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Video 3-1
• Foundations of Atomic Theory
• Development of Atomic Models
• Forces in the Nucleus
Chapter 3
Atoms: The Building Blocks of Matter
I. Foundations of Atomic Theory
Several basic laws were introduced after
the 1790’s (emphasis on quantitative
analysis):
 Law of conservation of mass:
mass is neither created nor destroyed
during ordinary chemical or physical
processes.
I. Foundations of Atomic Theory
 Law of definite proportions:
chemical compounds contain the
same elements in exactly the same
proportions by mass regardless of the
size of the sample.
 Ex. NaCl always is composed of
39.34% sodium and 60.66% chlorine
by mass.
I. Foundations of Atomic Theory
 Law of multiple proportions: if two
or more different compounds are
composed of the same 2 elements,
the ratio of mass of the second
element combined with a certain
mass of the first is always a ratio of
small whole numbers.
I. Foundations of Atomic Theory
 Ex. CO and CO2: For the same mass
of carbon, the mass of the O in CO to
the mass of O in CO2 will be 1:2
 If you had 28 g of CO and 44 g of
CO2, both would contain 12 g of C.
The CO would contain 16 g of O and
the CO2 would contain 32 g of O.
I. Foundations of Atomic Theory
Masses in CO
Masses in CO2
Ratios
12 g C
16 g O
12 g C
32 g O
C= 12:12 = 1:1
O = 16: 32 = 1:2
II. Development of Atomic Models:
John Dalton (1808):
1. All matter is composed of extremely
small particles called atoms (cannot
be subdivided, created, nor
destroyed)
2. Atoms of the same element are
identical; atoms of different elements
are different
II. Development of Atomic Models:
John Dalton (1808):
3. Atoms combine in simple whole
number ratios to form compounds
4. In chemical reactions, atoms
combine, separate, or are
rearranged.
II. Development of Atomic Models:
John Dalton (1808):
Which of these were later proven wrong
and why?
 1) Atoms can be subdivided
 2) Atoms of the same element do
NOT have to be identical
II. Development of Atomic Models:
J.J. Thomson (1897) and Robert Millikan
(1909):
 Used cathode rays to determine that
atoms contained small negatively
charged particles called electrons.
 Atoms must also contain positive
charges to balance the negative
electrons
II. Development of Atomic Models:
II. Development of Atomic Models:
 Other particles must account for
most of the mass of the atom
 Millikan determine the size of the
charge on the electron (oil drop
experiment)
II. Development of Atomic Models:
Ernest Rutherford (1911):
 What was the structure of the atom?
 Gold Foil Experiment
 Thomson assumed mass and charged
particles were evenly distributed
throughout the atom (“plum-pudding”
model)
II. Development of Atomic Models:
II. Development of Atomic Models:
Ernest Rutherford (1911):
 Expected most of the particles to pass
with only slight deflection
 Most particles did, but some showed
wide-angle deflections (some almost
came back to the source).
II. Development of Atomic Models:
II. Development of Atomic Models:
Ernest Rutherford (1911):
  discovery of the NUCLEUS of the
atom
 small, dense, positively charged center
of the atom
 number of PROTONS in the nucleus
determines the atom’s identity
II. Development of Atomic Models:
 Rutherford Atomic Model (solar
system model)
III. Forces in the Nucleus
Repulsive forces should exist between
protons in the nucleus (like charges
repel).
 Why doesn’t the nucleus “fly apart”
due to the repulsive electromagnetic
force?
III. Forces in the Nucleus
Strong (nuclear) force:
 attractive force that acts over very
small distances in the nucleus
 causes proton-proton, protonneutron, neutron-neutron attractions
Note: gravitational force is present, but
negligible. Why?
Video 3-2
•
•
•
•
•
Atomic Dimensions
Properties of Atoms and Ions
Designating Isotopes
Elements on the Periodic Table
Average Atomic Mass
IV. Atomic Dimensions
How “big” are subatomic particles?
