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Chapter 2
!
Atoms and Ions
A History of Atomic Models
400 B.C.E.
atomists
1804 C.E.
Dalton
Early Philosophy of Matter
Some early philosophers believed that matter had !
ultimate, tiny, indivisible particles!
!
Leucippus and Democritus!
!
Other philosophers believed that matter was infinitely divisible!
!
Plato and Aristotle!
!
Because there was no experimental way of proving who!
was correct, the best debater was the person assumed!
correct, i.e., Aristotle!
Scientific Revolution
In the late 17th century, the scientific approach to understanding
nature became established.!
!
For the next 150+ years, observations about nature were made that
could not easily be explained by the infinitely divisible matter
concept.!
Law of Conservation of Mass
Antoine Lavoisier
1743-1794
In a chemical reaction, matter is neither created nor destroyed.
The total mass of the materials you have before the reaction must
equal the total mass of the materials you have at the end.
total mass of reactants = total mass of products
Law of Conservation of Mass
7.7 g Na
+ 11.9 g Cl2
19.6 g NaCl
Law of Definite Proportions
Joseph Proust
1754-1826
All samples of a given compound, regardless
of their source or how they were prepared,
have the same proportions of their
constituent elements.
Proportions of Na and Cl in Sodium Chloride
A 100.0 g sample of sodium
chloride contains 39.3 g of
sodium and 60.7 g of
chlorine
A 200.0 g sample of sodium
chloride contains 78.6 g of
sodium and 121.4 g of
chlorine
A 58.44 g sample of sodium
chloride contains 22.99 g of
sodium and 35.44 g of
chlorine
Law of Multiple Proportions
John Dalton
1766-1844
When two elements (call them A and B) form
two different compounds, the masses of B
that combine with 1 g of A can be expressed as
a ratio of small, whole numbers.
Law of Multiple Proportions
Carbon monoxide contains
1.33 g of oxygen for
every 1.00 g of carbon.!
1.33/1.00 = 4/3
Carbon dioxide contains 2.67 g
of oxygen for every 1.00 g
of carbon.!
2.67/1.00 = 8/3
Dalton’s Atomic Theory
1. Each element is composed of tiny,
indestructible particles called atoms.
2. All atoms of a given element have the same
mass and other properties that distinguish
them from atoms of other elements.
3. Atoms combine in simple, whole-number
ratios to form molecules of compounds.
4. In a chemical reaction, atoms are reorganized
in the way they are bound together.
5. Atoms of one element cannot change into
atoms of another element.
Some Notes on Charge
Two kinds of charge called +
and –!
Opposite charges attract!
+ attracted to –
Like charges repel!
+ repels +
– repels –
To be neutral, something must
have no charge or equal
amounts of opposite
charges
J.J. Thomson and Cathode Rays
+++++++++++
Cathode
(+)
(-)
-------------
-
Power Supply
+
Examined the electrically charge “particles” in cathode rays
Measured the amount of force necessary to deflect the paths of the particles
!
The particles have a negative charge.
The amount of deflection was related to two factors, the charge and mass of
the particles.
!
Every material tested contained these same particles.
Thomson’s Conclusions
If the particle has the same amount of charge as a hydrogen
ion, then it must have a mass almost 2000 x smaller than
hydrogen atoms!
!
The only way for this to be true is if these particles were
pieces of atoms.
!
Apparently, the atom is not unbreakable.
!
Thomson believed that these particles were the ultimate
building blocks of matter.
!
These cathode ray particles became known as electrons.
A New Theory of the Atom
Thomson proposed that instead of being a hard, marblelike unbreakable sphere, the way Dalton described it,
the atom actually had an inner structure.
!
The structure of the atom contains many negatively
charged electrons. These electrons are held in the
atom by their attraction for a positively charged
electric field within the atom.
A History of Atomic Models
400 B.C.E.
atomists
1804 C.E.
Dalton
1903
Thompson
Predictions of the Plum Pudding Atom
The mass of the atom is due to the mass of the electrons within
it. Electrons are the only particles in Plum Pudding atoms,
therefore the only source of mass.
!
The atom is mostly empty space and should not have a bunch of
negatively charged particles near each other as they would
repel.
Electron
Sphere of
charge positive
Radioactivity
In the late 1800s, Henri Becquerel and Marie Curie
discovered that certain elements would constantly emit
small, energetic particles and rays.
