Name: ______________________
Chemistry 1 Notes,
Topic 2 – Atomic Theory
Topics:
1.1.1 – Analyze the structure of atoms, isotopes and ions.
1.1.2 – Analyze an atom in terms of the location of electrons.
1.1.3 – Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model.
I.
Historical Development of Atomic Theory
EARLY CONTRIBUTIONS: DEMOCRITUS & DALTON
♦
By 400 BC, a Greek philosopher named Democritus had postulated that there
must be some basic unit of matter that could not be divided any further. He
called this basic unit an atomon (from the Greek, meaning indivisible units).
From this term we derive the English word atom.
Democritus
 There was disagreement over this idea for the next 2200 years, until the invention of the chemical
balance – the tool needed to study composition of pure substances quantitatively.
 Watershed event – 1803 – John Dalton postulated his atomic theory:
1) All matter is composed of extremely small particles called atoms.
2) Atoms of a given element are identical in size, mass, and other properties;
atoms of different elements differ in size, mass, and other properties.
3) Atoms cannot be subdivided, created, or destroyed.
4) Atoms of different elements can combine in simple, whole-numbered ratios
to form chemical compounds.
5) In chemical reactions, atoms are combined, separated, or rearranged.
John Dalton
 We now know there are exceptions to Dalton’s theory:
 Atoms can be subdivided (fission, radioactive decay), created (fusion, etc.) and destroyed (all of
these are complicated process, sometimes requiring much energy).
 Several different types of atoms for a given element can exist, all with different masses. These
are called isotopes).
 Still, Dalton’s theory = cornerstone in thinking about chemistry.
1
Name: ______________________
Chemistry 1 Notes,
EXPERIMENTS – THE ELECTRON
♦
CRT (Cathode Ray Tube) Experiments – Sir Joseph John Thomson
 In 1897, J.J. Thomson built a cathode ray tube, which generated a steady
stream of electrons. These electrons traveled down the glass tube, causing a
fluorescent glow which could be easily seen and measured. Thomson
concluded several things from his experiments:
 Thomson placed two oppositely charged plates on either side of the CRT
J.J. Thomson
(by hooking the plates up to a strong battery). When the electron beam
passed between the two plates, it was deflected toward the positive plate.
Thus, Thomson reasoned that electrons must be negative.
 Thomson also placed north and south magnetic poles on opposite sides of the CRT. When
the electron beam (travelling perpendicular to the poles) passed between the magnetic poles,
it was deflected in a particular direction. (Positively charged particles are deflected in the
opposite direction.)
 When the oppositely charged plates and magnetic poles were used simultaneously, Thomson
could cause the electron beam to continue travelling in a straight line.
 Thomson could not quantify the exact mass or charge (only that it was negative) of an
electron, but he could measure how much it was deflected in a magnetic field. Thus,
Thomson was able to calculate the mass to charge ratio (m/e) for an electron.
 Big picture: The 2 most important things Thomson discovered in the CRT Experiments:
 The electron is negatively charged.
 The mass to charge ratio (m/e) for an electron.
♦
An animation of Thomson’s experiments:
http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/animations_center.html#
2
Name: ______________________
Chemistry 1 Notes,
♦
Oil-Drop Experiment – Robert Millikan
 This led to the determination of actual mass and charge of the electron.
Robert Millikan
 Between 1908 and 1917, Robert Millikan measured the charge on an electron with the apparatus
shown above. In these experiments, the atomizer from a perfume bottle was used to spray oil
droplets into a sample chamber. Some of these droplets fell through a pinhole between two plates
of an electric field, where they could be observed through a microscope.
A source of X-rays was then used to ionize the air in the chamber by removing electrons from the
molecules in the air. Droplets that did not capture one of these electrons fell to the bottom of the
chamber due to the force of gravity. Droplets that captured one or more electrons were attracted
to the positive plate at the top of the viewing chamber and either fell more slowly or rose toward
the top.
By carefully studying individual droplets, Millikan was able to show that the charge on a drop
was always an integral multiple of a small, but finite value. When his data are converted to SI
units, the charge on a drop is always some multiple of 1.59 x 10–19 C. Combining this value for
the charge on a single electron with the mass to charge ratio (m/e) for the electron confirms
Thomson's hypothesis. The mass of an electron is at least 1000 times smaller than the lightest
atom.

