Download File

“Atomic Structure”
Atomic Structure


Atomic Structure Video
Milikan Oil Drop Experiment Video
In the Beginning

The Greek philosopher Democritus
(460 B.C.) was among the first to
suggest the existence of atoms (from the Greek
word “atomos”)


He believed that atoms were indivisible and
indestructible
His ideas did agree with later scientific theory, but did
not explain chemical behavior, and was not based on
the scientific method – but just philosophy
Dalton’s Atomic Theory (experiment based)
John Dalton
(1766 – 1844)
1) All elements are composed of
tiny indivisible particles called
atoms
2) Atoms of the same element are
identical. Atoms of any one
element are different from
those of any other element.
3) Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds (Law of Definite Proportions)
4) In chemical reactions, atoms are combined,
separated, or rearranged – but never
changed into atoms of another element.
Dalton’s Atom
Discovering the first
subatomic particle.
In 1897, J.J. Thomson used a
cathode ray tube.
What do you think Thomson
concluded when he saw the
following?
Discovering the first subatomic
particle
Discovery of the Electron
1) Cathode rays cast a well defined shadow:
must travel in straight lines.
2) Rays deflected by magnetic and electrical
fields: So! they must be electrical
charges in motion.
3) Direction of deflection indicated charge was
negative.
Mass of the Electron
Mass of the
electron is
9.11 x 10-28 g
The oil drop apparatus
1916 – Robert Millikan determines the mass
of the electron: 1/1840 the mass of a
hydrogen atom; has one unit of negative
charge
Conclusions from the Study
of the Electron:
a) Cathode rays have identical properties
regardless of the element used to
produce them. All elements must contain
identically charged electrons.
b) Atoms are neutral, so there must be
positive particles in the atom to balance
the negative charge of the electrons
c) Electrons have so little mass that atoms
must contain other particles that account
for most of the mass
Thomson’s Atomic Model
Thomson believed that the electrons
were like plums embedded in a
positively charged “pudding,” thus it
was called the “plum pudding” model.
Ernest Rutherford’s Gold Foil
Experiment
Rutherford’s problem:
In the following pictures, there is a target
hidden by a cloud. To figure out the shape of
the target, we shot some beams into the cloud
and recorded where the beams came out. Can
you figure out the shape of the target?
Target
#1
Target
#2
The Answers:
Target #1
Target #2
Ernest Rutherford’s Gold Foil
Experiment
http://wps.prenhall.com/wps/media/objects/439
/449969/Media_Portfolio/index.html
The Rutherford Atomic Model
Problems with
Rutherford’s Model
There is a major problem with this nuclear model
of the atom. This relates to the fact that the
electron, in orbiting the nucleus, undergoes an
acceleration (recall from grade 10 that an
acceleration arises from a change in velocity,
which can be a change in speed and/or a
change in the direction of motion)
Problems with
Rutherford’s Model
Electrostatic attraction
Centripetal Force
5+
Energy Lost
Problems with
Rutherford’s Model

Thus in this model we would expect the
electron to be continuously emitting
radiation.

Therefore the electron will slow down
(because it is losing energy) and crash
into the nucleus.
Problems with
Rutherford’s Model

At first these were thought to be good
features of the model, as it was known that
atoms do emit radiation (light) when
heated, and also that many atoms,
particularly the heavier ones, do have a
finite lifetime.
Problems with
Rutherford’s Model

First problem: the light emitted by atoms of
a given element occurs only with a certain
set of discrete frequencies. (What would
you expect from the nuclear model?)
We will revisit this problem
Problems with
Rutherford’s Model

Second problem: the time for any atom to
collapse in this model turns out to be of
the order of 0.00000001 seconds. (Why
would that be a problem?)
We will revisit this problem also
Electromagnetic Spectrum

The electromagnetic spectrum is the
ordered series of all known types of
electromagnetic radiation, arranged by
wavelength ranging from the short cosmic
rays through gamma rays, X-rays,
ultraviolet radiation, visible radiation,
infrared radiation, microwaves, to the long
wavelengths of radio energy
Electromagnetic Spectrum
wavelength
crest
trough
Like waves in a piece of string
Electromagnetic Spectrum
Electromagnetic Spectrum
Electromagnetic Spectrum



Wavelength of light is dependent on the
amount of energy the light has.
Colour of light is defined by its wavelength
‫ ג‬ν
c – speed of light
‫ – ג‬wavelength
C=
ν - frequency
Problems with
Rutherford’s Model

First problem: the light emitted by atoms of
a given element occurs only with a certain
set of discrete frequencies. (What would
you expect from the nuclear model?)
Neils Bohr –
Hydrogen Emission Spectrum
Neils Bohr –
Hydrogen Emission Spectrum

Bohr used the colours in the emission spectrum
as evidence for his model of the atom. The
planetary model.

