The Atom
Atoms, Molecules, and Ions
1
A HISTORY OF THE
STRUCTURE OF THE ATOM
2
History
• Greek Philosopher Democritus (460-370
B.C.):
all matter composed of small atoms
atomos = indivisible
• 1803, John Dalton (brit.): atoms are the
fundamental building blocks of matter
3
Dalton's Postulates
Each element is composed of extremely small
particles called atoms.
4
Dalton's Postulates
All atoms of a given element are identical to one
another in mass and other properties, but the
atoms of one element are different from the
atoms of all other elements.
5
Dalton's Postulates
Atoms of an element are not
changed into atoms of a different
element by chemical reactions;
atoms are neither created nor
destroyed in chemical reactions.
6
Dalton’s Postulates
Compounds are formed when atoms of
more than one element combine; a
given compound always has the same
relative number and kind of atoms.
7
John Dalton’s Atomic Theory
(ca 1803)
1.
Each element is composed of extremely small
particles called atoms.
2.
All atoms of a given element are identical. The atoms
of different elements are different and have different
properties (including different masses).
Atoms of an element are not changed into different
types of atoms by chemical reactions. Atoms are
neither created nor destroyed in chemical reactions.
This is the Law of Conservation of Mass.
Compounds are formed when atoms of more than one
element combine. A given compound always has the
same relative number and kind of atoms. This is the
Law of Constant Composition.
3.
4.
8
John Dalton’s Atomic Theory
Led him to deduce the Law of Multiple Proportions:
When two or more elements combine to form more
than one compound, the relative masses of the
elements which combine will be in in the ratio of
small whole numbers.
In carbon monoxide, CO, 12 g carbon combine with
16 g oxygen. C:O ratio is 12:16 or 3:4.
In carbon dioxide, CO2, 12 g carbon combine with
32 g oxygen. C:O ratio is 12:32 or 3:8.
9
John Dalton’s Atomic Theory
Almost right. A good start.
very small
Structure of the atom after Dalton
(ca. 1810)
10
J.J. Thomson (1897):
Cathode Rays
Atoms subjected to high voltages
give off cathode rays.
11
J.J. Thomson: Cathode Rays
Cathode rays can be deflected by a magnetic field.
Cathode rays are negatively charged particles (electrons).
Electrons are in atoms.
12
J.J. Thomson – The Electron
“Plum pudding”
model: Negative
electrons are
embedded in a
positively charged
mass.
Unlike electrical
charges attract, and
that is what holds the
atom together.
Electrons (-)
Positively charged
mass
Structure of the atom after
Thomson (ca. 1900)
13
Radioactivity
• Radioactivity is the spontaneous emission
of radiation by an atom.
• First observed by Henri Becquerel
(1852-1908).
• Marie and Pierre Curie also studied it.
• Nobel Prize in 1903 (physics).
14
Studies of Natural Radioactivity
Some atoms naturally emit one or more of the following
types of radiation:
alpha (α) radiation (later found to be He2+ - helium nucleus)
beta (β) radiation (later found to be electrons)
gamma (γ) radiation (high energy light)
Alpha particles
α
γ
Positively charged
mass
α
Electrons (-)
γ
Somehow gamma
radiation is in
there, too.
Structure of the atom after Becquerel (early 1900s)
15
Radioactivity
• Three types of radiation were discovered by
Ernest Rutherford:
  particles (positive, charge 2+, mass 7400 times
of e-)
  particles (negative, charge 1-)
  rays (high energy light)
16
Ernest Rutherford (1910)
Scattering experiment: firing alpha particles at a gold foil
17
The Nuclear Atom
Some alpha particles
bounce off the gold foil.
This means the mass of the
atom must be concentrated
in the center and is
positively charged!
Thompson’s model could not
be correct.
18
Ernest Rutherford
The Nucleus and the Proton
The mass is not spread evenly throughout the atom, but is
concentrated in the center, the nucleus.
The positively
charged particles in
the nucleus are
protons.
Electrons (-) are now
outside the nucleus.
Structure of the atom after
Rutherford (1910)
19
James Chadwick – The Neutron
In the nucleus with the protons are particles of similar mass
but no electrical charge called neutrons.
The positively
charged particles in
the nucleus are
protons.
+ nn
Electrons (-) are now
outside the nucleus
in quantized energy
states called
orbitals. (From Niels
Bohr and quantum
mechanics)
Structure of the atom after
Chadwick (1932)
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Structure of the Atom
proton (+)
neutron
electrons responsible for the
volume and size of
the atom, negatively
charged
10-10 m
10-14
m
nucleus - responsible for the mass
of the atom, positively charged 21
Subatomic Particles
• Protons and electrons are the only particles that
have a charge.
• Protons and neutrons have essentially the same
mass.
• The mass of an electron is so small we ignore it.
22
Atomic Facts
Feature
Size
Mass
electron 10-18 m
(-)
???