Good animations of Millikan’s experiment:
http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/004_MILLIKANOIL.MOV
http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/animations_center.html#
3
Name: ______________________
Chemistry 1 Notes,
EXPERIMENTS – THE NUCLEUS (PROTONS & NEUTRONS)
♦
Gold Foil Experiment – Ernest Rutherford
 Rutherford’s experiments led to discovery of the atomic nucleus, and that it
is small, dense and positively charged.
 Fast-moving positively-charged particles (-particles) were shot through
a thin layer of gold foil, only a few atoms thick.
 Most passed straight through the gold foil in a straight line. Some were
Ernest Rutherford
slightly deflected from their straight-line course. But, roughly 1 in 8000
ricocheted back toward the source.
 Rutherford recognized that these -particles must be repelled by a very dense object in the
middle of the atom, which takes up a very small amount of space. Thus, atoms must be
mostly empty space. The nucleus must also have a positive charge (because like charges
repel).
 Rutherford concluded that the atomic nucleus was the positively charged, dense central
portion of the atom that contains nearly all of its mass, but takes up only an insignificant
fraction of its volume
 Another very good animation of Rutherford’s experiment:
http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/006_RUTHERFORD.MOV
 A good animation of Rutherford’s gold foil experiment:
http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/animations_center.html#
4
Name: ______________________
Chemistry 1 Notes,
♦
Beryllium–Wax Experiments – James Chadwick
 Following Rutherford’s experiments, scientists knew that there was
“extra” mass somewhere in the atom (they knew electrons have
essentially no mass, and the mass of an atom was greater than the
apparent number of protons in the atom). But scientists couldn’t
precisely characterize the source of this mass. In the early 1930s,
James Chadwick smashed alpha particles into beryllium, causing it to
release a different type of radiation (neutrons), which hit another
target: paraffin wax. When the beryllium radiation hit hydrogen
James Chadwick
atoms (protons) in the wax, the atoms (protons) were dislodged and
sent into a detecting chamber. In physics, it is known that only a particle having almost the same
mass as a hydrogen atom could affect hydrogen in that manner. Chadwick’s experiment showed
that a collision with beryllium atoms would release these neutral particles, which Chadwick
named neutrons. This provided the answer for hidden mass in atoms. Chadwick published his
findings in the journal Nature in 1932.
Chadwick proposed the following reaction, where 01 n represents the "new" particle, the neutron:
This is the detector used
by James Chadwick.
5
II.
Atomic Arrangement
Name: ______________________
mass
charge
Chemistry 1 Notes,
1
+1
proton
1
0
neutron
–1
electron 1/2000
 Within the nucleus are protons and neutrons:
 Protons (p+) have a positive charge equal to the negative charge of an electron.
Protons are about 2000x heavier than electrons.
 Neutrons are electrically neutral particles with effectively the same mass as protons.
 Electrons (e–) are negatively-charged subatomic particles that orbit the nucleus of an atom in
“clouds”. We need to discuss this electron cloud in terms of probability model of where the
electron(s) are most likely to be found at any given point in time…
 atomic number (Z) – the number of protons in the nucleus of each atom of that element
 # of protons = # of electrons (since all atoms are electrically neutral)
 mass number (A) – the total number of protons + neutrons in the nucleus of an isotope
 # of neutrons = mass # – atomic #
 Isotopes are atoms of the same element that have different masses.
 Isotopes have different masses due to different #s of neutrons.
There are three isotopes of hydrogen:
protium (99.985%), deuterium (0.015%), and tritium (very rare, radioactive)
6
Name: ______________________
Chemistry 1 Notes,
 2 ways to name isotopes:
mass#
1) nuclear symbol: atomic# symbol
ex:
=
A
ZE
238
92 U
2) hyphen notation: element–mass #
ex: uranium–238 or U–238
carbon–12 or C–12
 So, the three isotopes of hydrogen could be named:
 protium: 11 H or hydrogen–1
 deuterium:
 tritium:
3
1H
2
1H
or hydrogen–2
or hydrogen–3
 atomic mass unit (u) – 1/12 the mass of a carbon-12 atom
 relative atomic mass – mass of an atom expressed in atomic mass units
(it’s relative to the mass of carbon-12)
 Non-integer atomic masses (average atomic masses) result from the existence of several
different isotopes of that element.
