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Transcript
Before you begin this chapter,
review the following concepts
and skills:
Atoms and electrons
■
modern model of the atom. You will also learn about the model itself and
how it relates to periodic trends and the periodic table.
identifying subatomic
particles and their
properties (from previous
studies)
green colours
in the northernit
lights
comeresult of streams of
The northern lights, also called theThe
aurora
borealis,
the
from the interaction of accelerated electrons
withWhen
oxygen atoms.the
The redelectrons
colours usually
protons and electrons from the Sun.
interact with
come from nitrogen atoms. How do these
describing the structure
interactions lead to atoms
of different
elements
gaseous atoms and
in organization
Earth’sof the
upper atmosphere,
the
atoms
emit colored light.
emitting light with characteristic wavelengths?
■
periodic table (from
previous studies)
■
explaining periodic trends
for properties such as
atomic radius, first
ionization energy, and
electron affinity (from
previous studies)
118 MHR • Unit 2 Structure and Properties
In the twentieth century scientists developed a revolutionary new model of
the atom. This model helped to explain many phenomena that the existing
theories had failed to explain.
n modified form, both of these inventions are
in every chemistry course around the world.
ed the periodic table, he was well-acquainted
He knew nothing, however, about subatomic
electron, which is the foundation for the
nctive shape. Because the original periodic
imental observations, chemists did not need
structure to develop it. (As you will see in
iodic table easily accommodates details about
will learn that the modern periodic table’s
l consequence of atomic structure.)
spectrum, quantum, photons
The atomic model - Dalton
In your notebook, list the
main ideas in Dalton’s atomic
theory. Explain how this theory
enabled chemists to explain
the three mass laws: the law
of conservation of mass,
the law of definite proportions,
and the law of multiple
proportions.
John Dalton’s atomic theory (1809) suggested that
the atom was a tiny, solid uncuttable sphere.
e Modern Atomic Model
omic theory to advance their understanding
uring chemical reactions. His atomic model,
explaining the behaviour of substances.
d a system of symbols to show how atoms
ances. Figure 3.2 on the next page shows
you will no doubt notice, Dalton correctly
rbon dioxide and sulfur trioxide, but ran into
mmonia, and methane. Dalton’s attempt at
hts a crucial limitation with his atomic
se it to explain why atoms of elements
ch they do. This inability did not prevent
studies. It did, however, suggest the need
omic model.
The model of the
atom in 1809. The atom, as Dalton
pictured it, was a tiny, solid,
indestructible sphere.
Figure 3.1
Dalton’s atomic theory no longer applies in its
original form
Chapter 3 Atoms, Electrons, and Periodic Trends • MHR
119
The atomic model - Dalton
Dalton’s atomic model, however, was inadequate for explaining the behavior
of substances.
He designed a system of symbols to show how atoms combine to form
other substances.
hydrogen
sulfur
nitrogen
water
carbon dioxide
sulfur trioxide
carbon
ammonia
oxygen
methane
Examples of Dalton’s system of symbols for atoms and molecules
Figure 3.2
Examples of Dalton’s system of symbols for atoms and molecules
Dalton correctly predicted the formulas for carbon dioxide and sulfur
The Discovery of the Electron Requires a New Atomic Model
trioxide,Inbut
ran into serious trouble with water, ammonia, and methane
1897, Dalton’s idea of an indivisible atom was shattered with a startling
announcement. A British scientist, Joseph John Thomson, had discovered
1
the existence of a negatively charged particle with mass less than 1000
that of a hydrogen atom. This particle was, of course, the electron.
1
(Later calculations showed that the mass of an electron is 1837
that
The atomic model - Thomson
oxygen
methane
Figure 3.2 Examples of Dalton’s system of symbols for atoms and molecules
The Discovery of the Electron Requires a New Atomic Model
In 1897, Dalton’sIn 1897,
ideaDalton’s
of an
shattered
with a startling
ideaindivisible
of an indivisibleatom
atom waswas
shattered
with a startling
A British scientist,
JosephJohn
John Thomson,
had discovered
announcement. Aannouncement.
