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Chapter 4
Atomic Structure
Anything in black letters = write it in
your notes (‘knowts’)
4.1 – Defining the Atom
Atom - smallest particle of an element that still
has the properties of the element
comes from the Greek word atomos
which means uncuttable or indivisible
Democritus (460 B.C. – 370 B.C.)
one of the first to propose the idea of the atom;
based on pure speculation
John Dalton (~1800)
proposed 1st atomic theory
Daltons Atomic Theory (~1800) p. 103)
1. All elements are composed of tiny indivisible particles called atoms.
2. Atoms of the same element are identical. The atoms of any one
element are different from those of any other element.
3. Atoms of different elements can physically mix together or can
chemically combine in simple whole-number ratios to form compounds.
4.
Chemical reactions occur when atoms are separated from each other, joined,
or rearranged in different combinations. Atoms of one element are never
changed into atoms of another element as a result of a chemical reaction.
Atoms of
element A
Atoms of
element B
Mixture of atoms of
elements A and B
Compound made by
atoms of elements A
and B
Caveats of Dalton’s Atomic Theory
1. All elements are composed of tiny indivisible particles called atoms.
Atoms are not indivisible – they are made of subatomic particles
2. Atoms of the same element are identical. The atoms of any one
element are different from those of any other element.
Every atom has at least one isotope; one atom’s isotope is NOT
identical to another isotope of the same atom.
3. Atoms of different elements can physically mix together or can
chemically combine in simple whole-number ratios to form compounds.
This is known as the Law of Definite Proportions – very important.
4.
Chemical reactions occur when atoms are separated from each other, joined,
or rearranged in different combinations. Atoms of one element are never
changed into atoms of another element as a result of a chemical reaction.
Atoms of one element can change into an atom of another element
as a result of a nuclear reaction.
4.2 – Structure of the Nuclear Atom
Subatomic Particles
Electron – discovered in 1897 by J.J.
Thomson while experimenting with cathode ray
tubes
Thomson performed experiments that involved passing electric current
through gases at low pressure.
The result was a glowing beam, or cathode ray, that traveled from the cathode to
the anode.
Thomson found that a cathode ray is deflected by electrically charged metal plates.
Thompson knew that opposite charges attract and like charges repel, so he
hypothesized that a cathode ray is a stream of tiny negatively charged
particles moving at high speed; now called electrons.
To test his hypothesis,
Thompson set up an
experiment to measure
the ratio of an electron’s
charge to its mass.
Also, the charge-to-mass
ratio of electrons did not
depend on the kind of gas
in the cathode-ray tube or
the type of metal used for
the electrodes.
A cathode ray can also be deflected by a magnet.
Proton – observed in cathode ray tubes also in
1886 by Eugen Goldstein.
Neutron – confirmed in 1932 by James Chadwick.
Properties of Subatomic Particles
Particle
Symbol
Relative
charge
Relative mass
(mass of proton = 1)
Actual mass
(g)
Electron
e–
1–
1/1840
9.11  10–28
Proton
p+
1+
1
1.67  10–24
Neutron
n0
0
1
1.67  10–24
The Atomic Nucleus
How are atoms structured?
Democritus
Dalton
Thomson’s Plum
Pudding Model
In 1911, Ernest Rutherford and others performed
the Gold Foil Experiment to test the plum pudding
model
Ernest
Rutherford
The Gold Foil Experiment
The Gold Foil Experiment
The results…
It was expected that alpha
particles would pass through
the plum pudding model of
the gold atom undisturbed.
Expected
It was observed that a small
portion of the alpha particles
were deflected, indicating a
small, concentrated positive
charge (the nucleus!)
Actual
It was quite the most incredible event that has ever happened to me
in my life. It was almost as incredible as if you fired a 15-inch shell
at a piece of tissue paper and it came back and hit you. On
consideration, I realized that this scattering backward must be the
result of a single collision, and when I made calculations I saw that
it was impossible to get anything of that order of magnitude unless
you took a system in which the greater part of the mass of the atom
was concentrated in a minute nucleus. It was then that I had the
idea of an atom with a minute massive center, carrying a charge.
—Ernest Rutherford
Nucleus – tiny positively charge core of an atom
Rutherford’s Nuclear
Model of the Atom
Is this the current
model of the atom?
NO…
• If an atom were the size of a football stadium,
the nucleus would be about the size of a marble
ASSIGN:
Read 4.2
Lesson Check 4.2; #9-15 (page 109)
4.3 – Distinguishing Among Atoms
Atomic Number (Z) The number of protons in an atom; identifies
the element.
In a neutrally charged atom, the number of
protons (p+) equals the number of electrons (e-)
Mass Number (A) The number of protons (p+) and neutrons
(n0) in an atom.
The mass number is NOT the atomic mass.
Element
H
O
Ca
Atomic
Protons
Number (Z)
(p+)
Electrons
(e-)
Neutrons
(n0)
1
1
1
???
8
8
8
???
20
20
20
???
The number of n0 depends on the
mass number of the isotope
Isotopes Atoms of an element that have a different number
of neutrons.
Element
H
O
Ca
Atomic
Protons
Number (Z)
(p+)
Electrons
(e-)
1
1
1
8
8
8
20
20
20
Neutrons
(n0)
Zoom for detail
Chemical Symbols for Isotopes
A is the superscript
20
10
Ne
21
10
Ne
22
10
Z is the subscript
Ne
Determining the Composition of an Atom
How many protons, electrons, and neutrons
are in each atom?
9
4
a. Be
20
10
b. Ne
23
11
c. Na
Naturally Occurring Isotopes of Neon
20
10
Ne
21
10
Ne
22
10
Ne
Percent Abundance in Nature
90.48%
0.27%
9.25%
The masses of atoms are rarely expressed in grams.
The C-12 isotope has been given a mass of exactly
12 atomic mass units (amu)
The masses of all other elements are based on the
mass of the C-12 isotope.
Why is the mass of a carbon atom 12.011 amu?
Atomic Mass –
Weighted average of all the naturally occurring
isotopes of the element.
12
6
C
13
6
C
12.000 amu
13.003 amu
98.93 %
1.07 %
14
6
C
14.003 amu
0.0000000001 %
(12 x 0.9893)  (12.003 x 0.0107)  12.011
Atomic Mass of Carbon = 12.011 amu
12.000 amu
13.003 amu
98.93 %
1.07 %
14.003 amu
0.0000000001 %
Atomic Mass of Carbon = 12.011 amu
No atom of carbon actually weighs 12.011 amu. But
a typical carbon atom averages 12.011 amu.
Atomic masses are weighted averages.
There are 2 stable isotopes of silver
Silver-107; 106.905097 amu; 51.84%
Silver-109; 108.904752 amu; 48.16%
Calculate the atomic weight of silver.
Atomic Weight of Silver = 107.868 amu
Despite differences in the number of neutrons,
isotopes of an element are chemically similar.
Neutrons do not determine chemical reactivity; the
electrons do.
ASSIGN:
Lesson Check 4.3 (#26-34); p. 119
ASSIGN:
#53-59, 61, 64-67