 average atomic mass – the weighted average of the masses of the naturally occurring
isotopes of an element
 For example, there are 2 main isotopes of chlorine:
75% of all chlorine is
35
17 Cl ,
35
17 Cl
&
with a mass of 35u and 25% is
37
17 Cl
37
17 Cl
, with a mass of 37u
So the average atomic mass for chlorine is (.75 x 35u) + (.25 x 37u) = 35.5u
 Take carbon as a second example:
carbon-12 = 12.0u
carbon-13 = 13.003355u
98.90%, in nature
1.10%, in nature
average atomic mass for carbon = (0.9890 x 12.0u) + (0.0110 x 13.003355u) = 12.011 u
7
Name: ______________________
Chemistry 1 Notes,
III. Electrons and Light
WAVE-PARTICLE NATURE OF LIGHT AND ELECTRONS
 At the time of Rutherford (~1900), electrons were pictured as particles, and light was pictured as
waves. But this doesn’t explain all properties of electrons or light.
 In the early 1900’s the wave-particle theory was proposed to fully explain the properties of
electrons.
 Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through
space. Shown below is the full electromagnetic spectrum, arranged according to decreasing
wavelength. It is a continuous spectrum, showing all wavelengths of the full electromagnetic
spectrum. Note wavelength () and frequency () are inversely proportional.
 There is also an inset showing the continuous spectrum of visible light (ROY G. BIV, like a
rainbow).
 All electromagnetic radiation travels at 3.0 x 108 m/s (the speed of light, c).
 c =  (speed of light = wavelength x frequency)
 Again, note that wavelength and frequency are inversely proportional. Since their product (c)
is a constant, if one goes up, the other must go down.
 A line spectrum shows only certain, specific wavelengths of light (like a fingerprint).
 Every element has its own unique line spectrum.
 That is how we can observe starlight through a spectrophotometer, look at the wavelengths of
light, and infer what elements are contained in that star.
8
Name: ______________________
Chemistry 1 Notes,
 The emission spectrum of an element is the unique line spectrum produced by an element when it is
burned in a flame. Shown here is the emission spectrum of hydrogen, and a gas discharge tube filled
with hydrogen, with an electric current running through it, producing the characteristic lavender color
of hydrogen. (We’ll look at how these spectra form next…)
♦
Here is a good animation of several flame tests for different metals:
http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/039_FlameTestsMet.MOV
QUANTUM THEORY – MAX PLANCK
♦
Here is a quick synopsis of Quantum Theory, first suggested by Max Planck in 1900:
Quantum Theory describes the particles that make up matter and how they interact
with each other and with energy. The name “quantum theory” comes from the fact
that it describes matter and energy in the universe in terms of single indivisible units
called quanta.
Max Planck
9
Name: ______________________
Chemistry 1 Notes,
Quantum theory is different from classical physics. Classical physics is an approximation of the set
of rules and equations in quantum theory, and can help explain the motion of a car accelerating or of a
ball flying through the air.
Quantum theory, on the other hand, can accurately describe the behavior of the universe on a much
smaller scale, that of atoms and smaller particles. The rules of classical physics do not explain the
behavior of matter and energy on this small scale. Quantum theory is more general than classical
physics, and in principle, it could be used to predict the behavior of any physical, chemical, or
biological system. However, explaining the behavior of the everyday world with quantum theory is
too complicated to be practical.
Quantum theory not only specifies new rules for describing the universe but also introduces new ways
of thinking about matter and energy. The tiny particles that quantum theory describes do not have
defined locations, speeds, and paths like objects described by classical physics. Instead, quantum
theory describes positions and other properties of particles in terms of the chances that the property
will have a certain value. For example, it allows scientists to calculate how likely it is that a particle
will be in a certain position at a certain time (see the Heisenberg Principle below).
Quantum theory describes all of the fundamental forces—except gravity—that physicists have found
in nature. The forces that quantum theory describes are the electrical, the magnetic, the weak, and the
strong.
One of the striking differences between quantum theory and classical physics is that quantum theory
describes energy and matter both as waves and as particles. For example, classical physics considers
light to be only a wave, and it treats matter strictly as particles. Quantum theory acknowledges that
both light and matter can behave like waves and like particles.
 Max Planck said that when an object (like an electron) loses energy, it does not do so continuously
(which it would if radiation were in the form of waves).
 Energy is radiated off in small specific amounts called quanta.
 A quantum is a finite quantity of energy that can be gained or lost by an atom.
 A photon is a quantum of light – a particle of radiation.
 Radiation is absorbed and emitted only in whole numbers of photons.