British
scientist,
Joseph
Thomson,
had
discovered the
1
the existence of a negatively charged particle with mass less than 1000
existence of a negatively
charged
with
mass1000
that of a hydrogen
atom. Thisparticle
particle was,
of course,
the electron. times smaller than
1
(Later calculations showed that the mass of an electron is 1837
that
the mass of a hydrogen
atom.
This
particle
was
the
electron.
of a hydrogen atom.) It took several years for chemists to consider the
consequences of this discovery. They realized that if atoms contain
electrons, atoms must also contain a positive charge of some kind to
balance the negative charge. The atomic model that Thomson eventually
proposed is shown in Figure 3.3. Keep in mind that scientists had not
yet discovered the proton. Therefore, in this model, the entire sphere
carries a uniform, positive charge.
Chemists realized that if atoms contain electrons, atoms must also contain a
positive charge of some kind to balance the negative charge. This was the
atomic model that Thomson proposed.
electrons
Negative electrons
CHEM
FA C T
Scientists initially described
adioactivity solely in terms
of radiation. The idea of
adioactive particles first
appeared around the turn of
he twentieth century. In 1909,
Ernest Rutherford reported
confidently that the alpha
positively
Positively charged sphere
Figure 3.3 The atomic model in 1903. Thomson viewed the atom as a positively
charged sphere embedded with sufficient numbers of electrons to balance (neutralize)
the total charge.
The atomic model - Rutherford
•
End of XIX century
Marie and Pierre Curie discover “radioactive” elements: element
whose atoms emit positively charged particles (α-particles), negatively
charged particles (β-particles), and energy (γ-radiation)
•
From 1898 to 1907
Ernest Rutherford studies the chemistry of radioactive elements
•
In 1909
Rutherford and two students perform the gold-foil experiment and
demolish Thomson’s atomic model
rutherford-scattering simulation
The atomic model - Rutherford
gold-foil experiment, shown in Figure 3.4. On the basis of this experiment, Rutherford suggested that the deflections he and his students
observed were caused by an encounter between an alpha particle and
an intense electric field at the centre of the atom.
The gold-foil experiment
A
Hypothesis: Expected result
based on Thomson’s model
incoming
α particles
B Experiment
C
1 Radioactive sample emits
beam of α particles
2 Beam of α particles
meets gold foil
little or no
deflection
cross-section of gold foil
showing Thomson’s atoms
gold foil
5 Major deflections
are observed rarely
4 Minor deflections (exaggerated here)
are observed occasionally
Figure 3.4 The hypothesis, experiment, and results of Rutherford’s gold foil experiment.
The experimental hypothesis and design owed much to the contributions of Rutherford’s
students, Hans Geiger (of Geiger-counter fame) and Ernest Marsden.
Rutherford performed several calculations that led him to an inescapable
Actual Result
incoming
α particles
major
deflection
minor
deflection
cross-section of gold foil
composed of atoms with
tiny, massive nucleus
3 Flashes of light observed when α
particles strike screen show that
most are transmited with little
or no deflection
The atomic model - Rutherford
The gold-foil experiment
1. Rutherford aimed α-particles (+) at extremely thin metal foils.
2. A small number of the α-particles were deflected significantly
(∡>90°) by the atoms of the metal foils.
3. Following Thomson’s model, Rutherford expected the αparticles to pass through the metal atoms with low deflection.
4. Rutherford suggested that the deflections he observed were
caused by an encounter between an α-particles and an intense
electric field at the centre of the atom.
The atomic model - Rutherford
Rutherford performed several calculations that led him to an inescapable
conclusion:
“the atom is made up mainly of empty space, with a small, massive region of
concentrated charge at the centre”.
Soon afterward, the charge on this central region was determined to be
positive, and was named the atomic nucleus.
Because Rutherford’s atomic model pictures electrons in motion around an
atomic nucleus, chemists often call this the nuclear model of the
planetary atom.