!
These energetic particles could penetrate matter.
Ernest Rutherford discovered that there were three
different kinds of emissions:
!
alpha rays made of particles with a mass 4 x H atom and + charge
!
beta rays made of particles with a mass ~1/2000th H atom and – charge
!
gamma rays that are energy rays, not particles
Explaining Results of α-particle Scattering Experiments
Rutherford’s expectation!
!
Small, positively charged α-particles should
pass through the nebulous, positively charged
cloud of the Thomson atomic model largely
undeflected. Some would be slightly deflected
by passing near electrons (present to
neutralize the positive charge of the cloud).
Explaining Results of α-particle Scattering Experiments
Actual Result
Rutherford’s expectation!
!
Small, positively charged α-particles should
pass through the nebulous, positively charged
cloud of the Thomson atomic model largely
undeflected. Some would be slightly deflected
by passing near electrons (present to
neutralize the positive charge of the cloud).
To minimize alpha loss by scattering from air
molecules, the experiment was carried out in a
fairly good vacuum, the metal box being
evacuated through a tube T (see below). The
alphas came from a few milligrams of radium
(to be precise, its decay product radon 222) at R
in the figure below, from the original paper,
which goes on: "By means of a diaphragm
placed at D, a pencil of alpha particles was
directed normally on to the scattering foil F. By
rotating the microscope [M] the alpha particles
scattered in different directions could be
observed on the screen S." Actually, this was
more difficult than it sounds. A single
alpha caused a slight fluorescence
on the zinc sulphide screen S at
the end of the microscope. This
could only be reliably seen by
dark-adapted eyes (after half an
hour in complete darkness) and
one person could only count the
flashes accurately for one minute
before needing a break, and counts
above 90 per minute were too fast for reliability.
The experiment accumulated data from
hundreds of thousands of flashes.
Rutherford's partner in the initial phase of this work was Hans Geiger, who later developed the Geiger
counter to detect and count fast particles. Many hours of staring at the tiny zinc sulphide screen in the dark
must have focused his mind on finding a better way!
!
In 1909, an undergraduate, Ernest Marsden, was being trained by Geiger. To quote Rutherford (a lecture
he gave much later):
"I had observed the scattering of alpha-particles, and Dr. Geiger in my laboratory had examined it in
detail. He found, in thin pieces of heavy metal, that the scattering was usually small, of the order of one
degree. One day Geiger came to me and said, "Don't you think that young Marsden, whom I am training in
radioactive methods, ought to begin a small research?" Now I had thought that, too, so I said, " Why not
let him see if any alpha-particles can be scattered through a large angle?" I may tell you in confidence that
I did not believe that they would be, since we knew the alpha-particle was a very fast, massive particle
with a great deal of energy, and you could show that if the scattering was due to the accumulated effect of
a number of small scatterings, the chance of an alpha-particle's being scattered backward was very small.
Then I remember two or three days later Geiger coming to me in great excitement and saying "We have
" It was quite the
most incredible event that 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."
been able to get some of the alpha-particles coming backward …
Rutherford’s Results
Over 98% of the α particles went straight through
!
About 2% of the α particles went through but were
deflected by large angles
!
About 0.005% of the α particles bounced off the
gold foil
!
“...as if you fired a 15” cannon shell at a piece of tissue
paper and it came back and hit you.”
Rutherford’s Interpretation – the Nuclear Model
The atom contains a tiny dense center called the nucleus.
The amount of space taken by the nucleus is only about 1/109
the volume of the atom.
!
The nucleus has essentially the entire mass of the atom.
The electrons weigh so little they give practically no mass to the
atom.
!
The nucleus is positively charged.
The amount of positive charge balances the negative charge of
the electrons.
!
The electrons are dispersed in the empty space of the
atom surrounding the nucleus.
Structure of the Nucleus
Rutherford proposed that the nucleus had a particle
that had the same amount of charge as an
electron but opposite sign – the proton.
Because protons and electrons have the same amount of
charge, for the atom to be neutral there must be
equal numbers of protons and electrons.
Some Problems
Electrons - negatively charged particles found in all atoms.
Cathode rays are made of streams of electrons.
The electron has a charge of −1.60 x 1019 C.
The electron has a mass of 9.1 x 10−28 g.
