8 O
80
d. 35 Br
32
b. 16 S
108
c. 47 Ag
207
e. 82 Pb
18. Use the table on page 4 to express the composition of each atom in shorthand form.
a. carbon–12
b. fluorine–19
c. beryllium–9
C. Isotopes.
1. Isotopes—
Practice Problems
19. Three isotopes of oxygen are oxygen–16, oxygen–17, and oxygen–18. Write the symbol for each, including
the atomic number and mass number.
20. Three isotopes of chromium are chromium–50, chromium–52, and chromium–53. How many neutrons are
in each isotope, given that chromium has an atomic number of 24?
Chemistry
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Chapter 4 Notes
 Because isotopes of an element have different numbers of neutrons, they also have different mass
numbers.
D. Atomic mass.
1. Writing the mass of individual atoms involves using exponents in the range of 10–22 to 10–24g.
2. To simplify matters, scientists have developed a system that is proportional based on carbon–12
being equal to exactly 12 amu.
3. Atomic mass unit—
4. Atomic mass—
5. The atomic mass is based on both the percentage abundance of the different isotopes element and the
weight of the individual isotopes.
 To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance,
expressed as a decimal, and then add the products.
Practice Problems
23. The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative
abundance and atomic masses are 69.2% for mass = 62.93 amu, and 30.8% for mass= 64.93% amu.
Calculate the average atomic mass of copper.
24. Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative
abundance of 78.92 amu (50.69%) and 80.92 amu (49.31%).
E. The Periodic Table—A Preview.
1. Periodic table—
2. Period—
3. Group—
PRACTICE FOR THE TEST:
Standardized Test Prep, page 125
Chemistry
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Chapter 4 Notes
Internet Practice:
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fix=0000 Click on the arrow above your Chemistry book.
Your grades can be accessed through http://web.me.com/ejulien/Mr._Juliens_Homepage/Welcome.html
Chemistry
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Chapter 4 Notes