The atomic model - Bohr
in 1913 Niels Bohr, a Danish physicist and student of Rutherford, proposed a new model for
the hydrogen atom.
Bohr’s atomic model pictures electrons in discrete orbit around a central
nucleus, unlike Rutherford’s model, in which electrons may move anywhere within the
space around the nucleus.
The Bohr Atomic Model
1. Each atom has specific energy levels, corresponding to the atom’s
electrons orbits.
2. While in one of its energy levels, electrons do not emit energy.
3. Electron change energy level by emitting or absorbing a specific quantity
of energy.
that the frequency symbol, ν,
is the Greek letter nu, not the
letter v.)
λ
8s
(8 Hz)
The atomic model - The Problem of Atomic Spectra
C The amplitude (
wave
λ
wavelength (nm)
10−2
100
102
104
16 s−1 (16 Hz)
106
108
1010
A Despite differences in
wave represent
A wave with a h
amplitude
has a
1012
intensity (is brig
wave with a low
visible
ultra-wavelengths and frequency,
gamma
Infrared
infrared
x-ray
violet
microwave
radio frequency
ray
all electromagnetic radiation
rays
travelsFigure
at the3.6
sameThe
speed
electromagnetic spectrum and its properties
in a vacuum: 3.00 × 10−8 m/s.
12
1010
1020
1018
1016
1014
108
106
104
10
1012
This value, the speed of light,
−1
The visible
portion
is a constant
represented
by of the electromagnetic spectrum is called a continuous frequency (s )
the symbol
c.
spectrum,
because the component colours are indistinct. They appear
400
500
600
700
750 nm
“smeared” together into a continuum of colour. According to nineteenthcentury physics, part of the energy emitted by electrons should be
observable as a continuous spectrum. This is not, however, the case.
Instead, when atoms absorb energy (for example, when they are exposed
visible light region
to an electric current), you observe a pattern of discrete (distinct),
coloured lines separated by spaces of varying length. See Figure 3.7. You
The
relationship
between
B
higher
lower by
When atoms absorb
energy,
we this
canline
observe
a pattern
of
distinct,
linesinseparated
frequency
(ν) in
s−1and
(Hz) forcolored
can
also
observe
spectrum
for
hydrogen,
other atoms,
frequency (symbol ν, measured
amplitude
amplitude
spaces of varying
length.
(brighter)
(dimmer)
1 second
in Hz)Investigation
and wavelength
(symbol λ,
3-A.
usually measured in nm or m,
wavelength
λ
depending
As
400 on the energy). 500
600
750 nm
4 s−1 (4 Hz)
wavelength increases, frequency
decreases. As wavelength
decreases, frequency increases.
3.7= 1The
This1spectrum
characteristic
of lines
H atoms.
Nospectrum
other atoms
display this pattern
coloured
of this
are characteristic
of
(Note Figure
that
Hz
s−1is
.discrete,
Note
also
−1 (8 Hz)
8
s
λ
that the
frequency
symbol,
ν,
of coloured
lines.
hydrogen
atoms.
No other atoms display this pattern of coloured lines.
wavelength, λ
is the Greek letter nu, not the
letter v.)
C The amplitude (height) of a
wave represents its intensity.
16 s−1 (16 Hz)
λ
A wave with a higher
amplitude has a greater
intensity (is brighter) than a
The atomic model - The Problem of quanta
In 1900, a physicist named Max Planck suggested that matter, at the atomic
level, can absorb or emit only discrete quantities of energy. Each of these specific
quantities is called a quantum of energy.
The era of quantum physics was beginning.
A quantum is an extremely small “packet” of energy.
quantum
n=4
n=3
n=2
n=1
Energy
levels
The atomic model - The Problem of photons
region
In 1905, Albert Einstein theorized that light comesvisible
in energylight
packets
(quanta) which are called photons that have particle-like properties. Light is
emitted as photons of energy.
ured
mbol λ,
m,
As
uency
ases.
e also
ν,
he
frequency (ν) in s−1 (Hz)
1 second
wavelength
λ
4 s−1 (4 Hz)
ph
λ
8 s−1 (8 Hz)
How Bohr’s Atomic Model Explains the Spectrum
Bohr proposed that the energy that is emitted and absorbed by an atom
must have specific values.