How could Bohr
have connected the
colours of light he
saw to this model of
the atom?
1+
Neils Bohr to the rescue
1913


When an electrical discharge is passed
through hydrogen gas at low pressure it
emits a violet light; examing this light
through a spectroscope splits the light into
series of coloured lines called a line
spectrum.
Different gases produce characteristic line
spectra.
Spectra of Gas Discharges

Hydrogen

Helium

Carbon
Spectra of Gas Discharges

Neon

Iron

Xenon
Neils Bohr to the rescue
1913

Atomic Excitation
1+
electron n=1
1+
electron n=2
Neils Bohr to the rescue
1913

As the electron drops back to a lower level
it releases energy observed as a specific
colour of light.
Neils Bohr to the rescue
1913

Atomic De-excitation
1+
electron n=2
1+
electron n=1
Neils Bohr to the rescue
1913
Theory for orbits ("states") of electrons in atoms:
1.Electrons have certain allowed states in which
they can move without radiating.
2.The allowed states have well-defined
(‘quantized’) energies that can be determined
with normal classical physics.
Bohr’s model (con’d)


Predicts that an electron with a particular
energy travels along a three-dimensional
pathway called an orbit or shell.
These orbits are designated by the
principal quantum number, n, where n is
any positive integer from 1 to infinity.
Bohr’s model
(assumptions made by Bohr)



An electron moves in an energy level without
energy lose. (What is wrong with this
assumption?)
The greater the distance between the nucleus
and the energy level in an atom, the greater the
energy required for an electron to travel in the
energy level.
An electron cannot exist between orbits. It can
move to a higher unfilled orbit if it absorbs a
specific amount of energy and to a lower unfilled
orbit if it loses energy.
Bohr’s model
(assumptions made by Bohr)




The addition of energy to an electron in a given
energy level and the subsequent jump to a
higher unfilled energy level is referred to as a
transition.
The energy required for the transition is equal to
the difference in energy between the energy
levels; it is called a quantum of energy.
A release in the amount of energy by the
electron drops the electron to a lower energy
level.
When an electron is in its lowest energy level
that it can occupy it is said to be in its ground
state.
Neils Bohr –
Hydrogen Emission Spectrum



The emission spectrum of hydrogen is the
simplest emission spectrum because there
is only one electron.
It is not continuous but concentrated into
bright lines. (sunlight has a continuous
spectrum)
This indicates the existence of only certain
allowed energy levels within the atom
Emission spectrum of Hydrogen
Hydrogen Emission Spectrum





As you observe the spectrum is divided into a
number of distinct series, which correspond to
different regions of the electromagnetic
spectrum.
Each series represents a transition in which the
electron falls to a particular energy level.
All transitions to the n = 1 level include the large
n = 1 to n = 2 energy difference and so they are
all high energy transitions found in the UV region
All transitions to the n = 2 level are in the visible
region
All transitions to the n = 3 level are in the
infrared region
Bohr’s brilliant insight is immediately able to
explain the observed frequencies of the
emission spectrum of atomic hydrogen
Hydrogen Emission Spectrum





Each series has a very similar structure of lines
that become closer together going towards
higher frequencies; this is as a result of the
convergence of energy levels.
The higher the energy the smaller the difference
in energy between successive energy levels.
This means the lines in a spectrum will converge
The limit of this convergence indicates the
energy required to completely remove the
electron from the atom.
For example in the case of the Lyman series, to
completely remove an electron (n = infinity) from
a hydrogen atom at n = 1, this energy should be
equal to…
Limits to Bohr’s Model



Bohr made enormous contributions to
atomic theory, especially in developing the
concept of energy levels.
Bohr’s theory was limited in its ability to
predict the line spectra of atoms with more
than one electron.
The development of the quantum
mechanical model of the atom resulted
from Bohr’s early work.
Quantum Mechanical Model

Retains the concept that electrons fill successive
shells, each of which is designated by a principal
quantum number, n.

Quantum mechanics was built on the early idea
that electrons within an atom possess only discrete
quantities of energy, therefore electrons are
arranged in specific energy levels and sub-levels.
Quantum Mechanical Model



Each sub-level is divided into orbitals each of
which hold 2 electrons. The number of orbitals
that each energy level contains is determined by
the formula n2, where n is the principal quantum
number. 2n2 indicates the number of electrons
per energy level.
i.e. For n = 1; n2 = 12 = 1. This means that for n =
1 there is 1 orbital that can hold a maximum of 2n2
electrons, or 2(1)2 = 2 electrons.
Try to calculate the orbitals and electrons in n = 2.
Quantum Mechanical Model
Energy Level
n
1
Total number of
orbitals
n2
1
Electron
capacity
2n2
2
2
4
8
3
9
18
Quantum Mechanical Model

The energy level closest to the nucleus
only contains one sub-level and one
orbital, determined by n2. It is an orbital
with spherical symmetry and is known as
an s orbital. Since it is in n = 1 this orbital
is referred to as a 1s-orbital. It can hold 2
electrons.
Quantum Mechanical Model
http://www.humboldt.edu/~rap1/Chem_resrc/AOSup.htm
Quantum Mechanical Model
(Aufbau Diagram)
E
N
E
R
G
Y
Bohr Model
Protons
1s
Quantum Mechanical Model


