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Dalton’s Atomic Theory (experiment based!)
John Dalton
(1766 – 1844)
1) All elements are composed of
tiny indivisible particles called
atoms
2) Atoms of the same element are
identical. Atoms of any one
element are different from
those of any other element.
3) Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds
4) In chemical reactions, atoms are combined,
separated, or rearranged – but never
changed into atoms of another element.
Thomson’s Atomic Model
J. J. Thomson
Thomson believed that the electrons
were like plums embedded in a
positively charged “pudding,” thus it
was called the “plum pudding” model.
Conclusions from the Study of
the Electron:
a) Cathode rays have identical properties
regardless of the element used to
produce them. All elements must contain
identically charged electrons.
b) Atoms are neutral, so there must be
positive particles in the atom to balance
the negative charge of the electrons
c) Electrons have so little mass that atoms
must contain other particles that account
for most of the mass
The Rutherford Atomic Model
• Based on his experimental evidence:
–The atom is mostly empty space
–All the positive charge, and almost all
the mass is concentrated in a small area
in the center. He called this a “nucleus”
–The nucleus is composed of protons
and neutrons (they make the nucleus!)
–The electrons distributed around the
nucleus, and occupy most of the volume
–His model was called a “nuclear model”
QUICK CHECK
•
Subatomic Particles
Particle
Electron
(e-)
Charge Actual Mass (g)
Relative
Mass
(amu)
-1
9.11 x 10-28
0.0005486
Proton
(p+)
+1
1.67 x 10-24
1.007276
Neutron
(no)
0
1.67 x 10-24
1.008665
Measuring Atomic Mass
• Instead of grams, the unit we use
is the Atomic Mass Unit (amu)
• It is defined as one-twelfth the
mass of a carbon-12 atom.
– Carbon-12 chosen because of its isotope purity.
• Each isotope has its own atomic
mass, thus we determine the
average from percent abundance.
HOW DO ATOMS DIFFER
• OBJECTIVES:
–Explain what makes
elements and isotopes
different from each other.
Atomic Number
• Atoms are composed of identical
protons, neutrons, and electrons
– How then are atoms of one element
different from another element?
• Elements are different because they
contain different numbers of PROTONS
• The “atomic number” of an element is
the number of protons in the nucleus
• # protons in an atom = # electrons
Atomic Number
Atomic number (Z) of an element is
the number of protons in the nucleus
of each atom of that element.
Element
# of protons
Atomic # (Z)
Carbon
6
6
Phosphorus
15
15
Gold
79
79
Mass Number
Mass number is the number of
protons and neutrons in the nucleus
of an isotope: Mass # = p+ + n0
p+
n0
e- Mass #
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Nuclide
Oxygen - 18
Using a periodic table and what you know about atomic number,
mass, isotopes, and electrons, fill in the chart:
Element
Symbol
Atomic
Number
Atomic
Mass
# of
protons
# of
neutron
# of
electron
8
8
8
39
Potassium
Br
0
45
30
35
Atomic Number = Number of Protons
Number of Protons + Number of Neutrons = Atomic Mass
Atom (no charge) : Protons = Electrons
charge
0
30
Using a periodic table and what you know about atomic number,
mass, isotopes, and electrons, fill in the chart:
ANSWER KEY
Element
Symbol
Atomic
Number
Atomic
Mass
# of
protons
# of
neutron
# of
electron
charge
O
8
16
8
8
8
0
Potassium
K
19
39
19
20
19
0
Bromine
Br
35
80
35
45
35
0
Zinc
Zn
30
35
30
65
30
0
Oxygen
Atomic Number = Number of Protons
Number of Protons + Number of Neutrons = Atomic Mass
Atom (no charge) : Protons = Electrons
Isotopes are atoms of the same element having
different masses, due to varying numbers of
neutrons.
Isotope
Protons Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
1
1
2
Hydrogen-3
(tritium)
Nucleus
Isotopes
• Atoms of the same element with different
mass numbers.
• Nuclear symbol:
Mass #
12
Atomic #
6
• Hyphen notation: carbon-12
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
C
Isotopes
Neutron
+
Electrons
Nucleus
+
+
+
+
+
Nucleus
Carbon-12
Neutrons 6
Protons
6
Electrons 6
Nucleus
Proton
Proton
+
+
+
+
Neutron
Electrons
+
+
Carbon-14
Neutrons 8
Protons
6
Electrons 6
Nucleus
17
Cl
Isotopes
37
• Chlorine-37
– atomic #:
17
– mass #:
37
– # of protons:
17
37
– # of electrons:
17
– # of neutrons:
17
20
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Cl
Relative Atomic Mass
•
12C
atom = 1.992 × 10-23 g
• atomic mass unit (amu)
• 1 amu = 1/12 the mass of a 12C atom
• 1 p = 1.007276 amu
1 n = 1.008665 amu
1 e- = 0.0005486 amu
Neutron
+
Electrons
Nucleus
+
+
+
+
+
Nucleus
Carbon-12
Neutrons 6
Protons
6
Electrons 6
Proton
6Li
7Li
3 p+
3 n0
3 p+
4 n0
2e– 1e–
2e– 1e–
Neutron
Neutron
Electrons
Electrons
+
Nucleus
+
+
Nucleus
+
Proton
+
Nucleus
Nucleus
Lithium-6
Neutrons
Protons
Electrons
Lithium-7
Neutrons
Protons
Electrons
3
3
3
+
4
3
3
Proton
Check the atomic weight of elements in the
periodic table.
If the number of protons and neutrons are whole
numbers, why is the atomic mass NOT
a whole number?
Calculating averages
• You have five rocks, four with a mass of 50 g, and one
with a mass of 60 g. What is the average mass of the
rocks?
• Total mass = (4 x 50) + (1 x 60) = 260 g
• Average mass = (4 x 50) + (1 x 60) = 260 g
5
5
• Average mass = 4 x 50 + 1 x 60 = 260 g
5
5
5
California WEB
Average Atomic Mass
• weighted average of all isotopes
• on the Periodic Table
• round to 2 decimal places
Avg.
(mass)(%) + (mass)(%)
Atomic =
100
Mass
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Average Atomic Mass
• EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20%
18O.
Avg.
(16)(99.76) + (17)(0.04) + (18)(0.20)
Atomic =
100
Mass
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
=
16.00
amu
Isotopes
• Because of the existence of isotopes, the mass of a
collection of atoms has an average value.
• Average mass = ATOMIC WEIGHT
• Boron is 20% B-10 and 80% B-11.
11 is 80 percent abundant on earth.
That is, B-
• For boron atomic weight
= 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu
Atomic Mass
Calculate the atomic mass of copper if copper has two isotopes.
69.1% has a mass of 62.93 amu and the rest has a mass of 64.93
Percent
amu.
Isotope
Mass
Abundance
Cu-63
69.1
62.93
43.48463
Cu-65
30.9
64.93
20.06337
63.548
Average atomic mass (AAM)  (% " A" )(mass " A" )  (% " B" )(mass " B" )  ...
A.A.M.  (0.691)(62.93 amu)  (0.309)(64.93 amu)
A.A.M.  43.48463 amu  20.06337 amu
A.A.M.  63.548 amu for Copper
Cu
63.548
29