Unit 2

Atomic Theory

John Dalton









Ernest Rutherford
Robert Millikan
J.J. Thompson




Atomic Structure






Law of Conservation of Mass
Law of Definite Proportions
Law of Multiple Proportions
Protons, neutrons, electrons
Atomic number
Isotopes
Mass number
Average atomic mass
Wave nature of light


Electromagnetic Spectrum
C = λv
Bohr Models
Photoelectric effect
Heisenberg Uncertainty
Principle
Quantum numbers




Absorption/emission
E = hc/ λ
Pauli Exclusion Principle
Hund’s Rule
Aufbau Principle
Configurations (orbital,
electron, noble gas)


Paramagnetism/diamagnetism
Exceptions
B.C.
400 B.C. Democritus and Leucippos use the term "atomos”
2000 years of Alchemy
1500's
 Georg Bauer: systematic metallurgy
 Paracelsus: medicinal application of minerals
1600's
Robert Boyle:The Skeptical Chemist. Quantitative experimentation,
identification of elements
1700s'
 Georg Stahl: Phlogiston Theory
 Joseph Priestly: Discovery of oxygen
 Antoine Lavoisier: The role of oxygen in combustion, law of conservation of
mass, first modern chemistry textbook
1800's
Joseph Proust: The law of definite proportion (composition)
 John Dalton: The Atomic Theory, The law of multiple proportions
Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules
Amadeo Avogadro: Molar volumes of gases
Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements
 Dmitri Mendeleyev: The periodic table
 J.J. Thomson: discovery of the electron
 Henri Becquerel: Discovery of radioactivity
1900's
 Robert Millikan: Charge and mass of the electron
 Ernest Rutherford: Existence of the nucleus, and its relative size
 Meitner & Fermi: Sustained nuclear fission
 Ernest Lawrence: The cyclotron and trans-uranium elements
 400
BC
 Democritus




Matter consists of small particles
Called them “atomos”
Idea rejected by peers
No scientific proof
 Aristotle



All matter continuous
4 elements = earth, water, air, and fire
No scientific proof
 Idea
endured for
2000
years
School Teacher
 Atomic Theory

1.
2.
3.
4.
5.
All matter is composed of extremely small particles called
atoms. There are different kinds called elements.
Atoms of the same element are identical in size, mass, and
other properties; atoms of different elements differ in size,
mass, and other properties.
Atoms cannot be subdivided, created, or destroyed.
Atoms of different elements combine in simple, whole
number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or
rearranged but never destroyed/created.

Law of Conservation of Mass
 Total mass present before chemical reaction is same as
mass after chemical reaction
 2H2O  2H2 + O2 If you have 10 grams of water to start, you
will get 1.12 g of hydrogen and 8.88 g of oxygen

Law of Constant Composition (definite proportions)
 Relative numbers and kinds of atoms are constant
 Water is 88.8% oxygen and 11.2% hydrogen by mass no
matter how much you have

Law of Multiple Proportions
 If two elements combine to form more than one compound,
the masses of the two elements are in the ratio of small
whole numbers
 CO2 versus CO (mass ratio is 2 to 1 for oxygen)
 British
Physicist
 Discovered electron
 Cathode-ray experiment
 Plum pudding view of atom
 Electric
current sent through gases in glass
tube called cathode-ray tube
 Surface
of tube opposite the cathode glowed –
caused by stream of particles
 Ray

traveled from cathode to anode
Cathode rays deflected by
magnetic field away from
negatively charged object
(like a magnet)

Cathode rays concluded to
have negative charge
 American
Physicist
 Charge on each electron is same
 Charge of electron is -1.6022 x 10-19C
 Calculated mass of electron as
9.10x 10-31 kg
 Oil drop experiment

Drops of oil that had picked
up extra electrons allowed
to fall between two
electrically charged plates

Measured how voltage on
plates affected rate of fall

Calculated charges of drops
then deduced charge of a
single electron on the drops
 Discovered
nucleus
 Planetary model of the atom

Bombarded thin piece gold foil with
alpha particles (positively charged
particle 4 times mass of hydrogen
atom)


1 in 8000 particles deflected back
toward source


Expected to pass right through gold foil
“As if you fired 15-inch artillery shell at
a piece of tissue paper and it came back
and hit you”
Concluded most of atom is empty
space except for a very small force
within atom

Called positive bundle of matter the
“nucleus”
 Atom
consists of proton, neutron, and
electron



Proton charge = +1
Neutron charge = 0 (neutral)
Electron charge = -1
 Protons

and Neutrons located in nucleus
99.9% of atom’s mass is in nucleus
 Electrons
located outside the nucleus
47
Silver
Ag
107.87
Atomic number
Name of the element
Element Symbol
Atomic mass
 Atomic

