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ATOMIC THEORIES
Atomic theory is that theory that attempt to describe the compete structure of an atom
...
DALTON’S ATOMIC THEORY
Assumptions of Dalton’s atomic theory
Assumption 1: Matter is made up of very small indivisible particles called atoms
Assumption 2: Atoms can neither be created nor destroyed
Assumption 3: Atoms of the same element are all alike, that is, they have same mass, volume etc
...
Shortcomings of Dalton’s atomic theory
Assumption 1: This assumption is not valid due to existence of three smaller particles in an atom
...
Assumption 2: This assumption is not valid due to existence of both natural and artificial radioactivity
...
Assumptions 3: This assumption is not valid due to existence of isotopes which are atoms of the same
element with the same atomic number but differ in mass number due to difference in number of their
neutrons
...
Assumption 4: To large extent this assumption is valid
...
Assumption5: To large extent this assumption is valid and is supported by the law of chemical
combination
...
For example;
- Molecular formula of chlorophyll is C55 H7205 N4 Mg
- Molecular formula of haemoglobin is C2952 H4664 N812 O832 S8 F4
Also the element silicon occurs in some very complex silicates and thus reduces degree of correctness of
the assumption 5
...
- This beam could be attracted and deflected by a magnet, so it was not light but a stream of charged
particles which he named corpuscles which are nowadays known as electrons and assigned a
negative charge
...
- Thomson interpreted this to mean that electrons were a fundamental part of all matter
...
- He believed that the electrons were embedded in a positive solid matrix, the way that raisins stick in
raisin bun or plums stick in pudding
...
He suggested the atom to be
full of some positive fluid in which electrons were embedded
...
He further thought that the electrons
were arranged in the form of ring shells
...
In one sentence, Thomson atomic model may be explained as “An atom possesses a spherical shape in
which the positive charge is uniformly distributed with electrons embedded over it”
Advantage of Thomson’s atomic model
The model explains satisfactorily how heating a substance starts radiating light
...
According to the Thomson’s model, the atom would look like the inside of a pumpkin
...
The experiment confirms the presence of very small positively charged nucleus in the atom which is
completely against Thomson’s atomic model
...
Understand alternative terms for Thomson’s atomic model!
The Thomson’s model is sometimes known as plum-pudding atomic model (or raisin pudding
model), because it suggests the atom to be full of some positive fluid like a pudding in which electrons
were embedded like plums on that pudding
...
Figure: Thomson's model as watermelon model
By similar reasoning, Thomson’s atomic model is also known as pumpkin atomic model
...
Rutherford’s ∝ −𝐫𝐚𝐲 scattering experiment
In ∝ −𝐫𝐚𝐲 scattering experiment:
-
A narrow beam of ∝ −particles (positively charged particles) is passed through a thin gold foil
...
-
Whenever an ∝ −particle strike the zinc sulphide screen, a tiny flash of light will be produced at that
point
...
A small fraction of the ∝ −particles were either deflected by small angle or reflected (bounced back)
...
The positive charge in the atom is concentrated in very tiny central core called the nucleus (that is why
very few ∝ −particles were able to be either reflected or deflected)
...
That is not correct, it just virtue truth! What is the real
truth?
- The real truth is that, Rutherford discovered the nucleus (Not the proton!)
...
Assumption 2: The electrons are distributed around the nucleus such that there is a lot of empty space in
the atom
...
Assumption 4: Electrons revolve around the nucleus in a closed orbit at very high speed (as the planets do
around the sun)
...
Assumption 6: No energy is emitted when electrons revolve around the nucleus
...
- Any charged particle moving around another of oppositely charged continuously loses (emits) energy
giving off electromagnetic radiation (according to the laws of electrodynamics)
...
- This is simply because as the electrons emit radiation, the energy of electrons should gradually decrease
leading to a constant decrease in the radius of the electrons orbit and hence electrons fall into the
nucleus
...
The spectra provide the evidence of existence of energy levels contradicting with the model which
suggests the atomic spectra to be continuous radiation spectrum as the model does not recognise
presence of energy levels
...
Advantages of Rutherford’s atomic model
Assumptions 1: This assumption is correct
Assumption 2: This assumption is correct
BOHR’S ATOMIC MODEL
Bohr passed an electric current through a glass tube containing hydrogen gas
...
If Rutherford’s random electron model of the atom was correct, Bohr would have seen a continuous
spectrum (a rainbow) that contained all of the colours of light
...
-
Bohr interpreted this to mean that the electrons could only be certain, discrete distances from the
nucleus
...
Assumptions of Bohr’s atomic model
Assumptions 1: An atom consists of very tiny positively charged nucleus
...
Assumption 2: The electrons in an atom revolve around the nucleus in a certain permitted circular orbits
...
Assumption 3: In an atom there are energy levels or energy states called stationary states in which
electrons do not emit (radiate) energy, the energy of the orbits is quantized, that is they have fixed value of
energy to enable electrons keeps on moving in the same orbit
...
Definition of stationary state
Is any of several energy states an atom may occupy without emitting electromagnetic radiation
...
Advantages of Bohr’s atomic model
Bohr’s atomic model was successful in explaining spectra of hydrogen or ions consisting of one
electron only e
...
Hydrogen and Li2+
2
...
Shortcomings (drawbacks) of Bohr’s atomic model
1
...
2
...
It does not explain the relative intensity of spectral lines (that is why some spectral lines are brighter
while others are dimmer)
...
The model does not explain how covalent bonding makes a molecule stable because the model does
not recognise the existence of sub-energy levels
...
The model viewed an electron as being placed at certain distance from the nucleus but it was proved
by Werner Heisenberg in his Heisenberg’s uncertainty principle, which states that, “It is impossible
at any moment to predict simultaneously the exact position and velocity of an electron in an
atom”
5
...
