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Chapter 3 Notes
The Atom: From Idea to Theory
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
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
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 Cont’d
H
+
O
H2O
Wt.
of H
2
H
+
O
H2O2
2
Wt.
of O
16
32
The ratio of O in H2O2 to O in H2O = _______2 to 1______________
Dalton’s Atomic Theory
1. Everything is made of atoms.
2. Atoms of the same element are
identical. (NOT)
3. Atoms can not be broken down,
created, or destroyed. (NOT)
4. Atoms combine in simple whole
number ratios to form chemical
compounds.
5. A chemical reaction is the
combining, separation, or
rearrangement of atoms.
2:1
1:1
C (s) + O2 (g) --------> CO2 (g)
3.2 The Structure of the Atom
Updating Atomic Theory
 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 electrons
 20 years later, J.J. Thomson (English) repeated those
experiments and devised new ones. Thomson used a variety of
materials, so he figured cathode ray particles must be
fundamental to all atoms. 1897 - discovery of the positive
charge
 Plum Pudding Model
The Structure of the Atom Cont’d
o Charge and Mass of the electron
o Thomson and Milliken (oil drop experiment) worked
together to discover the charge and mass of the electron
 charge = 1.592
ever detected
× 10−19 coulombs
 mass = 9.1093821545
insignificant
× 10−31 g
this is the smallest charge
this weight is pretty
The Structure of the Atom Cont’d
1909 - Gold Foil Experiment (Rutherford New Zealand)
Nuclei are composed of ‘nucleons’:
protons and neutrons
Top: Expected results: alpha (+)
particles passing through the plum
pudding model of the atom
undisturbed.
Bottom: Observed results: a small
portion of the particles were
deflected, indicating a small,
concentrated positive charge.
Important Subatomic Particles
a.m.u.
Mass, kg
Charge
Locatio
n
Proton
(p+)
1
1.67265×10-27
Neutron
(n°)
1
1.67495×10-27
0
nucleus
Stabilize
atom
Electron
(e-)
0
9.10953×10-31
-1
clouds
Bonding
+1
Job
ID
nucleus
Important Subatomic Particles Cont’d
Electrostatic force - pulls nuclei apart:
protons and neutrons
Strong Nuclear Force- force holds
nuclei together
Weighing and Counting Atoms
 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 Atoms
 Atomic Number (Z)
 Number of protons in the nucleus
 Uniquely labels each element
 Mass Number (M)
 Number of protons + neutrons in the nucleus
Weighing and Counting Atoms
 Counting electrons
 Atoms
 Same number of electrons and protons
 Ions – lost or gained electrons
 Ionic charge (q) = #protons - #electrons
 Positive ions are cations
 Negative ions are anions
Weighing and Counting Atoms
If the mass # comes
from the p+ and n0 [each
with masses of exactly
1], why don’t the atomic
weights/masses of the
all elements turn out to
be whole numbers?
Because the atomic
weights/masses on the PTable are the “weighted
averages,” of the naturally
occurring isotopes of the
element. (remember:
ignore the mass of the e-,
it’s too small to care about.
Review of Formulas
 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
rounds to 4
example - # of neutrons in Li = 6.941-3 = 3.941
 Ion – a charged atom. Atoms become charged by gaining electrons
(become a negative charge) or losing electrons (become a positive charge)
Lets 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
 Two atoms of the same element (same # of p+) but with different
weights (different # of n0)
Average Atomic Mass (“weighted average”)
 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 12 - C is the standard upon which all other masses are
based.
1/12 of the mass of 1 atom of 12 - C = 1 amu
Isotope Calculations
 Carbon consists of two isotopes: 98.90% is C-12 (12.0000 amu).
The rest is C-12 (13.0034 amu). Calculate the average atomic
mass of carbon to 5 significant figures.
 12.011 amu
 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.
 35.46
 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.
The Mole, Avogadro’s number and Molar
Mass
The Mole
Atoms are tiny, so we count them in “bunches”.
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.
1 mole = 1 gram formula mass = 6.02 x 1023 particles
The Mole, Avogadro’s number and Molar
Mass
 Molar Mass
 Molar Mass - the sum of the atomic masses of all atoms in a
formula
 Round to the nearest tenth! (measured in amu or grams)
 ex - H2
H2O
Ca(OH)2
2.0g
18.0 g
74.1 g
The Mole, Avogadro’s number and Molar
Mass
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:
Formula mass (ionic compounds)
Molecular mass (covalent compounds and
diatomic elements)
Atomic weight, Atomic mass, grams formula
weight, etc.
The Mole, Avogadro’s number and Molar
Mass
 Examples: 1 mole Na = 6.02 x 1023 atoms = 23.0 g
1 mole O2 = 6.02 x 1023 molecules =
1 mole HCl = 6.02 x 1023 molecules =
1 mole NaCl =6.02 x 1023 formula units=
Mole Map
Liters
22.4 L
Mole
Grams
Atoms,
molecules,
particles
Examples
2 steppers
convert 13.8 g Li to moles
convert 2.0 moles Ne to g
convert 3.0 moles of Be to atoms
convert 44.8 L of O2 to moles
Examples
3 and 4 steppers
convert 1.2 x 1024 atoms of Magnesium to grams
convert 128 g of O2 to molecules of O2
convert 128 g of O2 to atoms of oxygen
Convert 100. g of Ar to liters of Ar