Particle
Symbol
electron e-
0
1
proton p+ 1
1
neutron n0
e
p
1
0
n
Relative
Charge
Mass
Number
Relative
Mass
(amu)
Actual
Mass (kg)
-1
0
0.0005486
9.109 x
10-31
+1
1
1.007276
1.673 x
10-27
0
1
1.008665
1.675 x
10-27
IV. Atomic Dimensions
How “big” are subatomic particles?
 Atomic radii: 40 to 270 pm
 Nuclear radii: about 0.001 pm
 Nuclear density: about 2 x 108 metric
tons/cm3
 1 amu (atomic mass unit) = 1.660540 x
10-27 kg
 Why?
V. Properties of Atoms and Ions
 atomic number (Z): number of protons in
an atom
 mass number: number of protons +
neutrons in an atom (number of nucleons—
particles in the nucleus)
 isotopes (nuclides): atoms of the same
element that have different masses
(different number of NEUTRONS)
V. Properties of Atoms and Ions
 ions: atoms with a charge (protons 
electrons)




charge = protons – electrons
atoms can only turn into ions by gaining or
losing ELECTRONS
cation: positive ion
anion: negative ion
VI. Designating Isotopes
 There are two ways to write symbols for an
isotope
1. name-(mass number)
massnumber
2. atomicnumber symbol
VI. Designating Isotopes
Examples:
 Hydrogen has 3 isotopes:
 Protium
 Deuterium
 Tritium
Hydrogen-1
Hydrogen-2
Hydrogen-3
1
1
2
1
3
1
H
H
H
 How many neutrons in each isotope?
 Note: mass number – atomic number =
number of neutrons
VII. Elements on the Periodic Table
 every periodic table will give you at least 3
pieces of information about elements:
Atomic Number
Symbol
Atomic mass (amu)
3
Li
6.941
VII. Elements on the Periodic Table
 What is the basis for the atomic mass unit
(amu)?
 1 amu = exactly 1/12 the mass of a
carbon-12 atom (6 protons, 6 neutrons, 6
electrons)
 All other atomic masses are based on
comparisons to C-12 (exactly 12 amu).
VII. Elements on the Periodic Table
 Example:
C-13 has a mass that is 1.083613 times
heavier than C-12. The mass of C-13 is
(1.083613) x 12 amu
= 13.003356 amu
VIII. Average Atomic Mass
 Carbon has 3 isotopes (nuclides):
C-12 (12 amu)
C-13 (13.003 amu)
C-14 (14.003 amu)
 Their average mass should be
(12 + 13.003 + 14.003) / 3
= 13.002 amu
VIII. Average Atomic Mass
 The atomic mass of carbon (periodic table)
is 12.011 amu.
 WHY?
VIII. Average Atomic Mass
 Atomic Mass of an element: weighted
average of all the atoms in a naturally
occurring sample of that element (NOTE:
not every atom in that sample has the
same mass)
 Ex. How would you determine the average
age of the students in this class?
VIII. Average Atomic Mass
 atomic mass = sum of (mass of each
isotope x percent abundance)
VIII. Average Atomic Mass
 Example:



C-12 (12 amu) 98.90%
C-13 (13.003 amu) 1.10%
C-14 (14.003 amu) trace
atomic mass of C = (12 amu)(0.9890) +
(13.003 amu)(0.0110) + (14.003 amu)(0)
= 12.011 amu
VIII. Average Atomic Mass
 NOTE: the atomic mass of most elements
will usually give you an idea of the most
common isotope of that element (mass
number that is closest to the atomic mass)
Video 3-3
• Counting Atoms and Stuff
IX. Counting Atoms and Stuff
 1 mole = 6.0221367 x 1023 things
 Ex. 1 mole of eggs contains 6.0221367 x
1023 eggs
 6.0221367 x 1023 = Avogadro’s number
(N)
1 mol = 6.022 x 1023
particles
IX. Counting Atoms and Stuff
 1 mole = 6.0221367 x 1023 things
 1 amu = 1.660540 x 10-24 g
 Suppose you had a sample of one mole of
particles. Each particle weighed exactly 1
amu.


