Download Ch 3 Notes Atoms

Ch 3: Atoms
Chapter 3 problem set:
Problem Set Ch 3: page 89-90
1-3, 7-9, 12-13, 17-20, 22-24
3.1 The Atom: From Idea to Theory
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Historical Background- In
approximately 400 BC, Democritus
(Greek) coins the term
“atom” (means indivisible). Before that
matter was thought to be one continuous
piece - called the continuous theory of
matter. Democritus creates the
discontinuous theory of matter. His
theory gets buried for thousands of years
18th century - experimental evidence
appears to support the idea of atoms.
Law of Conservation of Mass –
Antoine Lavosier (French) -1700’s
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The number of each kind of atoms
on the reactant side must equal the
number of each kind of atoms on
the product side
A +
B
+
C
—> ABC
Law of Multiple Proportions – John
Dalton (English) - 1803
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The mass of one element combines
with masses of other elements
simple in whole number ratios.
Water (H2O) is always: 11.2% H;
88.8% O
Sugar (C6H1206) is always: 42.1%
C; 6.5% H; 51.4% O
Law of Multiple Proportions Video
Law of Multiple Proportions – John
Dalton (English) - 1803
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Ex1:
wt. of H
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H +
H +
O  H 2O
0  H2O2
2
2
wt. of O
16
32
The ratio of O in H2O2 to O in H2O = 32/16 = 2:1
(small whole numbers
Dalton’s Atomic Theory
3.2 The Structure of the Atom
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Updating Atomic Theory
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1870’s - English physicist
William Crookes - studied
the behavior of gases in
vacuum tubes(Crookes
tubes - forerunner of picture
tubes in TVs).
Crookes’ theory was that
some kind of radiation or
particles were traveling from
the cathode across the tube.
He named them cathode
rays .
3.2 The Structure of the Atom
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20 years later, J.J. Thomson (English) repeated
those experiments and devised new ones. Cathode
Ray Tube
Thomson used a variety of materials, so he figured
cathode ray particles must be fundamental to all
atoms. 1897 - discovery of the electron.
3.2 The Structure of the Atom
Thomson and Milliken (oil drop experiment)
worked together (their data, not
them) to discover the charge and mass of
the electron
3.2 The Structure of the Atom
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Charge and Mass of the electron Thomson and Milliken (oil drop
experiment) worked together to
discover the charge and mass of the
electron Oil Drop
charge = 1.602 x 10-19 coulomb this is
the smallest charge ever detected
 mass
= 9/109 x 10-28 g this weight is
pretty insignificant
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3.2 The Structure of the Atom
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1909 - Gold Foil Experiment
(Rutherford - New Zealand)
Nuclei are composed of ‘nucleons’:
protons and neutrons
Alpha particles from
Polonium (in the lead
box) were released
towards a thin sheet of
gold foil. Most of the
particles went through
and were seen on the
detector screen. 1 in
20,000 alpha particles
bounced back.
Rutherford’s Conclusion
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Concluded:
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1–
the
2–
3–
4–
the positive portion of the atom is in
middle
most of the atom is empty
most of the mass is in the middle
electrons orbit the nucleus
Analogy: if an atom is the size of the
Linc, then the nucleus is the size of a
tennis ball floating in the middle of the
stadium.
Table: Subatomic particles
important in chemistry.
http://curriculum.media.pearsoncm
g.com/curriculum/science/cg_wsmw
_chem_12/UntamedSci/CHEM-UT04/player.html
Table: Subatomic particles
important in chemistry.
Nuclear Forces
Forces in Atoms
3.3 Weighing and Counting Atoms
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We look to the periodic table to give us
information about the number of particles
are in atoms and also to help us count
atoms in a sample.
