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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
cathode rays
THE STRUCTURE OF THE ATOM
 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 electron
 Plum Pudding Model/ chocolate chip cookie
THE STRUCTURE OF THE ATOM CONT’D
oCharge and Mass of the electron
oThomson and Milliken (oil drop experiment) worked together to
discover the charge and mass of the electron
 charge = 1.592
detected
 mass
× 10−19 coulombs this is the smallest charge ever
= 9.1093821545
× 10−31 g
this weight is pretty insignificant
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.
THE STRUCTURE OF THE ATOM CONT’D
Rutherford Concluded:
 1 – the positive portion of the atom is in the middle
 2 – most of the atom is empty
 3 – most of the mass is in the middle
 4 – electrons orbit the nuclues
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.
1932 – Chadwick confirmed the discovery of the neutron
IMPORTANT SUBATOMIC PARTICLES
a.m.u.
Mass, kg
Charge
Location
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
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)
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
ISOTOPE CALCULATIONS
Antimony consists of two natural isotopes 57.25% is Sb (120.9038).
Calculate the % and mass of the other isotope if the average atomic
mass is 121.8.
123.000 amu
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