How many amu would the sample weigh?
How many grams would the sample weigh?
IX. Counting Atoms and Stuff
 1 mole of amu = (6.0221367 x
1023)(1.660540 x 10-24 g)
= 1.00000 gram (exactly)
 What is the significance of this?
 How much will 1 mole of C-12 atoms weigh
(in grams)?
 12 grams (exactly)
IX. Counting Atoms and Stuff
Molar Mass:
 mass of one mole of a substance
 units: grams/mole
 equal to the ATOMIC MASS of the element
 for compounds, equal to the SUM of the
masses of all the elements in the
compound (multiply each elements’ atomic
mass by the subscript)
IX. Counting Atoms and Stuff
Example:
 Find the molar masses of the following:
 NaCl
 CO2
 Ca(C2H3O2)2
Na = 22.99
O = 16.00
58.44 grams/mol
44.01 grams/mol
158.168 grams/mol
Cl = 35.45
H = 1.008
C = 12.01
Ca = 40.08
IX. Counting Atoms and Stuff
1 mole of X = atomic mass of X (grams)
1 mole X = 6.022 x 1023 atoms
 These equalities will let you do
DIMENSIONAL ANALYSIS


Convert grams to moles and moles to grams
Convert moles to atoms and atoms to moles
 NEVER EVER put 6.022 x 1023 in front of
GRAMS
IX. Counting Atoms and Stuff
1 molecule C6H12O6 has 6 C atoms, 12 H
atoms, and 6 O atoms.
6 moles of C atoms,
1 mole of C6H12O6 has ___
12 moles of H atoms, and ___
6 moles of
____
O atoms.
IX. Counting Atoms and Stuff
1 mole XaYb = a moles X = b moles Y
 Ex.
1 mole Na2C2O4
= 2 moles Na
= 2 moles C
= 4 moles O
IX. Counting Atoms and Stuff
 1 mole X= 6.0221367 x 1023 atoms of X
 1 mole of X = atomic mass of X (grams)
 a moles X = 1 mole XaYb = b moles Y
IX. Counting Atoms and Stuff
atoms X
1 mole X= 6.0221367 x
1023 atoms
Moles X
1 mole of X =
atomic mass of X
grams X
molecules XaYb
Moles XaYb
grams XaYb
IX. Counting Atoms and Stuff
 Problem-solving:
 MAP the problem first (determine what
you are starting with, where you want to
end up, and the path to follow).
 Example:
What is the mass of 3.60 moles of Cu?
IX. Counting Atoms and Stuff
atoms X
START
1 mole X= 6.0221367 x
1023 atoms
Moles X
molecules XaYb
Moles XaYb
End
1 mole of X =
atomic mass of X
grams X
grams XaYb
IX. Counting Atoms and Stuff
 Map out the problem first (determine what
you are starting with, where you want to
end up, and the path to follow).
 Example:
What is the mass of 3.60 moles of Cu?
Moles of Cu  grams of Cu
3.60 moles Cu x
63.546 grams Cu
1
moles Cu
= 229 grams Cu
IX. Counting Atoms and Stuff
Map the following problems FIRST, then solve:
 How many moles are in 11.9 grams?
 How many atoms are in 3.60 x 10-10 grams
of gold?
 How many grams of sodium are in 2.34
moles of Na2CO3?
 How many grams of Fe are in 13.86 grams
of Fe2O3?
IX. Counting Atoms and Stuff
atoms X
1 mole X= 6.0221367 x
1023 atoms
Moles X
1 mole of X =
atomic mass of X
grams X
molecules XaYb
Moles XaYb
grams XaYb