Counting nucleons
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Atomic Number (Z) Atomic #
 Number of protons in the nucleus
 Uniquely labels each element
Mass Number (M) Mass #
 Number of protons + neutrons in the
nucleus
Counting nucleons
Counting electrons
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Atoms
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Same number of electrons and protons
Ions
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Ionic charge (q) = #protons #electrons
Positive ions are cations
Negative ions are anions
Review of formulas
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atomic # (Z) - (always a whole number, smaller
number on the periodic table) = # of protons in
the nucleus - also indicates the # of electrons if
the element is not charged
atomic mass – the average mass of all of the
isotopes of an element – is a number with a
decimal – is always the larger number on the
periodic table.
mass number (A) - sum of the protons and
neutrons in a nucleus
this number is rounded from atomic mass due to
the fact that there are isotopes
# neutrons = A - Z
example - # of
neutrons in Li = 6.941-3 = 3.941 rounds to 4
Ion – a charged atom. Atoms become charged by
gaining electrons (become a negative charge) or
losing electrons (become a positive charge)
Lots of Practice!!!
p+
e-
n°
Atomic # =
(# of p+)
Mass # =
(p+ + n0)
C
6
6
6
6
12
Ca
20
20
20
20
40
U
92
92
146
92
238
Cl
17
17
18
17
35
Mg
12
12
12
12
24
14C
6
6
8
6
14
S-2
16
18
16
16
32
Na+1
11
10
12
11
23
Isotopes
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Isotopes Isotope Video
Two atoms of the same element
(same # of p+) but with different
masses (different # of n0)
Average Atomic Mass (“weighted
average”)
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Definition - The average weight of the
natural isotopes of an element in their
natural abundance.
History lesson - originally H was the
basis of all atomic masses and was given
the mass of 1.0. Later, chemists changed
the standard to oxygen being 16.000
(which left H = 1.008). In 1961,
chemists agreed that 12C is the standard
upon which all other masses are based.
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1/12 of the mass of 1 atom of 12C = 1 amu
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Carbon consists of two isotopes:
98.90% is C-12 (12.0000 amu).
The rest is C-13 (13.0034 amu).
Calculate the average atomic
mass of carbon to 5 significant
figures.
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(.9890)(12.0000)+(.0110)(13.0034)=x
11.8680+.1430=12.011
Ex1: Chlorine consists of two natural isotopes, 35Cl (34.96885)
at 75.53% abundance and 37Cl (36.96590) at 24.47%
abundance. Calculate the average atomic mass of Chlorine.
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(.7553)(34.96885)+(.2447)(36.96590)=x
26.41+9.045=35.46
Ex2: Antimony consists of two natural isotopes
57.25% is 121Sb (120.9038). Calculate the % and
mass of the other isotope if the average atomic
mass is 121.8.
(.5725)( 120.9038)+(.4275)(x) =121.8
69.22+.4275x=121.8
-69.22- 69.22.4275x= 52.59x=123.0
The Mole, Avogadro’s number and
Molar Mass
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The Mole Mole Video
Atoms are tiny, so we count them in
“bunches”.
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A mole is a “bunch of atoms”.
The Mole (definition) -The amount of a
compound or element that contains
6.02 x 1023 particles of that
substance. Avogadro
1 mole = 1 gram formula mass =
6.02 x 1023 particles
Molar Mass - the sum of the atomic
masses of all atoms in a formula
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Round to the nearest tenth!
(measured in amu or grams)
ex - H2 H2O
Ca(OH)2
2.0g 18.0 g
74.1 g
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Molar mass is a term that can be used for
atoms, molecules (covalent compounds or
elements) and formula units (ionic
compounds)
Official names may also be:
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Formula mass (ionic compounds)
Molecular mass (covalent compounds and
diatomic elements)
Atomic weight, Atomic mass, grams formula
weight, etc.
Molar Mass
Examples:
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1 mole Na = 6.02 x 10 23atoms = 23.0g
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1 mole O2 = 6.02 x 1023 molecules=32.0g
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1 mole HCl = 6.02 x 1023 molecules = 36.5 g
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1 mole NaCl =6.02 x 1023 formula units = 58.5 g
Mole Relationships
1 mole
6.02 x 1023 atom/molecule
P-Table for gram
22.4
liter (at STP