 According to Planck, there is a relationship between the frequency of a particular radiation and the
energy with which it is associated.
 E = h , where “h” is known as Planck’s constant (h = 6.626 x 10–34 Js)
 As frequency increases, the energy of the radiation increases (frequency and energy are
directly proportional).
 Using the equation: c =  (speed of light = wavelength x frequency), you can rearrange the first
equation to relate energy with wavelength:
hc
E=
λ
 As wavelength increases, the energy of the radiation increases (wavelength and energy are
inversely proportional).
10
Name: ______________________
Chemistry 1 Notes,
 Consider the emission spectrum of hydrogen: it consists of only 4 lines (red, blue-green, blue, and
violet).
 Remembering the Bohr model of the H atom, now we can explain why these are the only colors
produced in an emission spectrum.

The state of an atom in which its electrons exist at their normal energy levels is called the
ground state, or E0. When electrons absorb energy (when they are heated, or two atoms
collide, etc.), they get “excited” or “promoted” up to higher energy levels (excited states or
E’). An electron must gain an amount of energy equal to the energy difference between
the two energy levels, in order to move from its ground state to the higher energy level.

When these excited electrons drop down to lower energy levels, they give off the excess
energy they no longer need as photons of electromagnetic energy. Specifically, they must
release an amount of energy equal to the energy difference between these two energy
levels as a photon of electromagnetic radiation.
11
Name: ______________________
Chemistry 1 Notes,


Look at the Bohr model for the hydrogen atom shown below, and make sure you understand
the specific details of the process. And note that we only use the Bohr model to describe the
hydrogen atom. Other atoms, with more electrons, become too complex to describe.
This basic idea explains why each element (with its own unique arrangement of electrons and
energy levels) would have its own unique emission spectrum.
12
Name: ______________________
Chemistry 1 Notes,
EVOLUTION OF THE ATOMIC MODEL
♦
Year
Scientist
Model
1803-07
Dalton
Billiard Ball
Model
1898
Thomson
Plum Pudding
Model
1910
Rutherford
1913
Bohr
Planetary
Model
1925
Schrodinger
Wave
Mechanical
Model
In 1911, Rutherford suggested that
electrons orbit the atomic nucleus like
planets round the Sun (to explain his
gold foil experiment results).
In 1914, Niels Bohr modified
Rutherford's model by introducing the
idea of energy levels:
1. Electrons move around nucleus in
orbits of definite potential energy.
2. The lowest energy orbit is the one
closest to the nucleus; the orbit
with the most energy is farthest
away from the nucleus.
3. An electron can neither gain nor
lose energy in its orbit, but could
move up or down into other orbits.
 Schrodinger’s quantum model of the
atom (wave mechanical model) –
electrons move around the nucleus in
orbitals (clouds) of increasing potential
energy, where they are most likely to
be found (wave-like behavior).
 Heisenberg Uncertainty Principle –
You cannot simultaneously measure
the position and velocity of an electron
Werner Heisenberg
♦
Good website that further summarizes development of atomic theory, including animations:
http://www.broadeducation.com/htmlDemos/AbsorbChem/HistoryAtom/page.htm
13
Name: ______________________
Chemistry 1 Notes,
♦
We mentioned earlier that both light and electrons exhibit particle-like and wave-like behavior. The
particle part is easy. An electron has mass and takes up space – not much of either, but it’s there.
The wave part is a bit more convoluted.
 Suffice it to say that a scientist named Louis de Broglie first
postulated the idea. He reasoned that if light could have waveparticle duality, then so should matter. He applied this idea to
electrons specifically, and his prediction received experimental
confirmation from two scientists experimenting with electron
diffraction in crystals. Thus, de Broglie’s theory has become
cemented into the foundations of modern physics.
 As illustrated in this diagram, you can even apply the concept
of wave-like behavior to an electron as it orbits around the
nucleus of an atom. This concept is important in more
advanced theories of molecular bonding.
14
Name: ______________________
Chemistry 1 Notes,
IV. Quantum numbers
 Quantum numbers – specify the properties of atomic orbitals and their electrons
 There are four quantum numbers:
1) Principal
2) Orbital
3) Magnetic
4) Spin
♦
Principal quantum number (n): indicates the energy levels (shells) surrounding the nucleus
 Values of n are whole #s only: 1-7
 As n increases, the distance of e- from the nucleus increases & so does the energy of the electrons
in that energy level.