Number
equal to number of protons in an atom
 Element


Symbol
First letter always capitalized
If second letter exists, it is lowercase
 Isotopes
are atoms of the same element
having different masses due to varying
numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
 Atomic
mass is the average of all the
naturally isotopes of that element.
Isotope
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Carbon = 12.011
 Mass
Number = Protons + Neutrons
 Not found on periodic table
 Isotopes have different mass numbers (due to
neutrons)
C– 12
Mass number
Atomic number
Mass number
 JJ
Thomson won the Nobel prize for
describing the electron as a particle
 His son, George Thomson won the Nobel prize
for describing the wave-like nature of the
electron.
The
electron is
a particle!
The electron
is an energy
wave!
Much of what has been learned about atomic
structure has come from observing the
interaction of visible light and matter.
 1924
De Broglie suggested that electrons have
wave properties to account for why their energy
was quantized.
 He
reasoned that the electron in the hydrogen
atom was fixed in the space around the nucleus.
 He
felt that the electron would best be
represented as a standing wave.
 As
a standing wave, each electron’s path must
equal a whole number times the wavelength.
The electron propagates
through space as an energy
wave. To understand the
atom, one must understand
the behavior of
electromagnetic waves.
Louis deBroglie
 Wavelength,

l
The distance for a wave to go through a
complete cycle.
 Amplitude

Half of the vertical distance from the top to the
bottom of a wave.
 Frequency,

n
The number of cycles that pass a point each
second.

Longer wavelength = lower frequency = lower energy

Shorter wavelength = higher frequency = higher energy
 The
SI unit of frequency (n) is the hertz, Hz
1 Hz = 1 s-1
 Wavelength
and frequency are related
c = ln
c is the speed of light, 2.998 x108 m/s
The wavelength of an argon laser's output is
488.0 nm. Calculate the frequency of this
wavelength of electromagnetic radiation.
c = ln

Convert nm to m
488 nm x (1 m / 109 nm) = 4.88 x 10-7 m

Then, substitute into c = λν
(4.88 x 10-7 m) (v) = 3.00 x 108 m s-1
v = 6.15 x 1014 s-1 = 6.15 x 1014 Hz
 Electromagnetic

Energy in the form of transverse magnetic and
electric waves.
 Electromagnetic

Spectrum
Contains all forms of electromagnetic radiation
 Visible

Radiation
spectrum
Portion of electromagnetic spectrum that we can see
(colors)
 ‘White’
light is actually a blend of all visible
wavelengths. They can separated using a
prism.
 Neils
Bohr studied the spectra produced when
atoms were excited in a gas discharge tube.
 Each
element produces its own set of
characteristic lines

Bohr proposed a model of how electrons moved
around the nucleus.

He wanted to explain why electrons did not fall in
to the nucleus.

He also wanted to account for spectral lines being
observed.

He proposed that the energy of the electron was
quantized - only occurred as specific energy levels.
the Bohr model,
electrons can only
exist at specific
energy levels
(orbit).
Energy
 In
 Each
energy level
was assigned a
principal quantum
number, n.

The Bohr model is a
‘planetary’ type
model.

Each principal
quantum represents a
new ‘orbit’ or layer.

The nucleus is at the
center of the model.

Absorption – Electromagnetic radiation is absorbed by an
atom causing electrons to jump to a higher energy state
(excited state).

Emission – Energy is released by an atom as particle of
light (photon) as electrons fall back to the lower energy
state (ground state).

Depending on
frequency of photon,
different colored light
may be seen
 Although
electromagnetic radiation has
definite wave properties, it also exhibits
particle properties.
 Photoelectric
•
•
•
•
effect.
First observed by Hertz and then later explained
by Einstein.
When light falls on Group IA metals, electrons
are emitted (photoelectrons).
As the light gets brighter, more electrons are
emitted.
The energy of the emitted electrons depends on
the frequency of the light.
 The
energy of a photon is proportional to the
frequency.
(Photon energy)
E= hn
 The
energy is inversely proportional to the
wavelength (remember c = λν so v = c/λ ).
E = hc /l
h is Plank’s constant, 6.626 x 10-34 J . S
c is the speed of light, 2.998 x108 m/s
 Determine
the energy, in kJ/mol of a photon of
blue-green light with a wavelength of 486 nm.
E=
hc
l
-34 J.s)(2.998 x 108 m.s-1)
(6.626
x
10
=
(4.86 x 10-7 m)
= 4.09 x 10-19 J
h
l =
mv




l
h
m
v
=
=
=
=
wavelength, meters
Plank’s constant
mass, kg
frequency, m/s
 Using
De Broglie’s equation, we can calculate the
wavelength of an electron.
l =
l=
h
mv
6.6 x 10-34 kg m2 s-1
(9.1 x 10-31 kg)(2.2 x 106 m s-1)
= 3.3 x 10-10 m
The speed of an electron had already been reported
by Bohr as 2.2 x 106 m s-1.
In order to observe an electron, one would need
to hit it with photons having a very short
wavelength.
 Short wavelength photons would have a high
frequency and a great deal of energy.
 If one were to hit an electron, it would cause
the motion and the speed of the electron to
change.
 According to Heisenberg, it is impossible to
know both the position and the speed of an
object precisely.