He was applying the new ideas of quanta in his model.
The Bohr Atomic Model
1. Each atom has specific energy levels, corresponding to the atom’s
electrons orbits.
2. While in one of its energy levels, electrons do not emit energy.
3. Electron change energy level by emitting or absorbing a specific
quantity of energy (quanta)
century physics, part of the energy emitted by electrons should be
observable as a continuous spectrum. This is not, however, the case.
Instead, when atoms absorb energy (for example, when they are exposed
againyou
theobserve
spectrum
H, the(distinct),
energy that is associated with
toAnalyzing
an electric current),
a patternfor
of discrete
coloured lines separated by spaces of varying length. See Figure 3.7. You
colored lines in this spectrum corresponds to the change in energy of
can also observe this line spectrum for hydrogen, and for other atoms, in
electron 3-A.
as it moves to higher or lower energy levels.
Investigation
How Bohr’s Atomic Model Explains the Spectrum
400
500
600
the
an
750 nm
The discrete, coloured lines of this spectrum are characteristic of
hydrogen atoms. No other atoms display this pattern of coloured lines.
Figure 3.7
When a H atom is exposed to an electromagnetic
radiation, its electron absorbs photons.
en=2
n=3
n=4
n=5
n=6
The atom is now said to be in an excited state.
Chapter 3 Atoms, Electrons, and Periodic Trends • MHR 123
When an electron is excited to a higher energy level, it
falls back to the original energy level and emits light
of certain energy.
E.g., an electron that falls from the third energy to the
second energy level emits a photon of red light with a
λ of 656 nm.
spectra are
called
emission
spectra.
century physics,
partalso
of the
energy
emitted
by electrons should be
3.9 shows
the energy
transitions
that are the
responsible
observable asFigure
a continuous
spectrum.
This
is not, however,
case. for the
coloured
lines
in hydrogen’s
emission
spectrum.
Notice
use of the
Instead, when
atoms
absorb
energy (for
example,
when they
are the
exposed
symbol
n to designate
the allowed
energy
levels
for lines
the hydrogen
atom: emission spectrum.
energy
transitions
responsible
for
colored
in hydrogen’s
toThe
an electric
current),
youare
observe
a pattern
ofthe
discrete
(distinct),
n = 1, n = 2, and so on. This symbol, n, represents a positive integer (such
coloured
lines “n”
separated
by spaces
of varying
length.
Seefor
Figure
3.7. You
The symbol
the possible
energy
levels
an hydrogen
as 1, 2,
ordesignate
3), and is called
a quantum
number.
You will
learn moreatom:
aboutn=1, n=2, and so on.
can also observe this line spectrum for hydrogen, and for other atoms, in
the significance of quantum numbers in section 3.2.
Investigation
3-A.
This symbol,
“n” is called a quantum number.
How Bohr’s Atomic Model Explains the Spectrum
400
400
500
600
600
500
750 nm
750 nm
n=6
n=5
n=4
The discrete, coloured lines of this spectrum are characteristic of
hydrogen atoms. No other atoms display this pattern of coloured lines.
Figure 3.7
n=3
e-n = 1
n=2
n=3
n=4
n=2
n=3
n=4
n=5
n=6
A Note: Orbits are not
n=5drawn to scale.
n=6
n=2
123
energy
Chapter 3 Atoms, Electrons, and Periodic Trends • MHR
n=1
InBitsNote:
unexcited
hydrogen‘s
electron
lays inelectron
the first is in
In itsstate,
unexcited
state,
hydrogen‘s
energy
Thislevel
is theclosest
lowest-possible
energy n
level,
thelevel.
energy
to the nucleus:
= 1. This is
representing
a
state
of
greatest
stability
for
the
hydrogen
the lowest-possible energy level, representing a atom
state
of greatest stability for the hydrogen atom.