 Orbital quantum number (l): indicates orbital shape (subshells or sublevels)
 orbital – 3-dimensional region around the nucleus where a particular e- can be located
s = sphere
p = peanut
d = double peanut
f = flower
an s-orbital

s, p, d, and f actually stand for sharp, principal, diffuse and fundamental, but we use sphere,
peanut, double peanut and flower to help remember and describe their shape.
15
Name: ______________________
Chemistry 1 Notes,
 Electrons with a given orbital quantum # (s, p, d, and f) can always be found orbiting the nucleus
somewhere in that shape/region.
 These orbitals are really the shapes of electron clouds, then. They describe paths electrons of
a given energy sweep out as they move around the nucleus.
 With this in mind, the actual motion of the electron cannot be defined. Remember the
Heisenberg uncertainty principle? We can know either the electron’s position or its velocity,
but not both simultaneously.
 How an electron exists in its orbital is essentially unknown. In fact, it only occupies this
“shape” about 90% of the time. The other 10%, the electron is elsewhere, somewhere in the
universe (typically limited by relativity) but not in the orbital. So these shapes can actually
be described as probability models, and electrons of very specific energy are found within
these orbitals about 90% of the time.
 Different orbitals have slightly different energies, so different electrons cannot move between
orbitals (unless the energy levels are degenerate, like the px, py, and pz orbitals, for a given
energy level) without absorbing or giving off energy.
 Magnetic quantum number (ml): indicates the orientation of an orbital around the nucleus
 s – only one orientation (it’s round)
 p – three axes (x, y, z)  three possible orientations – px, py, pz
 p-orbital takes the letter of the axis it’s centered on (see previous page)
 d – five possible orientations (see previous pages)
 f – seven possible orientations (see previous pages)
16
Name: ______________________
Chemistry 1 Notes,
 Spin quantum number (ms): indicates the two possible states of an electron in an orbital
 The two values of the spin quantum number are +½ and –½.
 Each orbital can hold no more than two electrons, which must have opposite spins.
 These electrons can be symbolized with an up arrow () and a down arrow ().
17
Name: ______________________
Chemistry 1 Notes,
V.
Electron Configurations
 Electron configuration – the arrangement of electrons in atoms
 Aufbau principle – an electron occupies the lowest-energy orbital that can receive it
 Electrons, like everything else in nature, want to be stable – arranged with the lowest possible
energy.
? Which orbitals have the lowest energy? – the ones closest or farthest away from the nucleus?
(closest – 1s, then 2s, then 2p, etc.)
 Hund’s rule – orbitals of equal energy are each occupied by one electron before any one
orbital receives a second electron, and all electrons in singly occupied orbitals
must have the same spin
 Ex: each of the three p orbitals gets one electron before any one gets a second electron.
 Each of these unpaired p electrons will have the same spin, either +½ or -½.
?
How many unpaired electrons can there be (max) in a d sublevel before any of the d orbitals gets
a second electron?
 Pauli exclusion principle – no two electrons in the same atom can have the same set of four
quantum numbers
 They are all unique – we name each of them differently, as they cannot have the same set of
quantum numbers.
 Electrons in the same orbital have different spin quantum #s.
♦
This website has a great animation explaining these rules / principles and some of the notations on the
next page of notes:
http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/043_ElectronConfig.MOV
18
Name: ______________________
Chemistry 1 Notes,
 There are several notations that can be used to indicate electron arrangement in an atom:
1) orbital notation
 An unoccupied orbital is represented by a line ___ and up or down arrows (since a pair of
electrons in an orbital will have opposite spins).
 Underneath the lines will be the principle quantum # (shell) and orbital (subshell).
2) electron configuration notation
 Write the principle quantum # and subshell and add a superscript indicating how many
electrons are present in that orbital.
 Shorthand electron configuration notation uses noble gases enclosed in brackets to
indicate what which orbitals have already been filled.
 The electron arrangement Mg could also be represented as: 2.8.2 or 2,8,2
to indicate the number of electrons occupying each successive energy level. Know this
arrangement, because it will definitely be on the IB exam.
3) electron-dot notation
 Shows only the electrons occupying the highest (or outermost) main energy level.
bromine =




Only s and p orbitals are in the highest energy level!!
Therefore, there can only be a max of 8 electrons in the highest energy level.
An atom with a full 8 electrons has a full octet.
All atoms would like to have a full octet, because that is the most stable arrangement.
19