Protons - subatomic particles found in the nucleus
!
The proton has a charge of +1.60 x 1019 C.
The proton has a mass of 1.67262 x 10−24 g.
How could beryllium have four protons stuck together in
the nucleus?
If a beryllium atom has four protons, then it should weigh 4
amu; but it actually weighs 9.01 amu! Where is the extra
mass coming from?
There Must Be Something Else!
To answer these questions, Rutherford and Chadwick
proposed that there was another particle in the
nucleus – it is called a neutron.
!
Neutrons are subatomic particles with a mass =
1.67493 x 10−24 g and no charge, and are found in the
nucleus
400 B.C.E.
atomists
1804 C.E.
Dalton
1903
Thompson
A History of Atomic Models
1932
Chadwick
1913
Bohr
1911
Rutherford
Subatomic
Particle
Mass
(g)
Mass
(amu)
Charge
Proton
1.67262 x
10
1.00727
+1
Electron
0.00091 x
10
0.00055
-1
e, e
Neutron
1.67493 x
10
1.00866
0
n, n
Symbol
p, p
1 amu = 1 atomic mass unit = 1/12
the mass of a carbon-12 atom
Quantum Model
The Nuclear Model
approximately 10-10 m
approximately 10-14 m
neutron
proton
Nucleus
nucleus
electrons
Atom
The Number of Protons Defines the Element
Helium nucleus:
2 protons, 2 neutrons
Carbon nucleus:
6 protons, 6 neutrons
Structure of the Nucleus
Soddy (1913) discovered that the same element could have atoms with
different masses, which he called isotopes.
!
The observed atomic mass is a weighted average of the weights of
all the naturally occurring atoms.
!
The percentage of an element that is one isotope is called the isotope’s
natural abundance.
Atomic Mass (Weight)
Carbon
98.99 % C-12
1.11 % C-13
trace % C-14
}
Weighted average = 12.01 amu
Chlorine
75.5 % Cl-35
24.5 % Cl 37
}
Weighted average = 35.45 amu
Isotopes
All isotopes of an element are chemically identical and undergo the
exact same chemical reactions.
!
All isotopes of an element have the same number of protons.
!
Isotopes of an element have different masses.
!
Isotopes of an element have different numbers of neutrons.
!
Isotopes are identified by their mass numbers, which is the sum of
all the protons and neutrons in the nucleus.
Isotopes
Atomic number
A
X
Z
Number of protons = Z
!
Mass Number
!
Protons + neutrons = whole number = A
!
Relative Abundance = relative amount found in a sample
# neutrons =
A-Z
a)
18e-
b)
18p+
10n0
25p+
30n0
38
d)
6p+
7n0
13
6C
e)
40e-
47e47p+
62n0
55
25Mn
Ar
18
6e-
25e-
c)
109
47 Ag
f)
28e-
40p+
50n0
28p+
33n0
90
40 Zr
61
28
Ni
Complete the table:
6
42
54
13
14
55
78
6
13
96
42 Mo
42
13
55
13
13
6C
27
27
13 Al
133
55 Cs
Ions
Charged Atoms
Ions are atoms which have acquired a charge
through loss or gain of electrons.
!
When atoms gain electrons, they become negatively
charged ions, called anions.
!
When atoms lose electrons, they become positively
charged ions, called cations.
a)
10e-
b)
8p+
9n0
17
8
d)
9p+
10n0
35p+
44n0
79
Br35
e)
10e7p+
8n0
15
3N
7
18e20p+
20n0
19
9 F-
O2-
36e-
10e-
c)
40
2+
Ca
20
f)
36e37p+
48n0
85
+
Rb
37
Complete the table
16
12
13
35
210
13
35
10
S2Mg2+
3+
3+
Al
Br-
Periodic Trend in Ion Formation
Nonmetals form anions.
!
Anions named by changing the ending of the name to -ide.
!
F + 1e−
fluorine
atom
F−
fluoride
ion
O + 2e−
oxygen
atom
O2−
oxide
ion
!
!
Periodic Trend in Ion Formation
Metals form cations.
!
Cations are named the same as the metal.
!
Na
sodium
atom
Na+ + 1e−
sodium
ion
Ca
calcium
atom
Ca2+ + 2e−
calcium
ion
!
Nonmetals
Metals