6
...
7
...
By definition:
Hyperfine spectral lines are extremely thin (fine) spectral lines which are formed after splitting one atomic
spectral line into two or more components
...
8
...
By definition:
Zeeman effect (or Zeeman splitting) is the splitting of a spectral line of atoms or molecules when subjected
to static magnetic field
...
Energy associated with an electron moving in a particular main energy level
Assume an electron (in the hydrogen atom) jump from one orbit of n energy level to new orbit of n1 energy
level with respective energy of E and E1
...
097 × 107 m−1
...
626 × 10−34 Js
c is the velocity of radiation = 2
...
626 × 10−34 × 2
...
097 × 107 (n2 −
1
1
)
n21
1
E1 − E = 2
...
e
...
6ev
n2
Then ∆E = E1 − E = 0 − E = −E
1
Thus –E = 2
...
179163×10−18
J
n2
But 1ev = 1
...
179163 ×10−18
J
n2
=
−2
...
6×10−19
Hence energy associated with an electron revolving in an energy level of n for hydrogen atom is given by:
−𝟏𝟑
...
6 ×𝟏𝟎−𝟏𝟗 J
−13
...
HEISENBERG’S UNCERTAINTY PRINCIPLE
It states that: It is not possible at any moment to predict accurately both position and velocity (or
momentum) of an electron in an atom simultaneously
...
- In another hand if radiations of low energy with large wavelength are used position of electrons cannot
be detected
...
To find the approximate position of an electron
Use ∆x × ∆p =
h
4π
Where ∆x is uncertainty in measuring in position which is given as an approximate distance from nucleus
in metres
...
The reader should understand that:
In actual sense the product ΔxΔp is always slightly greater than the value of
Heisenberg equation is written as:
ΔxΔp ≈
h
4π
(Or even ΔxΔp ≥
h
4π
)
h
4π
and hence sometimes
Digging Heisenberg uncertainty equation!
From the equation:
h
∆x∆p ≈ 4π ;
∆p = m∆v
Where ∆v is the uncertainty in the measurement of velocity (speed) of an electron
...
𝐡
It follows that: ∆𝐱 × 𝐦∆𝐯 ≈ 𝟒𝛑 … … … … … … … … … (𝐢)
Then from (i) above it is clearly understood that the right hand side of the equation, that is
h
4π
is always a
constant with a value of approximately 5
...
626 × 10−34 and π = 3
...
What
does this mean?
h
The fact that 4π is very small constant implies that if m could be a mass of large object like a car or stone,
𝐡
the product ∆𝐱 × 𝐦∆𝐯 would be very large compared to the value of
and hence making the
𝟒𝛑
Heisenberg uncertainty equation insignificant
...
Also the equation may be rearranged to give the following equation:
𝐡
∆𝐱 ≈ 𝟒𝛑×𝐦𝚫𝐯 … … … … … … …
...
-
That means if the velocity of the object is determined with more accuracy then its position will
determined with less accuracy as suggested in Heisenberg uncertainty principle
...
That means for macroscopic object whose mass is large, 𝚫𝐱 become very small and can be
neglected
...
So putting the two reasons together we may conclude that Heisenberg uncertainty principle is useless in the
daily life where macroscopic objects are commonly involved because:
-
h
For large mass, m, the product 4π × mΔv becomes very large compared to the value of 4π and
therefore making the Heisenberg uncertainity equation inapplicable
...
Be careful!
i)
Sometimes you may be given the constant ℏ(with value of about 1
...
626 × 10−34 Js )
...
∆P = 4π;
The equation may be re-written as
h
1
h
∆x∆p = 2 × 2π
h
2π
=ℏ
-
But
-
Hence ∆x∆p = ; where ℏ is the reduced Planck constant
...
To conclude:
Because of the Heisenberg uncertainty principle, the position of an electron moving with a definite velocity
cannot be determined exactly
...
The region in space around the nucleus where there is a maximum probability of finding an electron is
called the atomic orbital
...
Equating Plank’s and Einstein’s equation (Law of conservation of energy)
hf = mc2
c
But f = λ
Then
h
λ
hc
λ
= mc2
h
mc
= mc or λ =
Hence
λ=
𝐡
𝐦𝐜
The last equation is known as De Broglie equation and λ obtained according to the above formula is known
as De Broglie wavelength
...
-
In other words; De-Broglie concept is useless for macroscopic objects like stone because with large
mass the De-Broglie wavelength is too small to be measured
...
Wave-particle duality of an electron
An electron exhibit wave – particle duality also
...
Particle characteristics of an electron
As a particle an electron has mass, charge and momentum
...
Electron can also undergo diffraction and
can produce interference to justify its wave nature
...
It may behave as wave as well as particle
...
It shows the following
properties as a wave;
i)
It undergoes reflection
ii)
It undergoes refraction
iii)
It undergoes diffraction
iv)
It undergoes interference
v)
It undergoes polarisation
Despite of many common evidences of light behaving as a wave there are also some evidences which show
that the light is behaving like a particle
...
Presence of light particles (photons) can be verified by photoelectric effect
...
Photoelectric effect is very useful phenomenon on understanding wave- particle duality of electromagnetic
radiation
...
Wrong expectation!
It was thought that electrons were excited from the atom by absorbing the energy in the light wave (not
light particle!)
...
-
So with that point of view (of looking a light as merely a wave) it was expected that the more energy
the electrons could absorb, the more energy they could use to jump out
...
The only change which was noticed is that after increasing intensity of incident light the number of electrons
emitted was increased
...
-
A photon is almost massless particle carrying a small amount of energy
...
This justify that light comprises of particles (photons) and behave as a particles too
...