Download Chapter 3: Elements, Compounds, and the Periodic Table

Discovery of Subatomic Particles
Chapter 3:
Elements, Compounds,
and the Periodic Table
 Earliest theories about atoms, such as
Dalton’s atomic theory, imagined them
to be indestructible.
 Experiments performed in the late
1800’s and early 1900’s showed atoms
are composed of subatomic particles.
Graded Set: p.100 #38, 76, 78,
90, 92, 94, 102
Bonus Set: p.100 #2, 7, 10, 13,
19, 22, 51, 75, 77, 85
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Discovery of Electron
Discovery of Proton
JJ Thomson (1897)
 Detected using Mass Spectrometer
 Discovered negatively
charged particles
 Electrons (e–)
 Electron removal resulted in an atom with a
positive charge
 Mass is 1800 times heavier than electron
 Cathode Ray Tube
Experiment:
negatively charged
particles moved from cathode (-) to anode (+)
(read p.64-65)
 Youtube
Proton (p)
 Positively
charged particle
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FYI
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Discovery of Atomic Nucleus
Rutherford’s Alpha Scattering Experiment
Also known as: Gold Foil Experiment
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 Most alpha () particles (positively charged)
passed right through gold
 A few deflected off at an angle (when they hit
the tiny nucleus, also positively charged)
 Youtube
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Discovery of Neutron
Rutherford’s Nuclear Atom
 Discovered by Chadwick 1932
 Nuclear mass always double the number
of protons.
 Therefore, nucleus must contain must be
another type of particle neutrons (1n)
 Demonstrated that nucleus:
 has almost all of mass in atom
 has a positive charge
 is very tiny, located at center of atom
 Where protons and
neutrons are
located
 Has same mass as proton
 Electrically neutral
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Remember…
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Atomic Structure
 Electrons (e–)
 Very low mass
 Occupy most of atom’s space
 Attraction between protons (p) and
electrons (e–) holds electrons around
nucleus
 Repulsion between electrons helps
them spread out over volume of atom
 In a neutral atom, number of electrons
must equal number of protons
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Summary: Subatomic Particles
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Atomic Notation
Atomic number (Z)
Nucleus (protons
+ neutrons)
 Number of protons that atom has in nucleus
 Unique to each type of element
Mass number (A)
Electrons
Particle
Mass (g)
Electrical
Charge
 Mass number = protons + neutrons
Electron 9.10939  10–28
–1
1.67264  10–24
+1
0
1 e
1
1
1 H, 1 p
0
1
0n
Proton
Neutron 1.67495 
10–24
Atomic Symbols
Symbol
 Symbolized by A
Z
X
Ex. What is the atomic symbol for helium?
Z = 2, A = 4
11
4
2 He
12
2
Isotopes
Example:
What is the isotopic symbol for
Uranium-235?
 Atoms of same element with different mass
numbers
 Same number of protons
 Different number of neutrons
 Most elements are mixtures of 2 or more
isotopes
 Chemically, isotopes behave alike
 e.g. Three isotopes of hydrogen (H)
Hydrogen-1, Hydrogen-2, Hydrogen-3
 Chemical symbol = U
 Mass number (A) = 235
 Atomic number (Z ) = 92
235
92
U
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Atomic Mass Units
Learning Check:
 Uniform mass scale for atoms
 Symbol: amu or u
 Based on the most abundant isotope:
carbon-12
 1 atom of carbon-12 = 12 atomic mass
units
Fill in the blanks:
symbol
60Co
81Br
65
29 Cu
206
82
Pb
neutrons protons
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electrons
33
27
27
46
35
35
36
29
29
124
82
82
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Calculating Atomic Mass
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Learning Check
 Generally, elements are mixtures of isotopes
e.g. Hydrogen
Isotope
Mass
% Abundance
1H
1.007825 u
99.985
2H
2.0140 u
0.015
Atomic mass
Naturally occurring magnesium is a mixture of 3
isotopes; 78.99% of the atoms are 24Mg (atomic mass,
23.9850 u), 10.00% of 25Mg (atomic mass, 24.9858 u),
and 11.01% of 26Mg (atomic mass, 25.9826 u). From
these data calculate the average atomic mass of
magnesium.
0.7899 x 23.9850 u = 18.95 u
24Mg
 Weighted average of masses of all stable
isotopes of given element
0.1000 x 24.9858 u = 2.499 u
0.1101 x 25.9826 u = 2.861 u
Total mass of avg = 24.31 u
25Mg
 Use isotopic abundances and masses
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26Mg
18
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Your Turn!
A naturally occurring element consists of two
isotopes. The data on the isotopes:
isotope #1
68.5257 u
60.226%
isotope #2
70.9429 u
39.774%
Calculate the average atomic mass of this element.
0.60226 × 68.5257 u = 41.270 u
0.39774 × 70.9429 u = 28.217 u
69.487 u
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