Download Structure of Matter

Structure of Matter
Important Discoveries About the
Atom

Law of Conservation of Matter
◦ In 1774, Antoine Lavoisier performed experiments &
measurements that led to this law.
◦ In chemical reactions, matter cannot be created nor
destroyed

Law of Constant Composition
◦ In 1700, Joseph Proust mad measurements on
chemical reactions and compounds to develop this
law
◦ Each pure chemical compound always has the same
percentage composition of each element by mass.

These laws led John Dalton to develop his theory.
Important Discoveries About the
Atom

Dalton’s Atomic Theory states that:
◦ All matter is composed to tiny, indivisible
particles, called atoms, that cannot be created or
destroyed.
◦ Each element has atoms that are identical to each
other in all of their properties, and these
properties are different from the properties of all
other atoms.
◦ Chemical reactions are simple rearrangements of
atoms from one combination to another in small
whole-number ratios.
Important Discoveries About the
Atom

Dalton’s atomic theory led him to
propose the …
◦ Law of Multiple Proportions
 When 2 elements can be combined to made 2
different compounds, and if samples of these 2
compounds are taken so that the masses of one of
the elements in the 2 compounds are the same in
both samples, then the ratio of the masses of the
other element in these compounds will be a ratio
of small whole numbers.
Law of Multiple Proportions
Important Discoveries About the
Atom
In 1834, Michael Faraday showed that an
electric current could cause chemical
reactions to occur, demonstrating the
electric nature of the elements.
 In the 1870s the cathode ray tube was
developed by Sir William Crookes.

◦ He mistakenly thought that the cathode rays
were negatively charge molecules instead of
electrons.
Important Discoveries About the
Atom
In 1897, J.J. Thomson determined that
cathode rays were a fundamental part of
matter he called electrons.
 Thomson also discovered the electron’s
charge to mass ratio, (e/m = -1.76 × 108
coulombs per gram) by measuring the
deflection of the cathode rays in the
presence of electric and magnetic fields.

J.J. Thomson
Important Discoveries About the
Atom

Oil Drop Experiment
◦ In 1909, Robert Millikan performed this
experiment. He determine the charge of an
electron to be -1.60 × 10-19 coulomb.
◦ Using Thomson’s charge to mass ratio, he
then calculated the mass of an electron= 9.11
× 10-28g.
Important Discoveries About the
Atom

From those experiments, the plum
pudding model was developed, which has
electrons swimming in a sea of positive
charges.
Important Discoveries About the
Atom

Gold Foil Experiment
◦ Performed by Ernest Rutherford in 1910.
 He was interested in radioactive materials, alpha &
beta particles had already been discovered.
Rutherford’s Nuclear Model
Important Discoveries About the
Atom

In 1919, Rutherford also discovered the
proton with a mass of 1.67×10-24g.
◦ It’s 1836 times bigger than an electron

His student, James Chadwick, discovered
the nucleus in 1932. It has almost the
same mass as a proton.
Subatomic Particles
Name
Symbol
Electron
e or e-
Absolute
Absolute
charge
mass (g)
(coulombs)
-1.602×10-19
9.109×10-
Relative
charge
Relative
mass
-1
5.486×10-4
+1
1.0073
0
1.0087
28
Proton
p or p+
+1.602×10-19 1.673×1024
Neutron
n or n0
0
1.675×1024
Important Discoveries About the
Atom

While all of this stuff was happening,
other physicists were interested in the
interaction between light and matter.
(mid-1800s)
◦ One thing they discovered was that each
element, when heated or sparked with
electricity, gives off characteristic colors.
 A spectroscope was used to show these colors
consist of discrete wavelengths of light (line
spectra), not the uniform rainbow.
 The line spectra for most elements & compounds can be
complex but hydrogen’s is relatively simple.
Important Discoveries About the
Atom
Important Discoveries About the
Atom

In 1885, Johann Balmer found a
mathematical relationship between the
wavelengths of the lines in the visible light
region of the spectrum.
◦ Similar lines were discovered in the infrared
(Paschen series) and ultraviolet (Lyman series)
regions.
◦ Johannes Rydberg extended Balmer’s
equations so that all of the wavelengths could
be predicted.
Important Discoveries About the
Atom

Planetary or Solar System Model
◦ Proposed by Niels Bohr in 1913. He assumed
that electrons move around the nucleus in
only certain circular orbits.
◦ Max Planck described light as packets or
quanta of energy called photons.
Important Discoveries About the
Atom
In 1924, Louis de Broglie suggested that if light can be
considered as particles, then the small particles
(electrons) may also have the characteristics of
waves!
 In 1927, Erwin Schrödinger applied mathematical
equations for waves to electrons in the atom and
began the wave-mechanical model of the atom.