 Schrödinger
developed an equation to
describe the behavior and energies of
electrons in atoms.
 His
equation is similar to one used to
describe electromagnetic waves.
 Each
electron can be described in terms of
its quantum numbers.

Each electron in an atom has a unique set of 4
quantum numbers which describe it.

Principal quantum number

Angular momentum quantum number

Magnetic quantum number

Spin quantum number
 Principal

quantum number, n
Tells the size of an orbital and largely
determines its energy.
n = 1, 2, 3, ……
 Angular

momentum, l
The number of subshells that a principal level
contains. It tells the shape of the orbitals.
l = n – 1 to 0
 Orbitals


An orbital is a region within an energy level where there is a
probability of finding an electron
Orbital shapes are defined as the surface that contains
90% of the total electron probability.
 Magnetic

quantum number, ml
Describes the direction that the orbital projects
in space.
ml = -l to +l
(all integers, including zero)

For example, if l = 2, then ml would have values
of -2, -1, 0, 1 and 2.

Knowing all three numbers provide us with a
picture of all of the orbitals.
 Pauli
added one additional quantum number
that would allow only two electrons to be in
an orbital.
 Spin

It can have values of +1/2 and -1/2
 Pauli

quantum number, ms.
exclusion principle
Pauli proposed that no two electrons in an atom
can have the same set of four quantum numbers
 Aufbau

Principle
Electrons are placed into orbitals, subshells, and
shells in order of increasing energy
 Hund’s

Rule
The most stable arrangement of electrons in a
subshell is the one in which electrons have the
most number of parallel spins possible.
 Graphical
representation of an electron
configuration
 One arrow represents one electron
 Shows spin and which orbital within a
sublevel
 Follow all rules(Aufbau principle, two
electrons in each orbital, etc. etc.)
 Use
atomic number as number of electrons in
an atom
He
Be
Ne
Mg
Si
 Diamagnetism



Elements have all of their electrons spin paired
All of an element’s subshells are completed
Not affected by magnetic fields
 Paramagnetism



Not all electrons are spin paired in an element
Most elements are this
Affected by magnetic fields
A
list of all the electrons in an atom (or
ion)
Must go in order (Aufbau principle)
 2 electrons per orbital, maximum

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.
 We
need electron configurations so that
we can determine the number of
electrons in the outermost energy level.

These are called valence electrons.
4
2p
Number of electrons in
the sublevel
Energy Level
Sublevel
He, 2: 1s2
Ne, 10: 1s2 2s2 2p6
Ar, 18: 1s2 2s2 2p6 3s2 3p6
Kr, 36: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
 Orbitals
grouped in s, p, d, and f orbitals
(sharp, proximal, diffuse, and fundamental)
s orbitals
d orbitals
p orbitals
f orbitals

d and f orbitals require LARGE amounts of energy

It’s better (lower in energy) to skip a sublevel
that requires a large amount of energy (d and f
orbtials) for one in a higher level but lower
energy





A way of abbreviating long electron configurations
Since we are only concerned about the outermost
electrons, we can skip to places we know are
completely full (noble gases), and then finish the
configuration
Find the closest noble gas to the atom (or ion),
WITHOUT GOING OVER the number of electrons in the
atom (or ion). Write the noble gas in brackets [ ].
Step 2: Find where to resume by finding the next
energy level.
Step 3: Resume the configuration until it’s finished.
Example: [Ne] 3s2 3p5

Remember d and f orbitals require LARGE
amounts of energy

If we can’t fill these sublevels, then the next
best thing is to be HALF full (one electron in
each orbital in the sublevel)

There are many exceptions, but the most
common ones are

For the purposes of this class, we are going to
assume that ALL atoms (or ions) that end in d4 or
d9 are exceptions to the rule. This may or may
not be true, it just depends on the atom.
 d4
is one electron short of being HALF full
 In order to become more stable (require
less energy), one of the closest s
electrons will actually go into the d,
making it d5 instead of d4.
For example: Cr = [Ar] 4s2 3d4
 Since this ends exactly with a d4 it is an
exception to the rule. Thus, Cr = [Ar] 4s1 3d5

 Remember,
half full is good… and when an
s loses 1, it too becomes half full!
 d9 works the same way