The second energy level has two sub-levels,
defined by the principal quantum number n;
where n = 2 and four orbitals defined by n2;
where n2 = 22 = 4. 8 electrons defined by 2n2.
The ‘s’ sub-level has one orbital with spherical
symmetry, called the 2s-orbital and the ‘p’ sublevel has 3 orbitals which have figure 8 electron
distribution. These orbitals are known as 2porbitals
Quantum Mechanical Model
http://www.humboldt.edu/~rap1/Chem_resrc/AOSup.htm
Quantum Mechanical Model

Owing to electron-electron repulsion, the
p-orbitals are at a slightly higher energy
than the s orbitals.
Quantum Mechanical Model
(Aufbau Diagram)
E
N
E
R
G
Y
Bohr Model
2p
2s
Protons
1s
Quantum Mechanical Model

The third energy level has three sub-levels
and n2 orbitals or 32 = 9 orbitals in all; one
3s orbital, three 3p-orbitals and five 3dorbitals for a total of 2n2 electrons or 2(3)2
= 18 electrons.
Quantum Mechanical Model
http://www.humboldt.edu/~rap1/Chem_resrc/AOSup.htm
Quantum Mechanical Model

The 3d-orbitals are at a higher energy than
the 3p-orbitals again due to electronelectron repulsion forces.
Quantum Mechanical Model
(Aufbau Diagram)
E
N
E
R
G
Y
3d
3p
Bohr Model
3s
2p
2s
Protons
1s
Quantum Mechanical Model

The fourth energy level, as well as the sorbital, the three p-orbitals, and five dorbitals, there are seven f-orbitals. At n =
4, n2 orbitals or 42 = 16, there are 16
orbitals and 32 electrons.
Quantum Mechanical Model
7f - orbitals
http://www.humboldt.edu/~rap1/Chem_resrc/AOSup.htm
Quantum Mechanical Model
(Aufbau Diagram)
4f
4d
4p
E
N
E
R
G
Y
4s
3d
3p
Bohr Model
3s
2p
2s
Protons
1s
Quantum Mechanical Model
Energy
Level
n
1
2
3
4
Types of
sublevels
n types
s
s, p
s, p, d
s, p, d, f
Total number of
orbitals
n2
1s=1
1s+3p=4
1s+3p+5d=9
1s+3p+5d+7f=16
Electron
capacity
2n2
2
8
18
32
Quantum Mechanical Model




There are three principles that govern the filling
of orbitals by electrons in the Aufbau diagram.
Aufbau principle: Electrons enter orbitals of
lowest energy first.
Hund’s rule: When electrons occupy orbitals of
equal energy, one electron enters each orbital
until all the orbitals contain one electron with
spins parallel.
Pauli exclusion principle: An atomic orbital
contains a maximum of two electrons which
must have different spins.
Quantum Mechanical Model
(Aufbau Diagram)
Quantum Mechanical Model

The periodic table can be split into its
subshells called “blocks”. There is an s, p,
d, and f block.
i.e. hydrogen, lithium, beryllium, sodium,
magnesium all have there last electron
falling into which subshell?
Quantum Mechanical Model
s
p
d
f
Quantum Mechanical Model
Quantum Mechanical Model

Subshells get filled on the basis of lowest
energy first. An easy way to remember
this is to read it directly from the periodic
table from left to right.
Quantum Mechanical Model
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
7d
6d
5d
4d
3d
7f
6f
5f
4f
Quantum Mechanical Model

Now when we discuss electron configuration we
can use the quantum mechanical model.
Bohr model
Li  2, 1
Mg  2, 8, 2
Cl  2, 8, 7
Quantum Mechanical model
Li
 1s22s1
Mg  1s22s22p63s2
Cl
 1s22s22p63s23p5
Problems with
Rutherford’s and Bohr’s Models

Second problem: the time for any atom to
collapse in this model turns out to be of
the order of 0.00000001 seconds. (Why
would that be a problem?)
Summary: Electron Arrangement


De Broglie first pointed out the wave-particle
duality of nature. His idea was that all particles
exhibit some wave characteristics, and vice
versa.
Heisenberg’s uncertainty principle concerns the
process of observing an electron’s position and
velocity. It is impossible to know accurately both
the position and the momentum of an electron at
the same time.
Summary: Electron Arrangement

Schrodinger developed a mathematical
equation which describes the behaviour of
the electron as a wave. The solutions
(quantum numbers) can be used to calculate
the probability of finding an electron at a
particular point in space.
Summary: Electron Arrangement
Quantum Numbers


The Principle Quantum Number, n
is the number of the energy level and
describes the relative electron cloud
size.
The Secondary Quantum Number, l
(0,1,2,3 or s,p,d,f) describes the shape
of the electron cloud. It gives the
number of sublevels associated with
each energy level
Summary: Electron Arrangement
Quantum Numbers


The Third Quantum Number (Magnetic), ml
is +l …-l, which describes the orientation
in space of each sublevel and gives the
number of orbitals associated with each
sublevel.
The Fourth Quantum Number (Spin), ms,
describes the spin direction of the electron
Ionization Energies and
Quantum Mechanical Model
Quantum Mechanical Model

Is able to explain:
- the irregularities in the ionization energy
chart
- exceeding the octet rule