◦ For the hydrogen atom, the results are very similar to
Bohr’s model of the atom except the electron doesn’t
follow a precise orbit. The position of the electron is
described as a probable location.
◦ Werner Heisenberg developed the uncertainty principle in
the 1920s that says the position & momentum of any
particle cannot be both known exactly at the same time.
As you know one more precisely, the other becomes less
certain.
Atomic Structure

An atom usually exists in the lowest possible
energy state – called the ground state.

An atom that has more energy than the
ground state is said to be in an excited state.

When at atom loses energy in going from an
excited state to the ground state, that
energy is emitted as light.
Atomic Structure
Atomic Structure

Wavelength, Frequency, & Energy of Light
◦ All EM radiation may be considered as waves
defined by wavelength (λ) & frequency (ν).
 Wavelength – distance between 2 repeating points
on a sine wave.
 Frequency – the # of waves that pass a point in
space each second
Atomic Structure

Wavelength and frequency are inversely
proportional to each other.
(wavelength)(frequency) = speed of light
 λν = c

c = 3.00×108m/s
 λ= meters
 ν = s-1 or 1/s

Atomic Structure
Max Planck discovered that the energy of
the EM waves is proportional to the
frequency & inversely proportional to the
wavelength
 hν = E
Planck’s constant (h)
h =E
h = 6.63×10-34Js
λ

Atomic Structure
The Bohr Model- requires the electrons
in the atom be confined to specific,
allowed orbits. Using physics, the energy
of an orbit with the number, n, is
 En = -2πme4 = -2.178×10-18joule
n2h2
n2



Where m = mass e-, e = charge on e-, h = Planck’s constant,
& n = principal quantum number
n also represents the number of each orbit in Bohr’s
model, starting with the one closet to the nucleus
Atomic Structure




Bohr’s orbits
Energy, in the form of light, is emitted from
an atom when an electron moves from an
orbit to a lower-numbered orbit.
When an electron is promoted to a higher
numbered orbit, energy must be added.
The energy difference between 2 orbits is
constant so you know how much energy is
being released when the electron drops
down to its original orbit. You use the
equation in the last slide to calculate the
energy difference.
Atomic Structure

Emitting light

Another representation
Hydrogen
Atomic Structure
Bohr also thought that the momentum (mass ×
velocity) of the electron be related to the size of
the electron’s orbit.
 Mv = nh
2πr
 For hydrogen (n = 1), the radius of the electron
was calculated to be 53pm, called the Bohr radius.
The radii of other orbits are all whole-number
multiples of it.

◦ Gave chemists a theoretical value for the size of a
hydrogen atom.

Bohr won a Noble Prize for all this. Go Bohr!
Atomic Structure

Wave-Mechanical Model of the Atom
◦ After Bohr’s achievement, Louis deBroglie
ascertained that electrons could also act as
waves, not just particles.
 1 way of looking at this duality, is to look at the equation
for the energy of a wave…
 E = hv = h = mc2
λ
 Rearranging this equation, you can see the relationship
between the mass of the electron and the wave
(frequency) … = h = mc2
λ

Atomic Structure

Describing the motion of an electron, however,
requires complex “wave equations” and higher
level calculus (differential equations) than is taught
in high school. But understanding the results of
these wave equations can be done:
◦ Wave equations require 3 numbers, called quantum
numbers, in order to reach a solution.
The principal quantum number, n
The azimuthal quantum number, l
The magnetic quantum number, ml
And a 4th, unique, quantum number, the spin quantum number,
ms
 There are specific rules for applying these numbers to
electrons but it isn’t required on the AP test.




Atomic Structure
◦ The wave equations changed the picture of
the atom completely. The fixed orbits of the
Bohr model are replaced with a cloud of
electrons around the nucleus. The modern
orbital is the region of space in which there is
the greatest (90%) probability of finding an
electron.
Bohr orbit as fixed rings
Probability plot of electrons
in the 1st orbit
Probability plot with
90% of electrons within
circle
Atomic Structure
◦ The circular orbits of the Bohr theory are
replaced with spherical electron clouds. The wave
equations have shown that most electron clouds
have shapes that are more complex than Bohr’s
orbits but are still simple geometric shapes.
◦ The arrangement of electrons deduced from the
wave equations agrees well with the periodic
table. Many physical & chemical properties of
elements & compounds are more fully
understood with knowledge gained about the
electronic structure & orbital shape.
Atomic Structure
◦ The results of the wave equations agree
completely with the Bohr model. Specifically,
the energy change for an electron moving
from one electron cloud to another. Also the
53pm radius found for the electron in the
hydrogen atom still holds true.
◦ The Heisenberg uncertainty principle is
fundamental to the wave mechanical model.
The exact location and momentum of an
electron cannot be known at the same time.
Atomic Structure

