Transcript Ch. 3 Notes (Atomic Structure) Accelerated teacher

Ch. 3: Atomic Structure
The Theory of the Atom
•
________________,
Democritus a famous Greek teacher who lived in the 4th Century B.C.,
first suggested the idea of the atom.
• ________
John __________
Dalton came up with his atomic theory based on the results of
his experiments. (See p. 56)
The Atom
•
element
The smallest particle of an ________________
is an atom.
•
subatomic particles.
The atom is made up of three ________________
1897 by J. J. Thomson by using a cathode
(1) The electron was discovered in _______
(−) charge. It’s mass is much smaller than the
ray tube. The electron has a _______
ignored
other 2 subatomic particles, therefore it’s mass is usually ______________.
Cathode Ray Tube
(+) charge, and it was discovered in
(2) The proton has a ______
1886
_________
by E. Goldstein.
(3) The neutron does not have a charge. In other words, it is
neutral It was discovered in 1932
________.
____ by James Chadwick.
The neutron has about the same _________
mass as the proton.
visible matter
• These three particles make up all the ____________________
in the Universe!
• There are other particles such as neutinos, positrons, and
quarks, but are typically left for 2nd year chemistry courses.
Nuclear Atomic Structure
• The atom is made up of 2 parts/sections:
(1) The ______________
--- (in the center of the atom)
nucleus
electron _________
cloud --- (surrounds the nucleus)
(2) The ____________
(p+ & n0)
e− cloud
The Nucleus
• Discovered by Ernest ________________
in ________.
Rutherford
1911
• He shot a beam of positively charged “alpha particles”, which
are ___________
nuclei, at a thin sheet of ______
helium
gold_____.
foil
• 99.9% of the particles went right on
through to the ______________.
detector
•
• Some were slightly deflected. Some
even ____________
________
bounced
back
towards the source!
• This would be like shooting a
cannon ball at a piece of tissue paper
and having it bounce off.
Rutherford’s Experiment
Conclusions about the Nucleus
empty ___________.
space
(1) Most of the atom is more or less _________
tiny
(2) The nucleus is very _________.
(Stadium Analogy)
dense
(3) The nucleus is very ___________.
(Large Mass ÷ Small Volume)
(4) The nucleus is ______________
positively
charged.
Counting Subatomic Particles in an Atom
protons in the nucleus.
• The atomic # of an element equals the number of ____________
protons
neutrons
• The mass # of an element equals the sum of the _____________
and _____________
in the nucleus.
electrons
• In a neutral atom, the # of protons = # of ______________.
subtract
mass
• To calculate the # of neutrons in the nucleus, ______________
the ___________
#
atomic #.
from the __________
Practice Problems
(1) Find the # of e-, p+ and n0 for sodium. (mass # = 23)
Atomic # = 11 = # e- = # p+
2)
# neutrons = 23-11 = 12
Find the # of e-, p+ and n0 for uranium. (mass # = 238)
Atomic # = 92 = # e- = # p+
# neutrons = 238-92 = 146
3) What is the atomic # and mass # for the following atom?
# e- = 15; # n0 = 16
Atomic # = 15 = # e- = # p+
Mass # = p+ + n0 = 15+16 =31
The element is phosphorus!
Isotopes
protons
• An isotope refers to atoms that have the same # of ___________,
but they have a
neutrons
different # of ___________.
mass
• Because of this, they have different _________
#’s (or simply, different
masses
___________.)
• Isotopes are the same element, but the atoms weigh a different amount because
neutrons
of the # of ______________.
Examples--->
(1) Carbon-12 & Carbon-13
(2) Chlorine-35 & Chlorine-37
(The # shown after the name is the mass #.)
atomic
• For each example, the elements have identical ___________
#’s, (# of p+) but
mass #’s, (# of n0).
different _________
• Another way to write the isotopes in shorthand is as follows:
12
C
6
35
17
Cl
The top number is the ________
mass #, and the bottom # is the __________
atomic number.
subtracting the #’s!
Calculating the # n0 can be found by _____________
Figure 3.10: Two isotopes of sodium.
More Practice Problems
(1) Find the # e-, p+ and n0 for Xe-131.
Atomic # = 54 = p+ = e−
2)
Find the #
e-,
p+
and
n0
for
n0 = 131 − 54 = 77
63
29
Cu
Atomic # = 29 = p+ = e−
n0 = 63 − 29 = 34
3) Write a shorthand way to represent the following isotope:
# e- = 1
# n0 = 0
# p+ = 1
Atomic # = p+ = e− = 1
H-1 or
mass # = n0 + p+ = 1+ 0 = 1
1
1
H
Ions
•
An atom can gain or lose electrons to become electrically charged.
•
Cation = (___)
+ charged atom created by ___________
losing e-’s.
–
Cations are ______________
than the original atom.
smaller
–
_____________
generally form cations.
Metals
•
Anion = (___)
e-’s.
− charged atom created by _____________
gaining
–
larger
Anions are ____________
than the original atom.
–
_______________
generally form anions.
Nonmetals
Practice Problems: Count the # of protons & electrons in each ion.
a) Mg+2
12
10
p+ = _____
e− = ______
b) F−1
9
10
p+ = _____
e− = ______
Atomic Mass
12
• Based on the relative mass of Carbon-12 which is exactly _______.
1 amu
1 atomic mass unit (amu) 1 n0 ≈ __
• 1 p+ ≈ __
0 amu
1e- ≈ __
• The atomic masses listed in the Periodic Table are a “weighted average” of all
the isotopes of the element.
Weighted Average
Practice Problems:
(1) Mr. Turkowski’s Algebra 1 semester grades are calculated using a weighted average
of three category scores:
Major Grades= 60% of your grade
Minor Grades= 30% of your grade
Semester Exam=10% of your grade
• If a student had the following scores, what would they receive for the semester?
Major= 80 (B − )
Minor= 60 (D −)
Semester Exam=65 (D)
Weighted Average
Step (1): Multiply each score by the % that it is weighted.
Step (2): Add these products up, and that is the weighted average!
60% x 80 = 48.0
30% x 60 = 18.0
10% x 65 = + 6.5
Add them up!!
72.5 (C−)
A “normal average” would be calculated by simply adding the raw
scores together and dividing by 3…
80 + 60 + 65 = 205 ÷ 3 = 68.3 = D
Average Atomic Mass
Practice Problems:
(2) In chemistry, chlorine has 2 isotopes:
Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance)
What is the weighted average atomic mass of chlorine?
35 x 0.758 = 26.53
37 x 0.2423 = + 8.9651
Add them up!!!
35.4951 amu
(3) Oxygen has 3 isotopes:
O-16 (99.76%) O-17 (0.037%)
Estimate oxygen’s average atomic mass.
Barely over 16.0 amu
O-18 (0.2%)
Average Atomic Mass
(4) Copper has an average atomic mass of 63.546 g/mole. It contains
only two natural isotopes, which are Cu-63, with an isotope mass of
62.940 and Cu-65 with an isotope mass of 64.928. What are the
percent of the two isotopes in naturally occurring copper?