0.0006
amu
10-15 m
1.0073
amu
neutron 10-15 m
(0)
1.0087
amu
proton
(+)
1 amu = 1 atomic mass
unit = 1.66054 x 10-24 g
+nn
Electrons are outside
the nucleus in
quantized energy
states called orbitals.
23
Symbols of Elements
Elements are symbolized by one or two letters.
24
Atomic Number
All atoms of the same element have the same
number of protons:
The atomic number (Z)
25
Atomic Mass
The mass of an atom in atomic mass units (amu)
is the total number of protons and neutrons in
the atom.
26
Atomic Number
• The number of protons in
the nucleus is called the
atomic number Z.
Carbon atom
• Z determines the
identity of an element.
• Saying “the atomic number of
an element is 6” is the same as
saying “carbon.”
• The number of electrons in the
atom is also Z (because atoms
have no net electric charge).
• How many neutrons are in C?
- proton
- neutron
27
Isotopes
•The number of protons and
neutrons (nucleons) in an element
is called the mass number A.
A
Z
12C
6
A = Z + number of neutrons.
•An element may have different
numbers of neutrons but NOT
different numbers of protons.
•Atoms of an element with different
numbers of neutrons are called
isotopes of that element.
- proton
- neutron
How many neutrons are in C? The answer is “it depends on the isotope.”
28
Isotopes
• Isotopes are atoms of the same element with
different masses.
• Isotopes have different numbers of neutrons.
11
C
6
12
C
6
13
C
6
14
C
6
29
Isotopes
number of number of
protons (Z) neutrons
mass
number of
number (A) electrons
6
6
12
6
8
14
8
8
16
92
146
238
6
symbol
12C
or C-12
6
6
14C
or C-14
6
8
16O
or O-16
8
92
238U
or U-238
92
30
Atomic Masses
Atomic masses are based on 12C.
The mass of 12C (or C-12) is
defined to be exactly 12 amu.
31
Atomic Masses
The mass (weight)
shown in the periodic
table is the mass of
the element as its
occurs naturally.
If the element has more
than one isotope, the
mass shown is the
weighted average of
the masses of the
isotopes.
Mg has 3 isotopes.
24Mg
78.99% 23.985 amu
25Mg 10.00% 24.986 amu
26Mg 11.01% 25.983 amu
weighted average of Mg:
0.7899x23.985 18.946
0.1000x24.986
2.499
0.1101x25.983 +2.861
24.31 amu
atomic weight of Mg based on
natural abundance: 24.31 amu
32
Ions
• Atoms can gain or lose electrons to become charged
particles called ions.
– A chemical particle that contains a positive or
negative charge
• Cations are positively charged ions.
– Formed when an atom loses electrons
• Anions are negatively charged ions.
– Formed when an atom gains electrons
33
Ions
Formation of a cation
1p
e-
Hydrogen atom
1p, 0 n, 1 e1
H
Net charge = 0
+
1p
e-
Hydrogen ion (cation)
1p, 0 n, 0 e1
H+
Net charge = +1
34
Ions
Formation of an anion
8p
8n
8e-
+
Oxygen atom
8p, 8 n, 8e16
O
Net charge = 0
2e-
8p
8n
10e-
Oxygen ion (anion)
8p, 8n, 10e16
O2-
Net charge = -2
35
Nuclear Symbols
Mass
Number
Atomic
Number
X
Charge
Charge = # p - # e36
Nuclear Symbols
• Using nuclear symbols to determine the
number of p, n, e, and total charge
16
8
O
Mass Number = 16
Atomic Number = 8
# protons = atomic number = 8
# neutrons = Mass # - Atomic # = 16 - 8 = 8
# electrons = # protons = 8
37
Nuclear Symbols
16
O
8
2-
Mass Number = 16
Atomic Number = 8
# protons = atomic number = 8
# neutrons = Mass # - Atomic # = 16 - 8 = 8
# electrons = # protons - charge = 8 - (-2) = 10
38
Nuclear Symbols
137
2+
Ba
56
Mass Number = 137
Atomic Number = 56
# protons = atomic number = 56
# neutrons = Mass # - Atomic # = 137 - 56 = 81
# electrons = # protons - charge = 56 - (+2) = 54
39
Nuclear Symbols - Atoms
Example: Write the nuclear symbol for the
following atoms:
1) 50 p, 70 n
120Sn
2) 17 e-, 20 n
37Cl
50
17
40
Nuclear Symbols - Ions
Practice writing nuclear symbols from information
given:
1) 53 p, 74 n, 54 e53 proton (= atomic number)  I
74 neutrons + 53 proton  mass number = 127
54 electrons (one more than protons)  1127 1I
53
41
2) 23 e-, 30 n, net charge = +3
# protons?
23 electrons, but charge of 3+
ie 3 more protons than electrons  p= 26
 Atomic number = 26  element = Fe
56
26
3+
Fe
42