Principal energy levels (shells)
◦ Another term for principal quantum number
◦ There are a maximum of 7 shells in an atom.
◦ 1st shell is closest to the nucleus, the 7th is
farthest away.
◦ As they become larger, the further away from
the nucleus, the more electrons they can hold.
Atomic Structure

Sublevels (subshells)
◦ Each principal energy level contains one or more
sublevels, or the azimuthal quantum number, l.
 l can never be greater than n-1.
◦ How many sublevels can each principal energy
level hold?
 The same number as the value of n for that energy level.
Ex. n = 3? It has 3 sublevels. n = 6? It has 6 sublevels.
However, for the 118 known elements, only 4 of the
sublevels are used.
Sublevels
Principal level, n
Sublevel number, l
Sublevel letter
1
0
s
2
0, 1
s, p
3
0, 1, 2
s, p, d
4
0, 1, 2, 3
s, p, d, f
5*
0, 1, 2, 3
s, p, d, f
6*
0, 1 ,2
s, p ,d
7*
0, 1
s, p
*Only sublevels used by known elements are shown here.
Atomic Structure

Orbitals
◦ Each sublevel may contain one or more
electron orbitals, a region of space that has a
high electron density.
 Each orbital may hold a maximum of 2 electrons.
 In order to share the orbital, the electrons must have
opposite spins.
 The number of orbitals a sublevel can have depends on its
azimuthal quantum number, l, and is equal to 2l +1.
Orbitals
Sublevel
number, l
Sublevel letter
Number of
orbitals 2l + 1
Number of
electrons per
Sublevel
0
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
Each orbital is given a magnetic quantum number, ml.Values range from –l
to +l.
orbital
ml. values
s
0
p
-1, 0, +1
d
-2, -1, 0, +1, +2
f
-3, -2, -1, 0, +1, +2, +3
Orbital shapes
Electronic Structure of the Atom
s
d
p
f
Orbital filling
diagram
7s
7p
6s
6p
6d
5s
5p
5d
5f
4s
4p
4d
4f
3s
3p
3d
2s
2p
1s
start
Electron Configurations

Some rules to abide by:
◦ Aufbau Principle – electrons fill the lowest
available energy level before moving to a
higher one.
◦ Some exceptions:
Element
Electron Configuration
Copper, Cu
1s22s22p63s23p64s13d10
Silver, Ag
1s22s22p63s23p64s23d104p65s14d10
Gold, Au
1s22s22p63s23p64s23d104p65s24d105p66s14f145d10
Chromium, Cr
1s22s22p63s23p64s13d5
Molybdenum.
Mo
1s22s22p63s23p64s23d104p65s14d5
Abbreviated Electron
Configurations- Use noble gases
1
2
3
4
5
6
7
s
d
p
f
Example
Fe = 1s22s22p63s23p64s23d6
Fe = [Ar]4s23d6
Valence Electrons

Many times, chemists are only interested
in the outermost electrons in an atom,
the valence electrons.
◦ Only s and p electrons are valence electrons.
Orbital Diagrams

Hund’s Rule
◦ p, d, or f orbitals in a sublevel must all be filled
with one electron each before a 2nd electron
is allowed to pair in any orbital.
Quantum Numbers
Principal quantum number, n = 1-7
Azimuthal quantum number, l = n-1 = 0-6
Magnetic quantum number ml = -l -0-+l
 Spin quantum number, ms = + ½ or – ½



Lowest possible values for each quantum number are
used 1st.
 Pauli exclusion principle- no 2 electrons can have the
exact same quantum number
 While writing the quantum numbers is not on the AP
test, knowing what they look like according to the
rules is.
 See handout.

Quantum numbers, what do they
mean?




Principal quantum number, n – represents
the average distance of the electron from
the nucleus, or the size of the principal
energy level.
Azimuthal quantum number, l – represents
the shape(s) of the orbitals within the
sublevel.
Magnetic quantum number ml – represents
the oreientation of each orbital in space on
an x, y, z axis.
Spin quantum number, ms – represents the
“spin” of the electrons.