Ch 10 Notes: Modern Atomic Theory
10.1 Rutherford’s Atom
Our understanding of atoms has changed over the past 2400 years:
It started with Democritus’ philosophical idea of the ultimately small piece of matter that cannot be
further subdivided.
Over 2000 years later Dalton’s atomic theory was useful in explaining how compounds have a
constant composition and how matter is neither created nor destroyed in chemical reactions, only rearranged.
Dalton’s atom was also indivisible.
Thompson went on to show the atom was composed of parts. He discovered the negatively-charged
electron and determined, logically, that there must also be a positively-charged component to the atom. His
model of the atom, though, was not very accurate.
Rutherford’s experiments confirmed the existence of relatively heavy positively-charged protons
located in the tiny, dense nucleus of the atom with the far less massive electrons orbiting in a large volume of
space around the nucleus. He also proposed the existence of neutrons, later discovered by Chadwick.
Rutherford’s atom (see fig 10.1) was mostly empty space and in some versions is still the popular model of
the atom we have today. However, his model did not demonstrate how the electrons could remain in stable
orbits around the nucleus. They should fall towards the protons in the nucleus atom and merge with it,
thereby collapsing the atom.
To understand the model of the atom further requires knowledge of light and how it transmits energy.
10.2
Electromagnetic Radiation (EMR)
EMR is energy transmitted from one place to another by light. The fundamental unit of light is the photon.
There are many forms of EMR: gamma rays, x-rays, UV rays, visible light, infra-red rays, microwaves,
radio waves (see fig. 10.4). Human eyes can only see a small portion of this spectrum – the visible light part.
EMR has both particle and wave properties. (The particle properties of EMR – i.e., light – are not discussed
in your text.)
The wave properties of EMR include wavelength, frequency, and speed.
Wavelength is the distance between two consecutive wave peaks. (Think of water waves as
analogous to EMR waves. (See fig 10.2, 10.3).
The frequency of the wave refers to how many wave peaks pass a given point per given time period.
The speed of a wave indicates how fast a given peak travels through its medium. For light, that speed
is 186,000 miles per second (in a vacuum).
Different wavelengths of EMR (light) carry different amounts of energy. For example, photons that
correspond to red light have less energy than photons that correspond to blue light. That’s because the
wavelength of red light is longer than that of blue light. (See fig 10.6) The longer the wavelength of light
the lower the energy carried; the shorter the wavelength the higher the amount of energy carried.
This is why we are not easily damaged by visible light rays (ROY G BIV) but can be injured by light of
shorter, and therefore more powerful; wavelengths.
UV light has enough energy to cause damage to our skin (suntan or sunburn).
X-Rays are even stronger – they can go through our bodies. Our exposure to them must be strictly
limited as they are powerful enough to damage cells and our DNA.
Gamma rays are even stronger and can seriously damage or kill organisms even with brief exposures.
But remember, even visible light, infra-red, microwaves and radio waves can cause damage if intense
enough. After all, microwaves can be used for cooking foods!
10.3
Emission of Energy by Atoms
This section simply illustrates how different atoms give off different colored light when heated. To
summarize, when atoms receive energy from a source they become “excited.” They then release this energy
by emitting light of a particular wavelength. The energy of the light emitted is equal to the energy received.
(See fig 10.8.)
10.4
The Energy Levels of Hydrogen
This section gives a specific example of hydrogen atoms releasing energy when excited. (See figs 10.9 and
10.10.) When excited hydrogen atoms release energy they do so by emitting light (photons) of only very
particular wavelengths, as shown in Fig 10.11. Hydrogen atoms never emit light with wavelengths between
these narrow bands. (Compare this spectrum with the full spectrum of visible light in fig.10.4.) Because
only certain wavelengths are emitted, only certain energy changes are taking place.
All hydrogen atoms have the same set of discreet energy levels. Therefore, the energy levels of hydrogen are
quantized, which means only certain values are allowed. If hydrogen atoms could be excited to any energy
state, then the light they would give off would be continuous.
See the analogy of the ramp versus steps (fig 10.15) that illustrates the difference between continuous and
discreet energy levels.
10.5
The Bohr Model of the Atom
How is the particular emission spectrum of hydrogen explained?
In 1913 Niels Bohr constructed a model of the hydrogen atom that had its electron moving in circular orbits
corresponding to allowed energy levels. The electron could jump to a different orbit by absorbing or
emitting a photon of light of the correct energy content (fig 10.17). When an electron jumped down to a
lower orbit after being energized it gave off a photon of light that corresponded to the observed emission
spectrum for hydrogen.
Bohr’s model worked well-enough for hydrogen, but not for other atoms. In fact, it is fundamentally flawed
because electrons do not move in orbits around a nucleus, as Bohr proposed. However, it paved he way for
later theories.
10.6
The Wave-Mechanical Model
It replaced the Bohr model of the atom.
The Bohr model simply put electrons with higher energy in an orbit with a greater radius (distance from the
nucleus). This model did not explain the behavior of atoms other than hydrogen.
In the 1920s Louis Victor de Broglie and Erwin Schroedinger proposed that electrons have both wave and
particle properties, just like photons. By mathematically treating electrons as waves, rather than just
particles, they came up with a model that works for all atoms, not just hydrogen.
Electrons, behaving as waves, do not move in orbits but are located in orbitals. Orbitals simply represent the
region of space, as mathematically defined, occupied by an electron around an atom. Electrons do not move
like planets around the sun.
Orbitals are a probability map of where electrons are likely to be found, not precise locations.
As you’ll see in Ch 10.7, and below, orbitals have different shapes that correspond to the different energy
states of electrons. Shown are the s, p, d, and f orbitals. The nucleus of the atom is located at the
intersection of the axes.
Some web sites on atomic orbitals:
http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/index.html
http://en.wikibooks.org/wiki/High_School_Chemistry/Shapes_of_Atomic_Orbitals
http://www.orbitals.com/orb/index.html
(A bit advanced, but you may want to watch)
http://www.youtube.com/watch?v=Ewf7RlVNBSA
http://www.youtube.com/watch?v=9E3QaRxqXZc